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THE PRINCIPLES OF CHEMISTRY

by

D. MENDELÉEFF

Translated from the Russian (Sixth Edition) by George Kamensky, A.R.S.M.
of the Imperial Mint, St Petersburg: Member of the Russian
Physico-Chemical Society

Edited by T. A. Lawson, B.Sc. PH.D.

Examiner in Coal-Tar Products to the City and Guilds of London
Institute Fellow of the Institute of Chemistry

In Two Volumes

VOLUME I.







Longmans, Green, and Co
39 Paternoster Row, London
New York and Bombay
1897

All rights reserved




                                 PREFACE

                                 TO THE

                           ENGLISH TRANSLATION


The first English edition of this work was published in 1891, and that a
second edition is now called for is, we think, a sufficient proof that
the enthusiasm of the author for his science, and the philosophical
method of his teaching, have been duly appreciated by English chemists.

In the scientific work to which Professor Mendeléeff's life has been
devoted, his continual endeavour has been to bring the scattered facts
of chemistry within the domain of law, and accordingly in his teaching
he endeavours to impress upon the student the _principles_ of the
science, the generalisations, so far as they have been discovered, under
which the facts naturally group themselves.

Of those generalisations the periodic law is perhaps the most important
that has been put forward since the establishment of the atomic theory.
It is therefore interesting to note that Professor Mendeléeff was led to
its discovery in preparing the first Russian edition of this book.

It is natural, too, that the further application and development of
that generalisation should be the principal feature of this, the latest
edition.

There are special difficulties in rendering the Russian language into
good English, and we are conscious that these have not been entirely
overcome. Doubtless also there are errors of statement which have
escaped correction, but we believe that the present edition will be
found better in both respects than its predecessor. We have thought
it our duty as translators to give as far as possible a faithful
reproduction of Professor Mendeléeff's work--the sixth Russian
edition--without amplifying or modifying his statements, and in this we
have the author's approval.

Although other duties have prevented Mr. Greenaway from undertaking the
care of the present edition, he has been kind enough to give us the
benefit of his suggestions on several points. We also wish to thank the
Managers of the Royal Institution for permission to reprint the lecture
delivered at the Royal Institution by Professor Mendeléeff (Appendix
I.), and to the Council of the Chemical Society for permission to
reprint the Faraday lecture which forms Appendix II.

In conclusion, we are indebted to Mr. F. Evershed, who has given us much
valuable assistance in revising the sheets for the press.

                                                                G. K.
                                                                T. A. L.
     _August 1897_




                            AUTHOR'S PREFACE

                                   TO

                        THE SIXTH RUSSIAN EDITION


This work was written during the years 1868-1870, its object being to
acquaint the student not only with the methods of observation, the
experimental facts, and the laws of chemistry, but also with the insight
given by this science into the unchangeable substratum underlying the
varying forms of matter.

If statements of fact themselves depend upon the person who observes
them, how much more distinct is the reflection of the personality of him
who gives an account of methods and of philosophical speculations which
form the essence of science! For this reason there will inevitably be
much that is subjective in every objective exposition of science. And as
an individual production is only significant in virtue of that which has
preceded and that which is contemporary with it, it resembles a mirror
which in reflecting exaggerates the size and clearness of neighbouring
objects, and causes a person near it to see reflected most plainly those
objects which are on the side to which it is directed. Although I have
endeavoured to make my book a true mirror directed towards the whole
domain of chemical changes, yet involuntarily those influences near to
me have been the most clearly reflected, the most brightly illuminated,
and have tinted the entire work with their colouring. In this way the
chief peculiarity of the book has been determined. Experimental and
practical data occupy their place, but the philosophical principles of
our science form the chief theme of the work. In former times sciences,
like bridges, could only be built up by supporting them on a few broad
buttresses and long girders. In addition to the exposition of the
principles of chemistry, it has been my desire to show how science has
now been built up like a suspension bridge, supported by the united
strength of a number of slender, but firmly-fixed, chains, which
individually are of little strength, and has thus been carried over
difficulties which before appeared insuperable. In comparing the science
of the past, the present, and the future, in placing the particulars
of its restricted experiments side by side with its aspirations after
unbounded and infinite truth, and in restraining myself from yielding to
a bias towards the most attractive path, I have endeavoured to incite
in the reader a spirit of inquiry, which, dissatisfied with speculative
reasonings alone, should subject every idea to experiment, encourage the
habit of stubborn work, and excite a search for fresh chains of evidence
to complete the bridge over the bottomless unknown. History proves that
it is possible by this means to avoid two equally pernicious extremes,
the Utopian--a visionary contemplation which proceeds from a current of
thought only--and the stagnant realism which is content with bare facts.
Sciences like chemistry, which deal with ideas as well as with material
substances, and create a possibility of immediately verifying that
which has been or may be discovered or assumed, demonstrate at every
step that the work of the past has availed much, and that without it it
would be impossible to advance into the ocean of the unknown. They also
show the possibility of becoming acquainted with fresh portions of this
unknown, and compel us, while duly respecting the teachings of history,
to cast aside classical illusions, and to engage in a work which not
only gives mental satisfaction but is also practically useful to all our
fellow-creatures.[1]

  [1] Chemistry, like every other science, is at once a means and an
      end. It is a means of attaining certain practical results. Thus,
      by its assistance, the obtaining of matter in its various forms
      is facilitated; it shows new possibilities of availing ourselves
      of the forces of nature, indicates the methods of preparing
      many substances, points out their properties, &c. In this sense
      chemistry is closely connected with the work of the manufacturer
      and the artisan, its sphere is active, and is a means of promoting
      general welfare. Besides this honourable vocation, chemistry
      has another. With it, as with every other elaborated science,
      there are many lofty aspirations, the contemplation of which
      serves to inspire its workers and adherents. This contemplation
      comprises not only the principal data of the science, but also the
      generally-accepted deductions, and also hypotheses which refer
      to phenomena as yet but imperfectly known. In this latter sense
      scientific contemplation varies much with times and persons, it
      bears the stamp of creative power, and embraces the highest forms
      of scientific progress. In that pure enjoyment experienced on
      approaching to the ideal, in that eagerness to draw aside the
      veil from the hidden truth, and even in that discord which exists
      between the various workers, we ought to see the surest pledges of
      further scientific progress. Science thus advances, discovering
      new truths, and at the same time obtaining practical results. The
      edifice of science not only requires material, but also a plan,
      and necessitates the work of preparing the materials, putting them
      together, working out the plans and the symmetrical proportions of
      the various parts. To conceive, understand, and grasp the whole
      symmetry of the scientific edifice, including its unfinished
      portions, is equivalent to tasting that enjoyment only conveyed by
      the highest forms of beauty and truth. Without the material, the
      plan alone is but a castle in the air--a mere possibility; whilst
      the material without a plan is but useless matter. All depends
      on the concordance of the materials with the plan and execution,
      and the general harmony thereby attained. In the work of science,
      the artisan, architect, and creator are very often one and the
      same individual; but sometimes, as in other walks of life, there
      is a difference between them; sometimes the plan is preconceived,
      sometimes it follows the preparation and accumulation of the
      raw material. Free access to the edifice of science is not only
      allowed to those who devised the plan, worked out the detailed
      drawings, prepared the materials, or piled up the brickwork, but
      also to all those who are desirous of making a close acquaintance
      with the plan, and wish to avoid dwelling in the vaults or in the
      garrets where the useless lumber is stored.

      Knowing how contented, free, and joyful is life in the realm of
      science, one fervently wishes that many would enter its portals.
      On this account many pages of this treatise are unwittingly
      stamped with the earnest desire that the habits of chemical
      contemplation which I have endeavoured to instil into the minds
      of my readers will incite them to the further study of science.
      Science will then flourish in them and by them, on a fuller
      acquaintance not only with that little which is enclosed within
      the narrow limits of my work, but with the further learning which
      they must imbibe in order to make themselves masters of our
      science and partakers in its further advancement.

      Those who enlist in the cause of science have no reason to fear
      when they remember the urgent need for practical workers in the
      spheres of agriculture, arts, and manufacture. By summoning
      adherents to the work of theoretical chemistry, I am confident
      that I call them to a most useful labour, to the habit of dealing
      correctly with nature and its laws, and to the possibility of
      becoming truly practical men. In order to become actual chemists,
      it is necessary for beginners to be well and closely acquainted
      with three important branches of chemistry--analytical, organic,
      and theoretical. That part of chemistry which is dealt with in
      this treatise is only the groundwork of the edifice. For the
      learning and development of chemistry in its truest and fullest
      sense, beginners ought, in the first place, to turn their
      attention to the practical work of analytical chemistry; in the
      second place, to practical and theoretical acquaintance with some
      special chemical question, studying the original treatises of the
      investigators of the subject (at first, under the direction of
      experienced teachers), because in working out particular facts the
      faculty of judgment and of correct criticism becomes sharpened; in
      the third place, to a knowledge of current scientific questions
      through the special chemical journals and papers, and by
      intercourse with other chemists. The time has come to turn aside
      from visionary contemplation, from platonic aspirations, and from
      classical verbosity, and to enter the regions of actual labour for
      the common weal, to prove that the study of science is not only
      air excellent education for youth, but that it instils the virtues
      of industry and veracity, and creates solid national wealth,
      material and mental, which without it would be unattainable.
      Science, which deals with the infinite, is itself without bounds.

Thus the desire to direct those thirsting for truth to the pure source
of the science of the forces acting throughout nature forms the
first and most important aim of this book. The time has arrived when
a knowledge of physics and chemistry forms as important a part of
education as that of the classics did two centuries ago. In those days
the nations which excelled in classical learning stood foremost, just
as now the most advanced are those which are superior in the knowledge
of the natural sciences, for they form the strength and characteristic
of our times. In following the above and chief aim, I set myself a
second object: to furnish a text-book for an elementary knowledge of
chemistry and so satisfy a want which undoubtedly exists among students
and those who have recourse to chemistry either as a source of truth or
welfare.[2] Hence, although the fundamental object of this work was to
express and embrace the general chemical teaching of the present day
from a personal point of view, I have nevertheless striven throughout
to maintain such a level as would render the 'Principles of Chemistry'
accessible to the beginner. Many aspects of this work are determined
by this combination of requirements which frequently differ widely. An
issue was only possible under one condition, _i.e._ not to be carried
away by what appears to be a plausible theory in explaining individual
facts and to always endeavour to transmit the simple truth of a given
fact, extracting it from the vast store of the literature of the
subject and from tried personal experience. In publishing a new edition
of this work I have striven to add any facts of importance recently
discovered[3] and to revise the former edition in the above spirit. With
this object I have entirely gone over this edition, and a comparison
of it with the former one will show that the additions and alterations
have cost as much labour as many chapters of the work. I also wished
to show in an elementary treatise on chemistry the striking advantages
gained by the application of the periodic law, which I first saw in
its entirety in the year 1869 when I was engaged in writing the first
edition of this book, in which, indeed, the law was first enunciated.
At that time, however, this law was not established so firmly as now,
when so many of its consequences have been verified by the researches
of numerous chemists, and especially by Roscoe, Lecoq de Boisbaudran,
Nilson, Brauner, Thorpe, Carnelley, Laurie, Winkler, and others. The,
to me, unexpectedly rapid success with which the teaching of the
periodicity of the elements has spread in our science, and perhaps also,
the perseverance with which I collected in this work, and upon a new
plan, the most important data respecting the elements and their mutual
relations, explained sufficiently the fact that the former (5th, 1889)
edition of my work has been translated into English[4] and German[5] and
is being translated into French.[6] Deeply touched by the favourable
opinions expressed by English men of science upon my book, I ascribe
them chiefly to the periodic law placed at the basis of my treatise and
especially of the second part of the book, which contains a large amount
of data having a special and sometimes quite unexpected, bearing from
the point of view of this law. As the entire scheme of this work is
subordinated to the law of periodicity, which may be illustrated in a
tabular form by placing the elements in series, groups, and periods, two
such tables are given at the end of this preface.

  [2] I recommend those who are commencing the study of chemistry with my
      book _to first read only what is printed in the large type_,
      because in that part I have endeavoured to concentrate all the
      fundamental, indispensable knowledge required for that study.
      In the footnotes, printed in small type (which should be read
      only after the large text has been mastered), certain details
      are discussed; they are either further examples, or debatable
      questions on existing ideas which I thought useful to lay
      before those entering into the sphere of science, or certain
      historical and technical details which might be withdrawn from the
      fundamental portion of the book. Without intending to attain in my
      treatise to the completeness of a work of reference, I have still
      endeavoured to express the principal developments of science as
      they concern the chemical elements viewed in that aspect in which
      they appeared to me after long continued study of the subject and
      participation in the contemporary advance of knowledge.

      I have also placed my personal views, suppositions, and arguments
      in the footnotes, which are chiefly designed for details and
      references. But I have endeavoured to avoid here, as in the
      text, not only all that I consider doubtful, but also those
      details which belong either to special branches of chemistry
      (for instance, to analytical, organic, physical, theoretical,
      physiological, agricultural, or technical chemistry), or to
      different branches of natural science which are more and more
      coming into closer and closer contact with chemistry. Chemistry, I
      am convinced, must occupy a place among the natural sciences side
      by side with mechanics; for mechanics treats of matter as a system
      of ponderable points having scarcely any individuality and only
      standing in a certain state of mobile equilibrium. For chemistry,
      matter is an entire world of life, with an infinite variety of
      individuality both in the elements and in their combinations.
      In studying the general uniformity from a mechanical point of
      view, I think that the highest point of knowledge of nature
      cannot be attained without taking into account the individuality
      of things in which chemistry is set to seek for general higher
      laws. Mechanics may be likened to the science of statesmanship,
      chemistry to the sciences of jurisprudence and sociology. The
      general universe could not be built up without the particular
      individual universe, and would be a dry abstract were it not
      enlivened by the real variety of the individual world. Mechanics
      forms the classical basis of natural philosophy, while chemistry,
      as a comparatively new and still young science, already strives
      to--and will, in the future introduce a new, living aspect into
      the philosophy of nature; all the more as chemistry alone is never
      at rest or anywhere dead--its vital action has universal sway, and
      inevitably determines the general aspect of the universe. Just
      as the microscope and telescope enlarge the scope of vision, and
      discover life in seeming immobility, so chemistry, in discovering
      and striving to discern the life of the invisible world of atoms
      and molecules and their ultimate limit of divisibility, will
      clearly introduce new and important problems into our conception
      of nature. And I think that its _rôle_, which is now considerable,
      will increase more and more in the future; that is, I think that
      in its further development it will occupy a place side by side
      with mechanics for the comprehension of the secrets of nature. But
      here we require some second Newton; and I have no doubt that he
      will soon appear.

  [3] I was much helped in gathering data from the various chemical
      journals of the last five years by the abstracts made for me by
      Mr. Y. V. Kouriloff, to whom I tender my best thanks.

  [4] The English translation was made by G. Kamensky, and edited by
      A. J. Greenaway; published by Longmans, Green & Co.

  [5] The German translation was made by L. Jawein and A. Thillot;
      published by Ricker (St. Petersburg).

  [6] The French translation has been commenced by E. Achkinasi and
      H. Carrion from the fifth edition, and is published by Tignol
      (Paris).

In this the sixth edition I have not altered any essential feature
of the original work, and have retained those alterations which were
introduced into the fifth edition.[7] I have, however, added many
newly discovered facts, and in this respect it is necessary to say a
few words. Although all aspects of the simplest chemical relations are
as far as possible equally developed in this book, yet on looking back
I see that I have, nevertheless, given most attention to the so-called
indefinite compounds examples of which may be seen in solutions. I recur
repeatedly to them, and to all the latest data respecting them, for
in them I see a starting point for the future progress of our science
and to them I affiliate numerous instances of definite compounds,
beginning with alloys and silicates and ending with complex acids.
There are two reasons for this. In the first place, this subject has
deeply interested me from my youth; I have devoted a portion of my
own researches to it, and therefore it occupied an important position
even in the first edition of my book; besides which all that has been
subsequently accomplished in our science, especially during the last
five or six years, shows that at the present day an interest in these
questions plays an important part in the minds of a large circle of
contemporary workers in chemistry. This personal attachment, if I may
so call it, to the question of solutions and such indefinite compounds,
must involuntarily have impressed itself upon my work, and in the later
editions I have even had to strive not to give this subject a greater
development than previously, so great was the material accumulated,
which however does not yet give us the right to consider even the most
elementary questions respecting solutions as solved. Thus, we cannot
yet say what a solution really is. My own view is that a solution is
a homogeneous liquid system of unstable dissociating compounds of
the solvent with the substance dissolved. But although such a theory
explains much to me, I cannot consider my opinion as proved, and
therefore give it with some reserve as one of several hypotheses.[8] As
a subject yet far from solved, I might naturally have ignored it, or
only mentioned it cursorily, but such a treatment of solutions, although
usual in elementary treatises on chemistry, would not have answered my
views upon the subject of our science, and I wished that the reader
might find in my book beyond everything an expression of all that a
study of the subject built up for me. If in solutions I see and can
frequently prove distinct evidences of the existence of those definite
compounds which form the more generalised province of chemical data, I
could not refrain from going into certain details respecting solutions;
otherwise, there would have remained no trace of that general idea, that
in them we have only a certain instance of ordinary definite or atomic
compounds, subject to Dalton's laws. Having long had this idea, I wished
to impress it upon the reader of my book, and it is this desire which
forms the second of those chief reasons why I recur so frequently to
solutions in this work. At present, my ideas respecting solutions are
shared by few, but I trust that by degrees the instances I give will
pave the way for their general recognition, and it is my hope that they
may find adherents among those of my readers who are in a position to
work out by experiment this difficult but highly interesting problem.

  [7] The fifth edition was not only considerably enlarged, compared with
      the preceding, but also the foundations of the periodic system
      of the elements were placed far more firmly in it than in the
      former editions. The subject-matter was also divided into
      text and footnotes, which contained details unnecessary for a
      first acquaintance with chemistry. The fifth edition sold out
      sooner than I expected, so that instead of issuing supplements
      (containing the latest discoveries in chemistry), as I had
      proposed, I was obliged to publish the present entirely new
      edition of the work.

  [8] This hypothesis is not only mentioned in different parts of this
      book, but is partly (from the aspect of the specific gravity of
      solutions) developed in my work, _The Investigation of Solutions
      from their Specific Gravity_, 1887.

In conclusion, I desire to record my thanks to V. D. Sapogenikoff, who
has corrected the proofs of the whole of this edition and compiled the
indexes which greatly facilitate the search for those details which are
scattered throughout the work.

                                                          D. MENDELÉEFF.


                                 TABLE I
           _Distribution of the Elements in Groups and Series_
     +--------+-------+----------+----------+----------+----------+
     | Group  |   I.  |    II.   |   III.   |    IV.   |    V.    |
     +--------+-------+----------+----------+----------+----------+
     |Series 1|     H |    --    |    --    |    --    |    --    |
     |        |       |          |          |          |          |
     |  "    2| Li    |  Be      |  B       |  C       |  N       |
     |        |       |          |          |          |          |
     |  "    3|    Na |       Mg |       Al |       Si |        P |
     |        |       |          |          |          |          |
     |  "    4| K     |  Ca      |  Sc      |  Ti      |  V       |
     |        |       |          |          |          |          |
     |  "    5|   (Cu)|       Zn |       Ga |       Ge |       As |
     |        |       |          |          |          |          |
     |  "    6| Rb    |  Sr      |  Y       |  Zr      |  Nb      |
     |        |       |          |          |          |          |
     |  "    7|   (Ag)|       Cd |       In |       Sn |       Sb |
     |        |       |          |          |          |          |
     |  "    8| Cs    |  Ba      |  La      |  Ce      |  Di?     |
     |        |       |          |          |          |          |
     |  "    9|    -- |       -- |       -- |       -- |       -- |
     |        |       |          |          |          |          |
     |  "   10| --    |  --      |  Yb      |  --      |  Ta      |
     |        |       |          |          |          |          |
     |  "   11|   (Au)|       Hg |       Tl |       Pb |       Bi |
     |        |       |          |          |          |          |
     |  "   12| --    |  --      |  --      |  Th      |  --      |
     |        |       |          |          |          |          |
     +--------+-------+----------+----------+----------+----------+
     |        |R_{2}O |R_{2}O_{2}|R_{2}O_{3}|R_{2}O_{4}|R_{2}O_{5}|
     |        |       |          |          |          |          |
     |        |  --   |RO        |    --    |RO_{2}    |    --    |
     |        |       |          |          |          |          |
     |        |  --   |    --    |    --    |RH_{4}    |RH_{3}    |
     |        |       |          |          |          |          |
     +--------+-------+----------+----------+----------+----------+

         +---------+----------+----------+--------------------+
         |  Group  |    VI.   |   VII.   |       VIII.        |
         +---------+----------+----------+--------------------+
         | Series 1|    --    |    --    |                    |
         |         |          |          |                    |
         |   "    2|  O       |  F       |                    |
         |         |          |          |                    |
         |   "    3|        S |       Cl |                    |
         |         |          |          |                    |
         |   "    4|  Cr      |  Mn      |    Fe Co Ni Cu     |
         |         |          |          |                    |
         |   "    5|       Se |       Br |                    |
         |         |          |          |                    |
         |   "    6|  Mo      |  --      |    Ru Rh Pd Ag     |
         |         |          |          |                    |
         |   "    7|       Te |       I  |                    |
         |         |          |          |                    |
         |   "    8|  --      |  --      |    -- -- -- --     |
         |         |          |          |                    |
         |   "    9|       -- |       -- |                    |
         |         |          |          |                    |
         |   "   10|  W       |  --      |    Os Ir Pt Au     |
         |         |          |          |                    |
         |   "   11|       -- |       -- |                    |
         |         |          |          |                    |
         |   "   12|  U       |  --      |                    |
         |         |          |          |                    |
         +---------+----------+----------+--------------------+
         |         |R_{2}O_{6}|R_{2}O_{7}| Higher oxides      |
         |         |          |          |                    |
         |         |RO_{3}    |    --    | RO_{4}             |
         |         |          |          |                    |
         |         |RH_{2}    |    RH    | Hydrogen compounds |
         +---------+----------+----------+--------------------+

                                TABLE II

          _Periodic System and Atomic Weights of the Elements_
            (_Giving the pages on which they are described_)
  +---------------------+--------------+--------------+--------------+
  |                     | 2nd Series,  |              |              |
  |                     |  Typical     |     4th      |     6th      |
  |                     |  elements    |    Series    |    Series    |
  +-------+-------------+--------------+--------------+--------------+
  |    I. |             |    Li 7      |     K 39     |    Rb 86     |
  |       |             |              |     ====     |              |
  |       |             | vol. i. 574  | vol. i. 558  | vol. i. 576  |
  |       |             |              |              |              |
  |   II. |             |    Be 9      |    Ca 40     |    Sr 88     |
  |       |             |              |    -----     |              |
  |       |             | vol. i. 618  | vol. i. 590  | vol. i. 614  |
  |       |             |              |              |              |
  |  III. |             |    B 11      |    Sc 44     |     Y 89     |
  |       |             |    ----      |              |              |
  |       |             | vol. ii. 60  | vol. ii. 94  | vol. ii. 93  |
  |       |             |              |              |              |
  |   IV. |             |    C 12      |    Ti 48     |    Zr 91     |
  |       |             |    ====      |              |              |
  |       |             | vol. i. 338  | vol. ii. 144 | vol. ii. 93  |
  |       |             |              |              |              |
  |    V. |             |    N 14      |     V 51     |    Nb 94     |
  |       |             |    ====      |              |              |
  |       |             | vol. i. 223  | vol. ii. 194 | vol. ii. 197 |
  |       |             |              |              |              |
  |   VI. |             |    O 16      |    Cr 52     |    Mo 96     |
  |       |             |    ----      |    -----     |              |
  |       |             | vol. i. 155  | vol. ii. 276 | vol. ii. 290 |
  |       |             |              |              |              |
  |  VII. |             |    F 19      |    Mn 55     |     ? 99     |
  |       |             |              |    =====     |              |
  |       |             | vol. i. 489  | vol. ii. 303 |              |
  |       |             |              |              |              |
  |       |             |              |    Fe 56     |    Ru 102    |
  |       |             |              |    =====     |              |
  |       |             |              | vol. ii. 317 | vol. ii. 369 |
  |       |             |              |              |              |
  | VIII. |             |              |    Co 59     |    Rh 103    |
  |       |             |              | vol. ii. 353 | vol. ii. 369 |
  |       |             |              |              |              |
  |       |             |              |    Ni 59·5   |    Pd 106    |
  |       |             |              | vol. ii. 353 | vol. ii. 369 |
  |       |             |              |              |              |
  |       |             |  3rd Series  |  5th Series  |  7th Series  |
  |       |             |              |              |              |
  |    I. |     H 1     |    Na 23     |     Cu 64    |    Ag 108    |
  |       |    ----     |              |              |              |
  |       | vol. i. 129 | vol. i. 533  | vol. ii. 398 | vol. ii. 415 |
  |       |             |              |              |              |
  |   II. |             |   Mg 24      |     Zn 65    |    Cd 112    |
  |       |             | vol. i. 590  | vol. ii. 39  | vol. ii. 47  |
  |       |             |              |              |              |
  |  III. |             |   Al 27      |     Ga 70    |    In 114    |
  |       |             | vol. ii. 70  | vol. ii. 90  | vol. ii. 91  |
  |       |             |              |              |              |
  |   IV. |             |   Si 28      |     Ge 72    |    Sn 119    |
  |       |             | vol. ii. 99  | vol. ii. 124 | vol. ii. 125 |
  |       |             |              |              |              |
  |    V. |             |    P 31      |     As 75    |    Sb 120    |
  |       |             |              |     -----    |              |
  |       |             | vol. ii. 149 | vol. ii. 179 | vol. ii. 186 |
  |       |             |              |              |              |
  |   VI. |             |    S 32      |     Se 79    |    Te 125    |
  |       |             | vol. ii. 200 | vol. ii. 270 | vol. ii. 270 |
  |       |             |              |              |              |
  |  VII. |             |   Cl 35·5    |     Br 80    |     I 127    |
  |       |             |              |     -----    |              |
  |       |             | vol. i. 459  | vol. i. 494  | vol. i. 496  |
  +-------+-------------+--------------+--------------+--------------+

          +-------+-------------+--------------+--------------+
          |       |             |              |              |
          |       |     8th     |     10th     |      12th    |
          |       |    Series   |    Series    |     Series   |
          +---- --+-------------+--------------+--------------+
          |    I. |    Cs 133   |      --      |      --      |
          |       | vol. i. 576 |              |              |
          |       |             |              |              |
          |   II. |    Ba 137   |      --      |      --      |
          |       |    ------   |              |              |
          |       | vol. i. 614 |              |              |
          |       |             |              |              |
          |  III. |    La 138   |    Yb 173    |      --      |
          |       | vol. ii. 93 | vol. ii. 93  |              |
          |       |             |              |              |
          |   IV. |    Ce 140   |     ? 178    |    Th 232    |
          |       | vol. ii. 93 |              | vol. ii. 148 |
          |       |             |              |              |
          |    V. |  ? Di 142   |    Ta 183    |      --      |
          |       | vol. ii. 93 | vol. ii. 197 |              |
          |       |             |              |              |
          |   VI. |     --      |     W 184    |     U 239    |
          |       |             | vol. ii. 290 | vol. ii. 297 |
          |       |             |              |              |
          |  VII. |     --      |      --      |      --      |
          |       |             |              |              |
          |       |     --      |    Os 192    |              |
          |       |             | vol. ii. 369 |              |
          |       |             |              |              |
          | VIII. |     --      |    Ir 193    |              |
          |       |             | vol. ii. 369 |              |
          |       |             |              |              |
          |       |     --      |    Pt 196    |              |
          |       |             |    ------    |              |
          |       |             | vol. ii. 369 |              |
          |       |             |              |              |
          |       |     9th     |     11th     |              |
          |       |    Series   |    Series    |              |
          |       |             |              |              |
          |    I. |     --      |    Au 197    |              |
          |       |             | vol. ii. 442 |              |
          |       |             |              |              |
          |   II. |     --      |    Hg 200    |              |
          |       |             | vol. ii. 48  |              |
          |       |             |              |              |
          |  III. |     --      |    Tl 204    |              |
          |       |             | vol. ii. 91  |              |
          |       |             |              |              |
          |   IV. |     --      |    Pb 207    |              |
          |       |             | vol. ii. 134 |              |
          |       |             |              |              |
          |    V. |     --      |    Bi 209    |              |
          |       |             | vol. ii. 189 |              |
          |       |             |              |              |
          |   VI. |     --      |      --      |              |
          |       |             |              |              |
          |  VII. |     --      |      --      |              |
          +-------+-------------+--------------+--------------+

_Note._--Two lines under the elements indicate those which are very
widely distributed in nature; one line indicates those which, although
not so frequently met with, are of general use in the arts and
manufactures.




                                CONTENTS
                                   OF
                            THE FIRST VOLUME


                                                                    PAGE

  TRANSLATORS' PREFACE                                                 v

  AUTHOR'S PREFACE TO THE SIXTH RUSSIAN EDITION                      vii

  TABLE OF THE DISTRIBUTION OF THE ELEMENTS IN GROUPS AND SERIES      xv

  TABLE OF THE PERIODIC SYSTEM AND ATOMIC WEIGHTS OF THE ELEMENTS    xvi

  INTRODUCTION                                                         1

  CHAP.

     I. ON WATER AND ITS COMPOUNDS                                    40

    II. THE COMPOSITION OF WATER. HYDROGEN                           113

   III. OXYGEN AND THE CHIEF ASPECTS OF ITS SALINE COMBINATIONS      152

    IV. OZONE AND HYDROGEN PEROXIDE. DALTON'S LAW                    198

     V. NITROGEN AND AIR                                             223

    VI. THE COMPOUNDS OF NITROGEN WITH HYDROGEN AND OXYGEN           246

   VII. MOLECULES AND ATOMS. THE LAWS OF GAY-LUSSAC AND
        AVOGADRO-GERHARDT                                            299

  VIII. CARBON AND THE HYDROCARBONS                                  338

    IX. COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN                 379

     X. SODIUM CHLORIDE. BERTHOLLET'S LAWS. HYDROCHLORIC ACID        417

    XI. THE HALOGENS: CHLORINE, BROMINE, IODINE AND FLUORINE         459

   XII. SODIUM                                                       513

  XIII. POTASSIUM, RUBIDIUM, CÆSIUM AND LITHIUM. SPECTRUM ANALYSIS   543

   XIV. THE VALENCY AND SPECIFIC HEAT OF THE METALS. MAGNESIUM,
        CALCIUM, STRONTIUM, BARIUM, AND BERYLLIUM                    581




                         PRINCIPLES OF CHEMISTRY




                              INTRODUCTION


The study of natural science, whose rapid development dates from
the days of Galileo ([+]1642) and Newton ([+]1727), and its closer
application to the external universe[1] led to the separation of
Chemistry as a particular branch of natural philosophy, not only owing
to the increasing store of observations and experiments relating to the
mutual transformations of substances, but also, and more especially,
because in addition to gravity, cohesion, heat, light and electricity
it became necessary to recognise the existence of particular internal
forces in the ultimate parts of all substances, forces which make
themselves manifest in the transformations of substances into one
another, but remain hidden (latent) under ordinary circumstances, and
whose existence cannot therefore be directly apprehended, and so for
a long time remained unrecognised. The primary object of chemistry is
the study of the homogeneous substances[2] of which all the objects of
the universe are made up, with the transformations of these substances
into each other, and with the phenomena[3] which accompany such
transformations. Every chemical change or reaction,[4] as it is called,
can only take place under a condition of most intimate and close contact
of the re-acting substances,[5] and is determined by the forces proper
to the smallest invisible particles (molecules) of matter. We must
distinguish three chief classes of chemical transformations.

  [1] The investigation of a substance or a natural phenomenon consists
      (_a_) in determining the relation of the object under examination
      to that which is already known, either from previous researches,
      or from experiment, or from the knowledge of the common
      surroundings of life--that is, in determining and expressing the
      quality of the unknown by the aid of that which is known; (_b_)
      in measuring all that which can be subjected to measurement,
      and thereby denoting the quantitative relation of that under
      investigation to that already known and its relation to the
      categories of time, space, temperature, mass, &c.; (_c_) in
      determining the position held by the object under investigation
      in the system of known objects guided by both qualitative and
      quantitative data; (_d_) in determining, from the quantities which
      have been measured, the empirical (visible) dependence (function,
      or 'law,' as it is sometimes termed) of variable factors--for
      instance, the dependence of the composition of the substance on
      its properties, of temperature on time, of time on locality, &c.;
      (_e_) in framing hypotheses or propositions as to the actual cause
      and true nature of the relation between that studied (measured
      or observed) and that which is known or the categories of time,
      space, &c.; (_f_) in verifying the logical consequences of the
      hypotheses by experiment; and (_g_) in advancing a theory which
      shall account for the nature of the properties of that studied in
      its relations with things already known and with those conditions
      or categories among which it exists. It is certain that it is
      only possible to carry out these investigations when we have
      taken as a basis some incontestable fact which is self-evident
      to our understanding; as, for instance, number, time, space,
      motion, or mass. The determination of such primary or fundamental
      conceptions, although not excluded from the possibility of
      investigation, frequently does not subject itself to our present
      mode of scientific generalisation. Hence it follows that in the
      investigation of anything, there always remains something which is
      accepted without investigation, or admitted as a known factor. The
      axioms of geometry may be taken as an example. Thus in the science
      of biology it is necessary to admit the faculty of organisms for
      multiplying themselves, as a conception whose meaning is as yet
      unknown. In the study of chemistry, too, the notion of elements
      must be accepted almost without any further analysis. However, by
      first investigating that which is visible and subject to direct
      observation by the organs of the senses, we may hope that in
      the first place hypotheses will be arrived at, and afterwards
      theories of that which has now to be placed at the basis of our
      investigations. The minds of the ancients strove to seize at once
      the very fundamental categories of investigation, whilst all the
      successes of recent knowledge are based on the above-cited method
      of investigation without the determination of 'the beginning of
      all beginnings.' By following this inductive method, the _exact
      sciences_ have already succeeded in becoming accurately acquainted
      with much of the invisible world, which directly is imperceptible
      to the organs of sense (for example, the molecular motion of all
      bodies, the composition of the heavenly luminaries, the paths
      of their motion, the necessity for the existence of substances
      which cannot be subjected to experiment, &c.), and have verified
      the knowledge thus obtained, and employed it for increasing the
      interests of humanity. It may therefore be safely said that _the
      inductive method of investigation_ is a more perfect mode of
      acquiring knowledge than the deductive method alone (starting from
      a little of the unknown accepted as incontestable to arrive at the
      much which is visible and observable) by which the ancients strove
      to embrace the universe. By investigating the universe by an
      inductive method (endeavouring from the much which is observable
      to arrive at a little which may be verified and is indubitable)
      the new science refuses to recognise dogma as truth, but through
      _reason_, by a slow and laborious method of investigation, strives
      for and attains to true deductions.

  [2] A substance or material is that which occupies space and has
      weight; that is, which presents a mass attracted by the earth and
      by other masses of material, and of which the _objects_ of nature
      are composed, and by means of which the motions and _phenomena_
      of nature are accomplished. It is easy to discover by examining
      and investigating, by various methods, the objects met with
      in nature and in the arts, that some of them are homogeneous,
      whilst others are composed of a mixture of several homogeneous
      substances. This is most clearly apparent in solid substances.
      The metals used in the arts (for example, gold, iron, copper)
      must be homogeneous, otherwise they are brittle and unfit for
      many purposes. Homogeneous matter exhibits similar properties in
      all its parts. By breaking up a homogeneous substance we obtain
      parts which, although different in form, resemble each other in
      their properties. Glass, pure sugar, marble, &c., are examples of
      homogeneous substances. Examples of non-homogeneous substances
      are, however, much more frequent in nature and the arts. Thus
      the majority of the rocks are not homogeneous. In porphyries
      bright pieces of a mineral called 'orthoclase' are often seen
      interspersed amongst the dark mass of the rock. In ordinary red
      granite it is easy to distinguish large pieces of orthoclase mixed
      with dark semi-transparent quartz and flexible laminæ of mica.
      Similarly, plants and animals are non-homogeneous. Thus, leaves
      are composed of a skin, fibre, pulp, sap, and a green colouring
      matter. As an example of those non-homogeneous substances which
      are produced artificially, gunpowder may be cited, which is
      prepared by mixing together known proportions of sulphur, nitre,
      and charcoal. Many liquids, also, are not homogeneous, as may be
      observed by the aid of the microscope, when drops of blood are
      seen to consist of a colourless liquid in which red corpuscles,
      invisible to the naked eye owing to their small size, are floating
      about. It is these corpuscles which give blood its peculiar
      colour. Milk is also a transparent liquid, in which microscopical
      drops of fat are floating, which rise to the top when milk is
      left at rest, forming cream. It is possible to extract from every
      non-homogeneous substance those homogeneous substances of which
      it is made up. Thus orthoclase may he separated from porphyry by
      breaking it off. So also gold is extracted from auriferous sand by
      washing away the mixture of clay and sand. Chemistry deals only
      with the homogeneous substances met with in nature, or extracted
      from natural or artificial non-homogeneous substances. The various
      mixtures found in nature form the subjects of other natural
      sciences--as geognosy, botany, zoology, anatomy, &c.

  [3] All those events which are accomplished by substances in time are
      termed 'phenomena.' Phenomena in themselves form the fundamental
      subject of the study of physics. Motion is the primary and
      most generally understood form of phenomenon, and therefore we
      endeavour to reason about other phenomena as clearly as when
      dealing with motion. For this reason mechanics, which treats of
      motion, forms the fundamental science of natural philosophy, and
      all other sciences endeavour to reduce the phenomena with which
      they are concerned to mechanical principles. Astronomy was the
      first to take to this path of reasoning, and succeeded in many
      cases in reducing astronomical to purely mechanical phenomena.
      Chemistry and physics, physiology and biology are proceeding in
      the same direction. One of the most important questions of all
      natural science, and one which has been handed down from the
      philosophers of classic times, is, whether the comprehension of
      all that is visible can be reduced to motion? Its participation
      in all, from the 'fixed' stars to the most minute parts of the
      coldest bodies (Dewar, in 1894 showed that many substances cooled
      to -180° fluoresce more strongly than at the ordinary temperature;
      _i.e._ that there is a motion in them which produces light) must
      now be recognised as undoubtable from direct experiment and
      observation, but it does not follow from this that by motion alone
      can all be explained. This follows, however, from the fact that
      we cannot apprehend motion otherwise than by recognising matter
      in a state of motion. If light and electricity be understood as
      particular forms of motion, then we must inevitably recognise
      the existence of a peculiar luminiferous (universal) ether as
      a material, transmitting this form of motion. And so, under
      the present state of knowledge, it is inevitably necessary to
      recognise the particular categories, motion and matter, and as
      chemistry is more closely concerned with the various forms of the
      latter, it should, together with mechanics or the study of motion,
      lie at the basis of natural science.

  [4] The verb 'to react' means to act or change chemically.

  [5] If a phenomenon proceeds at visible or measurable distances (as,
      for instance, magnetic attraction or gravity), it cannot be
      described as chemical, since these phenomena only take place at
      distances immeasurably small and undistinguishable to the eye or
      the microscope; that is to say, they are purely molecular.

1. _Combination_ is a reaction in which the union of two substances
yields a new one, or in general terms, from a given number of
substances, a lesser number is obtained. Thus, by heating a mixture of
iron and sulphur[6] a single new substance is produced, iron sulphide,
in which the constituent substances cannot be distinguished even by
the highest magnifying power. Before the reaction, the iron could be
separated from the mixture by a magnet, and the sulphur by dissolving it
in certain oily liquids;[7] in general, before combination they might
be mechanically separated from each other, but after combination both
substances penetrate into each other, and are then neither mechanically
separable nor individually distinguishable. As a rule, reactions of
direct combination are accompanied by an evolution of heat, and the
common case of combustion, evolving heat, consists in the combination
of combustible substances with a portion (oxygen) of the atmosphere,
the gases and vapours contained in the smoke being the products of
combination.

2. Reactions of _decomposition_ are cases the reverse of those of
combination, that is, in which one substance gives two--or, in general,
a given number of substances a greater number. Thus, by heating wood
(and also coal and many animal or vegetable substances) without access
to air, a combustible gas, a watery liquid, tar, and carbon are
obtained. It is in this way that tar, illuminating gas, and charcoal are
prepared on a large scale.[8] All limestones, for example, flagstones,
chalk, or marble, are decomposed by heating to redness into lime and a
peculiar gas called carbonic anhydride. A similar decomposition, taking
place, however, at a much lower temperature, proceeds with the green
copper carbonate which is contained in natural malachite. This example
will be studied more in detail presently. Whilst heat is evolved in the
ordinary reactions of combination, it is, on the contrary, absorbed in
the reactions of decomposition.

3. The third class of chemical reactions--where the number of re-acting
substances is equal to the number of substances formed--may be
considered as a simultaneous decomposition and combination. If, for
instance, two compounds A and B are taken and they react on each other
to form the substances C and D, then supposing that A is decomposed
into D and E, and that E combines with B to form C, we have a reaction
in which two substances A, or D E, and B were taken and two others
C, or E B, and D were produced. Such reactions ought to be placed
under the general term of reactions of '_rearrangement_,' and the
particular case where two substances give two fresh ones, reactions of
'_substitution_.'[9] Thus, if a piece of iron be immersed in a solution
of blue vitriol (copper sulphate), copper is formed--or, rather,
separated out, and green vitriol (iron sulphate, which only differs
from the blue vitriol in that the iron has replaced the copper) is
obtained in solution. In this manner iron may be coated with copper, so
also copper with silver; such reactions are frequently made use of in
practice.

  [6] For this purpose a piece of iron may be made red hot in a forge,
      and then placed in contact with a lump of sulphur, when iron
      sulphide will be obtained as a molten liquid, the combination
      being accompanied by a visible increase in the glow of the iron.
      Or else iron filings are mixed with powdered sulphur in the
      proportion of 5 parts of iron to 3 parts of sulphur, and the
      mixture placed in a glass tube, which is then heated in one place.
      Combination does not commence without the aid of external heat,
      but when once started in any portion of the mixture it extends
      throughout the entire mass, because the portion first heated
      evolves sufficient heat in forming iron sulphide to raise the
      adjacent parts of the mixture to the temperature required for
      starting the reaction. The rise in temperature thus produced is so
      high as to soften the glass tube.

  [7] Sulphur is slightly soluble in many thin oils; it is very soluble
      in carbon bisulphide and in some other liquids. Iron is insoluble
      in carbon bisulphide, and the sulphur therefore can be dissolved
      away from the iron.

  [8] Decomposition of this kind is termed 'dry distillation,' because,
      as in distillation, the substance is heated and vapours are
      given off which, on cooling, condense into liquids. In general,
      decomposition, in absorbing heat, presents much in common to a
      physical change of state--such as, for example, that of a liquid
      into a gas. Deville likened complete decomposition to boiling, and
      compared partial decomposition, when a portion of a substance is
      not decomposed in the presence of its products of decomposition
      (or dissociation), to evaporation.

  [9] A reaction of rearrangement may in certain cases take place with
      one substance only; that is to say, a substance may by itself
      change into a new isomeric form. Thus, for example, if hard yellow
      sulphur be heated to a temperature of 250° and then poured into
      cold water it gives, on cooling, a soft, brown variety. Ordinary
      phosphorus, which is transparent, poisonous, and phosphorescent
      in the dark (in the air), gives, after being heated at 270° (in
      an atmosphere incapable of supporting combustion, such as steam),
      an opaque, red, and non-poisonous isomeric variety, which is not
      phosphorescent. Cases of isomerism point out the possibility of an
      internal rearrangement in a substance, and are the result of an
      alteration in the grouping of the same elements, just as a certain
      number of balls may be grouped in figures and forms of different
      shapes.

The majority of the chemical changes which occur in nature and
are made use of technically are very complicated, as they consist
of an association of many separate and simultaneous combinations,
decompositions, and replacements. It is chiefly due to this natural
complexity of chemical phenomena that for so many centuries chemistry
did not exist as an exact science; that is so say, that although many
chemical changes were known and made use of,[10] yet their real nature
was unknown, nor could they be predicted or directed at will. Another
reason for the tardy progress of chemical knowledge is the participation
of gaseous substances, especially air, in many reactions. The true
comprehension of air as a ponderable substance, and of gases in general
as peculiar elastic and dispersive states of matter, was only arrived
at in the sixteenth and seventeenth centuries, and it was only after
this that the transformations of substances could form a science. Up
to that time, without understanding the invisible and yet ponderable
gaseous and vaporous states of substances, it was impossible to obtain
any fundamental chemical evidence, because gases escaped from notice
between the reacting and resultant substances. It is easy from the
impression conveyed to us by the phenomena we observe to form the
opinion that matter is created and destroyed: a whole mass of trees
burn, and there only remains a little charcoal and ash, whilst from
one small seed there grows little by little a majestic tree. In one
case matter seems to be destroyed, and in the other to be created. This
conclusion is arrived at because the formation or consumption of gases,
being under the circumstances invisible to the eye, is not observed.
When wood burns it undergoes a chemical change into gaseous products,
which escape as smoke. A very simple experiment will prove this. By
collecting the smoke it may be observed that it contains gases which
differ entirely from air, being incapable of supporting combustion or
respiration. These gases may be weighed, and it will then be seen that
their weight exceeds that of the wood taken. This increase in weight
arises from the fact that, in burning, the component parts of the wood
combine with a portion of the air; in like manner iron increases in
weight by rusting. In burning gunpowder its substance is not destroyed,
but only converted into gases and smoke. So also in the growth of a
tree; the seed does not increase in mass of itself and from itself, but
grows because it absorbs gases from the atmosphere and sucks water and
substances dissolved therein from the earth through its roots. The sap
and solid substances which give plants their form are produced from
these absorbed gases and liquids by complicated chemical processes. The
gases and liquids are converted into solid substances by the plants
themselves. Plants not only do not increase in size, but die, in a gas
which does not contain the constituents of air. When moist substances
dry they decrease in weight; when water evaporates we know that it
does not disappear, but will return from the atmosphere as rain, dew,
and snow. When water is absorbed by the earth, it does not disappear
there for ever, but accumulates somewhere underground, from whence it
afterwards flows forth as a spring. Thus matter does not disappear
and is not created, but only undergoes various physical and chemical
transformations--that is to say, changes its locality and form. Matter
remains on the earth in the same quantity as before; in a word it is, so
far as we are concerned, everlasting. It was difficult to submit this
simple and primary truth of chemistry to investigation, but when once
made clear it rapidly spread, and now seems as natural and simple as
many truths which have been acknowledged for ages. Mariotte and other
savants of the seventeenth century already suspected the existence of
the law of the indestructibility of matter, but they made no efforts
to express it or to apply it to the requirements of science. The
experiments by means of which this simple law was arrived at were
made during the latter half of the last century by the founder of
modern chemistry, LAVOISIER, the French Academician and tax farmer. The
numerous experiments of this savant were conducted with the aid of the
balance, which is the only means of directly and accurately determining
the quantity of matter.

  [10] Thus the ancients knew how to convert the juice of grapes
       containing the saccharine principle (glucose) into wine or
       vinegar, how to extract metals from the ores which are found
       in the earth's crust, and how to prepare glass from earthy
       substances.

Lavoisier found, by weighing all the substances, and even the apparatus,
used in every experiment, and then weighing the substances obtained
after the chemical change, that the sum of the weights of the substances
formed was always equal to the sum of the weights of the substances
taken; or, in other words: MATTER IS NOT CREATED AND DOES NOT DISAPPEAR,
or that, _matter is everlasting_. This expression naturally includes
a hypothesis, but our only aim in using it is to concisely express
the following lengthy period--That in all experiments, and in all the
investigated phenomena of nature, it has never been observed that the
weight of the substances formed was less or greater (as far as accuracy
of weighing permits[11]) than the weight of the substances originally
taken, and as weight is proportional to mass[11 bis] or quantity of
matter, it follows that no one has ever succeeded in observing a
disappearance of matter or its appearance in fresh quantities. The law
of the indestructibility of matter endows all chemical investigations
with exactitude, as, on its basis, an equation may be formed for every
chemical reaction. If in any reaction the weights of the substances
taken be designated by the letters A, B, C, &c., and the weights of the
substances formed by the letters M, N, O, &c., then

            A + B + C + ... ... ... = M + N + O + ... ... ...

Therefore, should the weight of one of the re-acting or resultant
substances be unknown, it may be determined by solving the equation.
The chemist, in applying the law of the indestructibility of matter,
and in making use of the chemical balance, must never lose sight of any
one of the re-acting or resultant substances. Should such an over-sight
be made, it will at once be remarked that the sum of the weights of the
substances taken is unequal to the sum of the weights of the substances
formed. All the progress made by chemistry during the end of the last,
and in the present, century is entirely and immovably founded on the
law of the indestructibility of matter. It is absolutely necessary in
beginning the study of chemistry to become familiar with the simple
truth which is expressed by this law, and for this purpose several
examples elucidating its application will now be cited.

  [11] The experiments conducted by Staas (described in detail in Chap.
       XXIV. on Silver) form some of the accurate researches, proving
       that the weight of matter is not altered in chemical reactions,
       because he accurately weighed (introducing all the necessary
       corrections) the reacting and resultant substances. Landolt
       (1893) carried on various reactions in inverted and sealed glass
       U-tubes, and on weighing the tubes before reaction (when the
       reacting solutions were separated in each of the branches of
       the tubes), and after (when the solutions had been well mixed
       together by shaking), found that either the weight remained
       perfectly constant or that the variation was so small (for
       instance, 0·2 milligram in a total weight of about a million
       milligrams) as to be ascribed to the inevitable errors of
       weighing.

  [11 bis] The idea of the mass of matter was first shaped into an exact
       form by Galileo (died 1642), and more especially by Newton (born
       1643, died 1727), in the glorious epoch of the development of
       the principles of inductive reasoning enunciated by Bacon and
       Descartes in their philosophical treatises. Shortly after the
       death of Newton, Lavoisier, whose fame in natural philosophy
       should rank with that of Galileo and Newton, was born on August
       26, 1743. The death of Lavoisier occurred during the Reign
       of Terror of the French Revolution, when he, together with
       twenty-six other chief farmers of the revenue, was guillotined
       on May 8, 1794, at Paris; but his works and ideas have made him
       immortal.

1. It is well known that iron rusts in damp air,[12] and that when
heated to redness in air it becomes coated with scoria (oxide), having,
like rust, the appearance of an earthy substance resembling some of the
iron ores from which metallic iron is extracted. If the iron is weighed
before and after the formation of the scoria or rust, it will be found
that the metal has increased in weight during the operation.[13] It
can easily be proved that this increase in weight is accomplished at
the expense of the atmosphere, and mainly, as Lavoisier proved, at the
expense of that portion which is called oxygen. In fact, in a vacuum,
or in gases which do not contain oxygen, for instance, in hydrogen or
nitrogen, the iron neither rusts nor becomes coated with scoria. Had the
iron not been weighed, the participation of the oxygen of the atmosphere
in its transformation into an earthy substance might have easily passed
unnoticed, as was formerly the case, when phenomena like the above were,
for this reason, misunderstood. It is evident from the law of the
indestructibility of matter that as the iron increases in weight in its
conversion into rust, the latter must be a more complex substance than
the iron itself, and its formation is due to a reaction of combination.
We might form an entirely wrong opinion about it, and might, for
instance, consider rust to be a simpler substance than iron, and explain
the formation of rust as the removal of something from the iron. Such,
indeed, was the general opinion prior to Lavoisier, when it was held
that iron contained a certain unknown substance called 'phlogiston,' and
that rust was iron deprived of this supposed substance.

  [12] By covering iron with an enamel, or varnish, or with unrustable
       metals (such as nickel), or a coating of paraffin, or other
       similar substances, it is protected from the air and moisture,
       and so kept from rusting.

  [13] Such an experiment may easily be made by taking the finest
       (unrusted) iron filings (ordinary filings must be first washed
       in ether, dried, and passed through a very fine sieve). The
       filings thus obtained are capable of burning directly in air
       (by oxidising or forming rust), especially when they hang (are
       attracted) on a magnet. A compact piece of iron does not burn in
       air, but spongy iron glows and smoulders like tinder. In making
       the experiment, a horse-shoe magnet is fixed, with the poles
       downwards, on one arm of a rather sensitive balance, and the iron
       filings are applied to the magnet (on a sheet of paper) so as to
       form a beard about the poles. The balance pan should be exactly
       under the filings on the magnet, in order that any which might
       fall from it should not alter the weight. The filings, having
       been weighed, are set light to by applying the flame of a candle;
       they easily take fire, and go on burning by themselves, forming
       rust. When the combustion is ended, it will be clear that the
       iron has increased in weight; from 5-1/2 parts by weight of iron
       filings taken, there are obtained, by complete combustion, 7-1/2
       parts by weight of rust.

[Illustration: FIG. 1.--Apparatus for the decomposition of red mercury
oxide.]

2. Copper carbonate (in the form of a powder, or as the well-known green
mineral called 'malachite,' which is used for making ornaments, or as an
ore for the extraction of copper) changes into a black substance called
'copper oxide' when heated to redness.[14] This black substance is also
obtained by heating copper to redness in air--that is, it is the scoria
or oxidation product of copper. The weight of the black oxide of copper
left is less than that of the copper carbonate originally taken, and
therefore we consider the reaction which occurred to have been one of
decomposition, and that by it something was separated from the green
copper carbonate, and, in fact, by closing the orifice of the vessel in
which the copper carbonate is heated with a well-fitting cork, through
which a gas delivery tube[15] passes whose end is immersed under water,
it will be observed that on heating, a gas is formed which bubbles
through the water. This gas can be easily collected, as will presently
be described, and it will be found to essentially differ from air in
many respects; for instance, a burning taper is extinguished in it as if
it had been plunged into water. If weighing had not proved to us that
some substance had been separated, the formation of the gas might easily
have escaped our notice, for it is colourless and transparent like air,
and is therefore evolved without any striking feature. The carbonic
anhydride evolved may be weighed,[16] and it will be seen that the sum
of the weights of the black copper oxide and carbonic anhydride is equal
to the weight of the copper carbonate[17] originally taken, and thus by
carefully following out the various stages of all chemical reactions we
arrive at a confirmation of the law of the indestructibility of matter.

  [14] For the purpose of experiment, it is most convenient to take
       copper carbonate, which may be prepared by the experimenter
       himself, by adding a solution of sodium carbonate to a solution
       of copper sulphate. The precipitate (deposit) so formed is
       collected on a filter, washed, and dried. The decomposition of
       copper carbonate into copper oxide is effected by so moderate
       a heat that it may be performed in a glass vessel heated by a
       lamp. For this purpose a thin glass tube, closed at one end, and
       called a 'test tube,' may be employed, or else a vessel called a
       'retort.' The experiment is carried on, as described in example
       three on p. 11, by collecting the carbonic anhydride over water,
       as will be afterwards explained.

  [15] Gas delivery tubes are usually made of glass tubing of various
       diameters and thicknesses. If of small diameter and thickness, a
       glass tube is easily bent by heating in a gas jet or the flame
       of a spirit lamp, and it may also be easily divided at a given
       point by making a deep scratch with a file and then breaking the
       tube at this point with a sharp jerk. These properties, together
       with their impermeability, transparency, hardness, and regularity
       of bore, render glass tubes most useful in experiments with
       gases. Naturally they might be replaced by straws, india-rubber,
       metallic, or other tubes, but these are more difficult to fix on
       to a vessel, and are not entirely impervious to gases. A glass
       gas delivery tube may be hermetically fixed into a vessel by
       fitting it into a perforated cork, which should be soft and free
       from flaws, and fixing the cork into the orifice of the vessel.
       To protect the cork from the action of gases it is sometimes
       previously soaked in paraffin, or it may be replaced by an
       india-rubber cork.

  [16] Gases, like all other substances, may be weighed, but, owing to
       their extreme lightness and the difficulty of dealing with them
       in large masses, they can only be weighed by very sensitive
       balances; that is, in such as, with a considerable load, indicate
       a very small difference in weight--for example, a centigram or a
       milligram with a load of 1,000 grams. In order to weigh a gas,
       a glass globe furnished with a tight-fitting stop-cock is first
       of all exhausted of air by an air-pump (a Sprengel pump is the
       best). The stop-cock is then closed, and the exhausted globe
       weighed. If the gas to be weighed is then let into the globe,
       its weight can be determined from the increase in the weight of
       the globe. It is necessary, however, that the temperature and
       pressure of the air about the balance should remain constant for
       both weighings, as the weight of the globe in air will (according
       to the laws of hydrostatics) vary with its density. The volume of
       the air displaced, and its weight, must therefore be determined
       by observing the temperature, density, and moisture of the
       atmosphere during the time of experiment. This will be partly
       explained later, but may be studied more in detail by physics.
       Owing to the complexity of all these operations, the mass of a
       gas is usually determined from its volume and density, or from
       the weight of a known volume.

  [17] The copper carbonate should be dried before weighing, as
       otherwise--besides copper oxide and carbonic anhydride--water
       will be obtained in the decomposition. Water forms a part of the
       composition of malachite, and has therefore to be taken into
       consideration. The water produced in the decomposition may be all
       collected by absorbing it in sulphuric acid or calcium chloride,
       as will be described further on. In order to dry a salt it must
       be heated at about 100° until its weight remains constant, or
       be placed under an air pump over sulphuric acid, as will also
       be presently described. As water is met with almost everywhere,
       and as it is absorbed by many substances, the possibility of its
       presence should never be lost sight of.

3. Red mercury oxide (which is formed as mercury rust by heating
mercury in air) is decomposed like copper carbonate (only by heating
more slowly and at a somewhat higher temperature), with the formation
of the peculiar gas, oxygen. For this purpose the mercury oxide is
placed in a glass tube or retort,[18] to which a gas delivery tube is
attached by means of a cork. This tube is bent downwards, as shown in
the drawing (Fig. 1). The open end of the gas delivery tube is immersed
in a vessel filled with water, called a pneumatic trough.[19] When the
gas begins to be evolved in the retort it is obliged, having no other
outlet, to escape through the gas delivery tube into the water in the
pneumatic trough, and therefore its evolution will be rendered visible
by the bubbles coming from this tube. In heating the retort containing
the mercury oxide, the air contained in the apparatus is first partly
expelled, owing to its expansion by heat, and then the peculiar gas
called 'oxygen' is evolved, and may be easily collected as it comes off.
For this purpose a vessel (an ordinary cylinder, as in the drawing) is
filled quite full with water and its mouth closed; it is then inverted
and placed in this position under the water in the trough; the mouth
is then opened. The cylinder will remain full of water--that is, the
water will remain at a higher level in it than in the surrounding
vessel, owing to the atmospheric pressure. The atmosphere presses on
the surface of the water in the trough, and prevents the water from
flowing out of the cylinder. The mouth of the cylinder is placed over
the end of the gas delivery tube,[20] and the bubbles issuing from it
will rise into the cylinder and displace the water contained in it.
Gases are generally collected in this manner. When a sufficient quantity
of gas has accumulated in the cylinder it can be clearly shown that it
is not air, but another gas which is distinguished by its capacity for
vigorously supporting combustion. In order to show this, the cylinder is
closed, under water, and removed from the bath; its mouth is then turned
upwards, and a smouldering taper plunged into it. As is well known, a
smouldering taper will be extinguished in air, but in the gas which
is given off from red mercury oxide it burns clearly and vigorously,
showing the property possessed by this gas for supporting combustion
more energetically than air, and thus enabling it to be distinguished
from the latter. It may be observed in this experiment that, besides the
formation of oxygen, metallic mercury is formed, which, volatilising at
the high temperature required for the reaction, condenses on the cooler
parts of the retort as a mirror or in globules. Thus two substances,
mercury and oxygen, are obtained by heating red mercury oxide. In this
reaction, from one substance, two new substances are produced--that
is, a decomposition has taken place. The means of collecting and
investigating gases were known before Lavoisier's time, but he first
showed the real part they played in the processes of many chemical
changes which before his era were either wrongly understood (as will be
afterwards explained) or were not explained at all, but only observed
in their superficial aspects. This experiment on red mercury oxide has
a special significance in the history of chemistry contemporary with
Lavoisier, because the oxygen gas which is here evolved is contained in
the atmosphere, and plays a most important part in nature, especially in
the respiration of animals, in combustion in air, and in the formation
of rusts or scoriæ (earths, as they were then called) from metals--that
is, of earthy substances, like the ores from which metals are extracted.

  [18] As the decomposition of red oxide of mercury requires so high a
       temperature, near redness, as to soften ordinary glass, it is
       necessary for this experiment to take a retort (or test tube)
       made of hard glass, which is able to stand high temperatures
       without softening. For the same reason, the lamp used must give
       a strong heat and a large flame, capable of embracing the whole
       bottom of the retort, which should be as small as possible for
       the convenience of the experiment.

  [19] [Illustration: FIG. 2.--Apparatus for distilling under a
       diminished pressure liquids which decompose at their boiling
       points under the ordinary pressure. The apparatus in which the
       liquid is distilled is connected with a large globe from which the
       air is pumped out; the liquid is heated, and the receiver cooled.]

       The pneumatic trough may naturally be made of any material
       (china, earthenware, or metal, &c.), but usually a glass one,
       as shown in the drawing, is used, as it allows the progress
       of the experiment to be better observed. For this reason, as
       well as the ease with which they are kept clean, and from the
       fact also that glass is not acted on by many substances which
       affect other materials (for instance, metals), glass vessels
       of all kinds--such as retorts, test tubes, cylinders, beakers,
       flasks, globes, &c.--are preferred to any other for chemical
       experiments. Glass vessels may be heated without any danger if
       the following precautions be observed: 1st, they should be made
       of thin glass, as otherwise they are liable to crack from the bad
       heat-conducting power of glass; 2nd, they should be surrounded
       by a liquid or with sand (Fig. 2), or sand bath as it is called;
       or else should stand in a current of hot gases without touching
       the fuel from which they proceed, or in the flame of a smokeless
       lamp. A common candle or lamp forms a deposit of soot on a
       cold object placed in their flames. The soot interferes with
       the transmission of heat, and so a glass vessel when covered
       with soot often cracks. And for this reason spirit lamps, which
       burn with a smokeless flame, or gas burners of a peculiar
       construction, are used. In the Bunsen burner the gas is mixed
       with air, and burns with a non-luminous and smokeless flame.
       On the other hand, if an ordinary lamp (petroleum or benzine)
       does not smoke it may be used for heating a glass vessel without
       danger, provided the glass is placed well above the flame in the
       current of hot gases. In all cases, the heating should be begun
       very carefully by raising the temperature by degrees.


  [20] In order to avoid the necessity of holding the cylinder, its open
       end is widened (and also ground so that it may be closely covered
       with a ground-glass plate when necessary), and placed on a stand
       below the level of the water in the bath. This stand is called
       'the bridge.' It has several circular openings cut through it,
       and the gas delivery tube is placed under one of these, and the
       cylinder for collecting the gas over it.

4. In order to illustrate by experiment one more example of chemical
change and the application of the law of the indestructibility of
matter, we will consider the reaction between common table salt and
lunar caustic, which is well known from its use in cauterising wounds.
By taking a clear solution of each and mixing them together, it will at
once be observed that a solid white substance is formed, which settles
to the bottom of the vessel, and is insoluble in water. This substance
may be separated from the solution by filtering; it is then found to be
an entirely different substance from either of those taken originally in
the solutions. This is at once evident from the fact that it does not
dissolve in water. On evaporating the liquid which passed through the
filter, it will be found to contain a new substance unlike either table
salt or lunar caustic, but, like them, soluble in water. Thus table salt
and lunar caustic, two substances soluble in water, produced, by their
mutual chemical action, two new substances, one insoluble in water, and
the other remaining in solution. Here, from two substances, two others
are obtained, consequently there occurred a reaction of substitution.
The water served only to convert the re-acting substances into a liquid
and mobile state. If the lunar caustic and salt be dried[21] and
weighed, and if about 58-1/2 grams[22] of salt and 170 grams of lunar
caustic be taken, then 143-1/2 grams of insoluble silver chloride and
85 grams of sodium nitrate will be obtained. The sum of the weights
of the re-acting and resultant substances are seen to be similar and
equal to 228-1/2 grams, which necessarily follows from the law of the
indestructibility of matter.

  [21] Drying is necessary in order to remove any water which may be held
       in the salts (_see_ Note 17, and Chapter I., Notes 13 and 14).

  [22] The exact weights of the re-acting and resulting substances are
       determined with the greatest difficulty, not only from the
       possible inexactitude of the balance (every weighing is only
       correct within the limits of the sensitiveness of the balance)
       and weights used in weighing, not only from the difficulty in
       making corrections for the weight of air displaced by the vessels
       holding the substances weighed and by the weights themselves, but
       also from the hygroscopic nature of many substances (and vessels)
       causing absorption of moisture from the atmosphere, and from the
       difficulty in not losing any of the substance to be weighed in
       the several operations (filtering, evaporating, and drying, &c.)
       which have to be performed before arriving at a final result. All
       these circumstances have to be taken into consideration in exact
       researches, and their elimination requires very many special
       precautions which are impracticable in preliminary experiments.

Accepting the truth of the above law, the question naturally arises as
to whether there is any limit to the various chemical transformations,
or are they unrestricted in number--that is to say, is it possible
from a given substance to obtain an equivalent quantity of any other
substance? In other words, does there exist a perpetual and infinite
change of one kind of material into every other kind, or is the cycle
of these transformations limited? This is the second essential problem
of Chemistry, a question of quality of matter, and one, it is evident,
which is more complicated than the question of quantity. It cannot be
solved by a mere superficial glance at the subject. Indeed, on seeing
how all the varied forms and colours of plants are built up from air
and the elements of the soil, and how metallic iron can be transformed
into colours such as inks and Prussian blue, we might be led to think
that there is no end to the qualitative changes to which matter is
susceptible. But, on the other hand, the experiences of everyday life
compel us to acknowledge that food cannot be made out of a stone, or
gold out of copper. Thus a definite answer can only be looked for in
a close and diligent study of the subject, and the problem has been
resolved in different way at different times. In ancient times the
opinion most generally held was that everything visible was composed of
four elements--Air, Water, Earth, and Fire. The origin of this doctrine
can be traced far back into the confines of Asia, whence it was handed
down to the Greeks, and most fully expounded by Empedocles, who lived
before 460 B.C. This doctrine was not the result of exact research,
but apparently owes its origin to the clear division of bodies into
gases (like air), liquids (like water), and solids (like the earth).
The Arabs appear to have been the first who attempted to solve the
question by experimental methods, and they introduced, through Spain,
the taste for the study of similar problems into Europe, where from
that time there appear many adepts in chemistry, which was considered
as an unholy art, and called 'alchemy.' As the alchemists were ignorant
of any exact law which could guide them in their researches, they
obtained most anomalous results. Their chief service to chemistry
was that they made a number of experiments, and discovered many new
chemical transformations; but it is well known how they solved the
fundamental problem of chemistry. Their view may be taken as a positive
acknowledgment of the infinite transmutability of matter, for they aimed
at discovering the Philosopher's Stone, capable of converting everything
into gold and diamonds, and of making the old young again. This solution
of the question was afterwards completely overthrown, but it must not,
for this reason, be thought that the hopes held by the alchemists were
only the fruit of their imaginations. The first chemical experiments
might well lead them to their conclusions. They took, for instance,
the bright metallic mineral galena, and extracted metallic lead from
it. Thus they saw that from a metallic substance which is unfitted for
use they could obtain another metallic substance which is ductile and
valuable for many technical purposes. Furthermore, they took this lead
and obtained silver, a still more valuable metal, from it. Thus they
might easily conclude that it was possible to ennoble metals by means of
a whole series of transmutations--that is to say, to obtain from them
those which are more and more precious. Having got silver from lead,
they assumed that it would be possible to obtain gold from silver. The
mistake they made was that they never weighed or measured the substances
used or produced in their experiments. Had they done so, they would have
learnt that the weight of the lead was much less than that of the galena
from which it was obtained, and the weight of the silver infinitesimal
compared with that of the lead. Had they looked more closely into
the process of the extraction of the silver from lead (and silver at
the present time is chiefly obtained from the lead ores) they would
have seen that the lead does not change into silver, but that it only
contains a certain small quantity of it, and this amount having once
been separated from the lead it cannot by any further operation give
more. The silver which the alchemists extracted from the lead was in the
lead, and was not obtained by a chemical change of the lead itself. This
is now well known from experiment, but the first view of the nature of
the process was very likely to be an erroneous one.[23] The methods of
research adopted by the alchemists could give but little success, for
they did not set themselves clear and simple questions whose answers
would aid them to make further progress. Thus though they did not
arrive at any exact law, they left nevertheless numerous and useful
experimental data as an inheritance to chemistry; they investigated, in
particular, the transformations proper to metals, and for this reason
chemistry was for long afterwards entirely confined to the study of
metallic substances.

  [23] Besides which, in the majority of cases, the first explanation of
       most subjects which do not repeat themselves in everyday
       experience under various aspects, but always in one form, or
       only at intervals and infrequently, is usually wrong. Thus
       the daily evidence of the rising of the sun and stars evokes
       the erroneous idea that the heavens move and the earth stands
       still. This apparent truth is far from being the real truth,
       and, as a matter of fact, is contradictory to it. Similarly, an
       ordinary mind and everyday experience concludes that iron is
       incombustible, whereas it burns not only as filings, but even as
       wire, as we shall afterwards see. With the progress of knowledge
       very many primitive prejudices have been obliged to give way to
       true ideas which have been verified by experiment. In ordinary
       life we often reason at first sight with perfect truth, only
       because we are taught a right judgment by our daily experience.
       It is a necessary consequence of the nature of our minds to reach
       the attainment of truth through elementary and often erroneous
       reasoning and through experiment, and it would be very wrong
       to expect a knowledge of truth from a simple mental effort.
       Naturally, experiment itself cannot give truth, but it gives the
       means of destroying erroneous representations whilst confirming
       those which are true in all their consequences.

In their researches, the alchemists frequently made use of two chemical
processes which are now termed 'reduction' and 'oxidation.' The rusting
of metals, and in general their conversion from a metallic into an
earthy form, is called 'oxidation,' whilst the extraction of a metal
from an earthy substance is called 'reduction.' Many metals--for
instance, iron, lead, and tin--are oxidised by heating in air alone,
and may be again reduced by heating with carbon. Such oxidised metals
are found in the earth, and form the majority of metallic ores. The
metals, such as tin, iron, and copper, may be extracted from these ores
by heating them together with carbon. All these processes were well
studied by the alchemists. It was afterwards shown that all earths and
minerals are formed of similar metallic rusts or oxides, or of their
combinations. Thus the alchemists knew of two forms of chemical changes:
the oxidation of metals and the reduction of the oxides so formed
into metals. The explanation of the nature of these two classes of
chemical phenomena was the means for the discovery of the most important
chemical laws. The first hypothesis on their nature is due to Becker,
and more particularly to Stahl, a surgeon to the King of Prussia.
Stahl writes in his 'Fundamenta Chymiæ,' 1723, that all substances
consist of an imponderable fiery substance called 'phlogiston' (materia
aut principium ignis non ipse ignis), and of another element having
particular properties for each substance. The greater the capacity of
a body for oxidation, or the more combustible it is, the richer it is
in phlogiston. Carbon contains it in great abundance. In oxidation or
combustion phlogiston is emitted, and in reduction it is consumed or
enters into combination. Carbon reduces earthy substances because it is
rich in phlogiston, and gives up a portion of its phlogiston to the
substance reduced. Thus Stahl supposed metals to be compound substances
consisting of phlogiston and an earthy substance or oxide. This
hypothesis is distinguished for its very great simplicity, and for this
and other reasons it acquired many supporters.[24]

  [24] It is true that Stahl was acquainted with a fact which directly
       disproved his hypothesis. It was already known (from the
       experiments of Geber, and more especially of Ray, in 1630) that
       metals increase in weight by oxidation, whilst, according to
       Stahl's hypothesis, they should decrease in weight, because
       phlogiston is separated by oxidation. Stahl speaks on this
       point as follows:--'I am well aware that metals, in their
       transformation into earths, increase in weight. But not only does
       this fact not disprove my theory, but, on the contrary, confirms
       it, for phlogiston is lighter than air, and, in combining
       with substances, strives to lift them, and so decreases their
       weight; consequently, a substance which has lost phlogiston
       must be heavier.' This argument, it will be seen, is founded
       on a misconception of the properties of gases, regarding them
       as having no weight and as not being attracted by the earth,
       or else on a confused idea of phlogiston itself, since it was
       first defined as imponderable. The conception of imponderable
       phlogiston tallies well with the habit and methods of the last
       century, when recourse was often had to imponderable fluids for
       explaining a large number of phenomena. Heat, light, magnetism,
       and electricity were explained as being peculiar imponderable
       fluids. In this sense the doctrine of Stahl corresponds entirely
       with the spirit of his age. If heat be now regarded as motion
       or energy, then phlogiston also should be considered in this
       light. In fact, in combustion, of coals for instance, heat and
       energy are evolved, and not combined in the coal, although the
       oxygen and coal do combine. Consequently, the doctrine of Stahl
       contains the essence of a true representation of the evolution
       of energy, but naturally this evolution is only a consequence of
       the combination occurring between the coal and oxygen. As regards
       the history of chemistry prior to Lavoisier, besides Stahl's
       work (to which reference has been made above), Priestley's
       _Experiments and Observations on Different Kinds of Air_, London,
       1790, and also Scheele's _Opuscula Chimica et Physica_, Lips.,
       1788-89, 2 vols., must be recommended as the two leading works
       of the English and Scandinavian chemists showing the condition
       of chemical learning before the propagation of Lavoisier's
       views, and containing also many important observations which lie
       at the basis of the chemistry of our times. A most interesting
       memoir on the history of phlogiston is that of Rodwell, in the
       _Philosophical Magazine_, 1868, in which it is shown that the
       idea of phlogiston dates very far back, that Basil Valentine
       (1394-1415), in the _Cursus Triumphalis Antimonii_, Paracelsus
       (1493-1541), in his work, _De Rerum Natura_, Glauber (1604-1668),
       and especially John Joachim Becher (1625-1682), in his _Physica
       Subterranea_, all referred to phlogiston, but under different
       names.

[Illustration: FIG. 3.--Lavoisier's apparatus for determining the
composition of air and the reason of metals increasing in weight when
they are calcined in air.]

Lavoisier proved by means of the balance that every case of rusting of
metals or oxidation, or of combustion, is accompanied by an increase
in weight at the expense of the atmosphere. He formed, therefore, the
natural opinion that the heavier substance is more complex than the
lighter one.[25] Lavoisier's celebrated experiment, made in 1774,
gave indubitable support to his opinion, which in many respects was
contradictory to Stahl's doctrine. Lavoisier poured four ounces of pure
mercury into a glass retort (fig. 3), whose neck was bent as shown in
the drawing and dipped into the vessel R S, also full of mercury. The
projecting end of the neck was covered with a glass bell-jar P. The
weight of all the mercury taken, and the volume of air remaining in
the apparatus, namely, that in the upper portion of the retort, and
under the bell-jar, were determined before beginning the experiment.
It was most important in this experiment to know the volume of air in
order to learn what part it played in the oxidation of the mercury,
because, according to Stahl, phlogiston is emitted into the air, whilst,
according to Lavoisier, the mercury in oxidising absorbs a portion of
the air; and consequently it was absolutely necessary to determine
whether the amount of air increased or decreased in the oxidation of
the metal. It was, therefore, most important to measure the volume of
the air in the apparatus both before and after the experiment. For this
purpose it was necessary to know the total capacity of the retort,
the volume of the mercury poured into it, the volume of the bell-jar
above the level of the mercury, and also the temperature and pressure
of the air at the time of its measurement. The volume of air contained
in the apparatus and isolated from the surrounding atmosphere could
be determined from these data. Having arranged his apparatus in this
manner, Lavoisier heated the retort holding the mercury for a period
of twelve days at a temperature near the boiling point of mercury. The
mercury became covered with a quantity of small red scales; that is,
it was oxidised or converted into an earth. This substance is the same
mercury oxide which has already been mentioned (example 3). After the
lapse of twelve days the apparatus was cooled, and it was then seen that
the volume of the air in the apparatus had diminished during the time
of the experiment. This result was in exact contradiction to Stahl's
hypothesis. Out of 50 cubic inches of air originally taken, there only
remained 42. Lavoisier's experiment led to other equally important
results. The weight of the air taken decreased by as much as the weight
of the mercury increased in oxidising; that is, the portion of the air
was not destroyed, but only combined with mercury. This portion of
the air may be again separated from the mercury oxide and has, as we
saw (example 3), properties different from those of air. It is called
'oxygen.' That portion of the air which remained in the apparatus and
did not combine with the mercury does not oxidise metals, and cannot
support either combustion or respiration, so that a lighted taper is
immediately extinguished if it be dipped into the gas which remains
in the bell-jar. 'It is extinguished in the residual gas as if it had
been plunged into water,' writes Lavoisier in his memoirs. This gas is
called 'nitrogen.' Thus air is not a simple substance, but consists of
two gases, oxygen and nitrogen, and therefore the opinion that air is
an elementary substance is erroneous. The oxygen of the air is absorbed
in combustion and the oxidation of metals, and the earths produced by
the oxidation of metals are substances composed of oxygen and a metal.
By mixing the oxygen with the nitrogen the same air as was originally
taken is re-formed. It has also been shown by direct experiment that
on reducing an oxide with carbon, the oxygen contained in the oxide is
transferred to the carbon, and gives the same gas that is obtained by
the combustion of carbon in air. Therefore this gas is a compound of
carbon and oxygen, just as the earthy oxides are composed of metals and
oxygen.

  [25] An Englishman, named Mayow, who lived a whole century before
       Lavoisier (in 1666), understood certain phenomena of oxidation
       in their true aspect, but was not able to develop his views
       with clearness, or support them by conclusive experiments; he
       cannot therefore be considered, like Lavoisier, as the founder of
       contemporary chemical learning. Science is a universal heritage,
       and therefore it is only just to give the highest honour in
       science, not to those who first enunciate a certain truth, but to
       those who are first able to convince others of its authenticity
       and establish it for the general welfare. But scientific
       discoveries are rarely made all at once; as a rule, the first
       teachers do not succeed in convincing others of the truth they
       have discovered; with time, however, a true herald comes forward,
       possessing every means for making the truth apparent to all, but
       it must not be forgotten that such are entirely indebted to the
       labours and mass of data accumulated by many others. Such was
       Lavoisier, and such are all the great founders of science. They
       are the enunciators of all past and present learning, and their
       names will always be revered by posterity.

The many examples of the formation and decomposition of substances which
are met with convince us that the majority of substances with which
we have to deal are compounds made up of several other substances. By
heating chalk (or else copper carbonate, as in the second example)
we obtain lime and the same carbonic acid gas which is produced by
the combustion of carbon. On bringing lime into contact with this gas
and water, at the ordinary temperature, we again obtain the compound,
carbonate of lime, or chalk. Therefore chalk is a compound. So also are
those substances from which it may be built up. Carbonic anhydride is
formed by the combination of carbon and oxygen; and lime is produced
by the oxidation of a certain metal called 'calcium.' By resolving
substances in this manner into their component parts, we arrive at
last at such as are indivisible into two or more substances by any
means whatever, and which cannot be formed from other substances. All
we can do is to make such substances combine together to act on other
substances. Substances which cannot be formed from or decomposed into
others are termed _simple substances_ (elements). Thus all homogeneous
substances may be classified into simple and compound substances. This
view was introduced and established as a scientific fact during the
lifetime of Lavoisier. The number of these elements is very small in
comparison with the number of compound substances which are formed
by them. At the present time, only seventy elements are known with
certainty to exist. Some of them are very rarely met with in nature, or
are found in very small quantities, whilst the existence of others is
still doubtful. The number of elements with whose compounds we commonly
deal in everyday life is very small. Elements cannot be transmuted
into one another--at least up to the present not a single case of such
a transformation has been met with; it may therefore be said that, as
yet, it is impossible to transmute one metal into another. And as yet,
notwithstanding the number of attempts which have been made in this
direction, no fact has been discovered which could in any way support
the idea of the complexity of such well-known elements[26] as oxygen,
iron, sulphur, &c. Therefore, from its very conception, an element is
not susceptible to reactions of decomposition.[27]

  [26] Many of the ancient philosophers assumed the existence of one
       elementary form of matter. This idea still appears in our times,
       in the constant efforts which are made to reduce the number of
       the elements; to prove, for instance, that bromine contains
       chlorine or that chlorine contains oxygen. Many methods, founded
       both on experiment and theory, have been tried to prove the
       compound nature of the elements. All labour in this direction has
       as yet been in vain, and the assurance that elementary matter
       is not so homogeneous (single) as the mind would desire in its
       first transport of rapid generalisation is strengthened from year
       to year. All our knowledge shows that iron and other elements
       remain, even at such a high temperature as there exists in the
       sun, as different substances, and are not converted into one
       common material. Admitting, even mentally, the possibility of one
       elementary form of matter, a method must be imagined by which
       it could give rise to the various elements, as also the _modus
       operandi_ of their formation from one material. If it be said
       that this diversitude only takes place at low temperatures, as is
       observed with isomerides, then there would be reason to expect,
       if not the transition of the various elements into one particular
       and more stable form, at least the mutual transformation of some
       into others. But nothing of the kind has as yet been observed,
       and the alchemist's hope to manufacture (as Berthollet puts it)
       elements has no theoretical or practical foundation.

  [27] The weakest point in the idea of elements is the negative
       character of the determinative signs given them by Lavoisier, and
       from that time ruling in chemistry. They do _not_ decompose, they
       do _not_ change into one another. But it must be remarked that
       elements form the limiting horizon of our knowledge of matter,
       and it is always difficult to determine a positive side on the
       borderland of what is known. Besides, there is no doubt (from the
       results of spectrum analysis) that the elements are distributed
       as far as the most distant stars, and that they support the
       highest attainable temperatures without decomposing.

The quantity, therefore, of each element remains constant in all
chemical changes: a fact which may be deduced as a consequence of
the law of the indestructibility of matter, and of the conception
of elements themselves. Thus the equation expressing the law of the
indestructibility of matter acquires a new and still more important
signification. If we know the quantities of the elements which occur in
the re-acting substances, and if from these substances there proceed, by
means of chemical changes, a series of new compound substances, then the
latter will together contain the same quantity of each of the elements
as there originally existed in the re-acting substances. The essence
of chemical change is embraced in the study of how, and with what
substances, each element is combined before and after change.

In order to be able to express various chemical changes by equations,
it has been agreed to represent each element by the first or some two
letters of its (Latin) name. Thus, for example, oxygen is represented by
the letter O; nitrogen by N; mercury (hydrargyrum) by Hg; iron (ferrum)
by Fe; and so on for all the elements, as is seen in the tables on
page 24. A compound substance is represented by placing the symbols
representing the elements of which it is made up side by side. For
example, red mercury oxide is represented by HgO, which shows that it
is composed of oxygen and mercury. Besides this, the symbol of every
element corresponds with a certain relative quantity of it by weight,
called its 'combining' weight, or the weight of an atom; so that the
chemical formula of a compound substance not only designates the nature
of the elements of which it is composed, but also their quantitative
proportion. Every chemical process may be expressed by an equation
composed of the formulæ corresponding with those substances which take
part in it and are produced by it. The amount by weight of the elements
in every chemical equation must be equal on both sides of the equation,
since no element is either formed or destroyed in a chemical change.

On pages 24, 25, and 26 a list of the elements, with their symbols and
combining or atomic weights, is given, and we shall see afterwards on
what basis the atomic weights of elements are determined. At present we
will only point out that a compound containing the elements A and B is
designated by the formula A_n_ B_m_, where _m_ and _n_ are the
coefficients or multiples in which the combining weights of the elements
enter into the composition of the substance. If we represent the
combining weight of the substance A by _a_ and that of the substance B by
_b_, then the composition of the substance A_n_ B_m_ will be expressed
thus: it contains _na_ parts by weight of the substance A and _mb_ parts
by weight of the substance B, and consequently 100 parts of our compound
contain _na_ 100/_na_ + _mb_ percentage parts by weight of the substance
A and _mb_ 100/_na_ + _mb_ of the substance B. It is evident that as a
formula shows the relative amounts of all the elements contained in a
compound, the actual weights of the elements contained in a given weight
of a compound may be calculated from its formula. For example, the
formula NaCl of table salt shows (as Na = 23 and Cl = 35·5) that 58·5
lbs. of salt contain 23 lbs. of sodium and 35·5 lbs. of chlorine, and
that 100 parts of it contain 39·3 per cent. of sodium and 60·7 per cent.
of chlorine.

What has been said above clearly limits the province of chemical
changes, because from substances of a given kind there can be obtained
only such as contain the same elements. Even with this limitation,
however, the number of possible combinations is infinitely great. Only
a comparatively small number of compounds have yet been described or
subjected to research, and any one working in this direction may easily
discover new compounds which had not before been obtained. It often
happens, however, that such newly-discovered compounds were foreseen by
chemistry, whose object is the apprehension of that uniformity which
rules over the multitude of compound substances, and whose aim is the
comprehension of those laws which govern their formation and properties.
The conception of elements having been established, the next objects
of chemistry were: the determination of the properties of compound
substances on the basis of the determination of the quantity and kind of
elements of which they are composed; the investigation of the elements
themselves; the determination of what compound substances can be formed
from each element and the properties which these compounds show; and
the apprehension of the nature of the connection between the elements
in different compounds. An element thus serves as the starting point,
and is taken as the primary conception on which all other substances are
built up.

When we state that a certain element enters into the composition of a
given compound (when we say, for instance, that mercury oxide contains
oxygen) we do not mean that it contains oxygen as a gaseous substance,
but only desire to express those transformations which mercury oxide
is capable of making; that is, we wish to say that it is possible to
obtain oxygen from mercury oxide, and that it can give up oxygen to
various other substances; in a word, we desire only to express those
transformations of which mercury oxide is capable. Or, more concisely,
it may be said that the _composition_ of a compound is the expression
of those transformations of which it is capable. It is useful in this
sense to make a clear distinction between the conception of an element
as a _separate_ homogeneous substance, and as a _material_ but invisible
_part_ of a compound. Mercury oxide does not contain two simple bodies,
a gas and a metal, but two elements, mercury and oxygen, which, when
free, are a gas and a metal. Neither mercury as a metal nor oxygen as
a gas is contained in mercury oxide; it only contains the substance of
these elements, just as steam only contains the substance of ice, but
not ice itself, or as corn contains the substance of the seed, but not
the seed itself. The existence of an element may be recognised without
knowing it in the uncombined state, but only from an investigation
of its combinations, and from the knowledge that it gives, under
all possible conditions, substances which are unlike other known
combinations of substances. Fluorine is an example of this kind. It was
for a long time unknown in a free state, and nevertheless was recognised
as an element because its combinations with other elements were known,
and their difference from all other similar compound substances was
determined. In order to grasp the difference between the conception of
the visible form of an element as we know it in the free state, and of
the intrinsic element (or 'radicle,' as Lavoisier called it) contained
in the visible form, it should be remarked that compound substances
also combine together forming yet more complex compounds, and that
they evolve heat in the process of combination. The original compound
may often be extracted from these new compounds by exactly the same
methods as elements are extracted from their corresponding combinations.
Besides, many elements exist under various visible forms whilst the
intrinsic element contained in these various forms is something which is
not subject to change. Thus carbon appears as charcoal, graphite, and
diamond, but yet the element carbon alone, contained in each, is one
and the same. Carbonic anhydride contains carbon, and not charcoal, or
graphite, or the diamond.

Elements alone, although not all of them, have the peculiar lustre,
opacity, malleability, and the great heat and electrical conductivity
which are proper to metals and their mutual combinations. But elements
are far from all being _metals_. Those which do not possess the physical
properties of metals are called _non-metals_ (or _metalloids_). It
is, however, impossible to draw a strict line of demarcation between
metals and non-metals, there being many intermediary substances. Thus
graphite, from which pencils are manufactured, is an element with the
lustre and other properties of a metal; but charcoal and the diamond,
which are composed of the same substance as graphite, do not show any
metallic properties. Both classes of elements are clearly distinguished
in definite examples, but in particular cases the distinction is not
clear and cannot serve as a basis for the exact division of the elements
into two groups.

The conception of elements forms the basis of chemical knowledge, and
in giving a list of them at the very beginning of our work, we wish to
tabulate our present knowledge on the subject. Altogether about seventy
elements are now authentically known, but many of them are so rarely met
with in nature, and have been obtained in such small quantities, that
we possess but a very insufficient knowledge of them. The substances
most widely distributed in nature contain a very small number of
elements. These elements have been more completely studied than the
others, because a greater number of investigators have been able to
carry on experiments and observations on them. The elements most widely
distributed in nature are:--

  Hydrogen,   H  = 1.      In water, and in animal and vegetable
                             organisms.
  Carbon,     C  = 12.     In organisms, coal, limestones.
  Nitrogen,   N  = 14.     In air and in organisms.
  Oxygen,     O  = 16.     In air, water, earth. It forms the greater
                             part of the mass of the earth.
  Sodium,     Na = 23.     In common salt and in many minerals.
  Magnesium,  Mg = 24.     In sea-water and in many minerals.
  Aluminium,  Al = 27.     In minerals and clay.
  Silicon,    Si = 28.     In sand, minerals, and clay.
  Phosphorus, P  = 31.     In bones, ashes of plants, and soil.
  Sulphur,    S  = 32.     In pyrites, gypsum, and in sea-water.
  Chlorine,   Cl = 35·5.   In common salt, and in the salts of sea-water.
  Potassium,  K  = 39.     In minerals, ashes of plants, and in nitre.
  Calcium,    Ca = 40.     In limestones, gypsum, and in organisms.
  Iron,       Fe = 56.     In the earth, iron ores, and in organisms.

Besides these, the following elements, although not very largely
distributed in nature, are all more or less well known from their
applications to the requirements of everyday life or the arts, either in
a free state or in their compounds:--

  Lithium,    Li = 7.     In medicine (Li_{2}CO_{3}), and in photography
                            (LiBr).
  Boron,      B = 11.     As borax, B_{4}Na_{2}O_{7}, and as boric
                            anhydride, B_{2}O_{3}.
  Fluorine,   F  = 19.    As fluor spar, CaF_{2}, and as hydrofluoric
                            acid, HF.
  Chromium,   Cr = 52.    As chromic anhydride, CrO_{3}, and potassium
                            dichromate, K_{2}Cr_{2}O_{7}.
  Manganese,  Mn = 55.    As manganese peroxide, MnO_{2}, and potassium
                            permanganate, MnKO_{4}.
  Cobalt,     Co = 59·5   In smalt and blue glass.
  Nickel,     Ni = 59·5   For electro-plating other metals.
  Copper,     Cu = 63.    The well-known red metal.
  Zinc,       Zn = 65.    Used for the plates of batteries, roofing, &c.
  Arsenic,    As = 75.    White arsenic (poison), As_{2}O_{3}.
  Bromine,    Br = 80.    A brown volatile liquid; sodium bromide, NaBr.
  Strontium,  Sr = 87.    In  fires (SrN_{2}O_{6}).
  Silver,     Ag = 109.   The well-known white metal.
  Cadmium,    Cd = 112.   In alloys. Yellow paint (CdS).
  Tin,        Sn = 119.   The well-known metal.
  Antimony,   Sb = 120.   In alloys such as type metal.
  Iodine,     I  = 127.   In medicine and photography; free, and as KI.
  Barium,     Ba = 137.   "Permanent white," and as an adulterant in
                            white lead, and in heavy spar, BaSO_{4}.
  Platinum,   Pt = 196.}
  Gold,       Au = 197.}
  Mercury,    Hg = 200.}  Well-known metals.
  Lead,       Pb = 207.}
  Bismuth,    Bi = 209.   In medicine and fusible alloys.
  Uranium,    U  = 239.   In green fluorescent glass.

The compounds of the following metals and semi-metals have fewer
applications, but are well known, and are somewhat frequently met with
in nature, although in small quantities:--

           Beryllium,  Be = 9.           Palladium,  Pd = 107.
           Titanium,   Ti = 48.          Cerium,     Ce = 140.
           Vanadium,   V  = 51.          Tungsten,   W  = 184.
           Selenium,   Se = 79.          Osmium,     Os = 192.
           Zirconium,  Zr = 91.          Iridium,    Ir = 193.
           Molybdenum, Mo = 96.          Thallium,   Tl = 204.

The following rare metals are still more seldom met with in nature, but
have been studied somewhat fully:--

          Scandium,   Sc = 44.           Germanium,  Ge = 72.
          Gallium,    Ga = 70.           Rubidium,   Rb = 86.
          Yttrium,    Y =  89.           Cæsium,     Cs = 133.
          Niobium,    Nb = 94.           Lanthanum,  La = 138.
          Ruthenium,  Ru = 102.          Didymium,   Di = 142.
          Rhodium,    Rh = 103.          Ytterbium,  Yb = 173.
          Indium,     In = 114.          Tantalum,   Ta = 183.
          Tellurium,  Te = 125.          Thorium,    Th = 232.

Besides these 66 elements there have been discovered:--Erbium, Terbium,
Samarium, Thullium, Holmium, Mosandrium, Phillipium, and several
others. But their properties and combinations, owing to their extreme
rarity, are very little known, and even their existence as independent
substances[28] is doubtful.

  [28] Possibly some of their compounds are compounds of other
       already-known elements. Pure and incontestably independent
       compounds of these substances are unknown, and some of them have
       not even been separated, but are only supposed to exist from the
       results of spectroscopic researches. There can be no mention
       of such contestable and doubtful elements in a short general
       handbook of chemistry.

It has been incontestably proved from observations on the spectra of
the heavenly bodies that many of the commoner elements (such as H, Na,
Mg, Fe) occur on the far distant stars. This fact confirms the belief
that those forms of matter which appear on the earth as elements are
widely distributed over the entire universe. But we do not yet know why,
in nature, the mass of some elements should be greater than that of
others.[28 bis]

  [28 bis] Clark in America made an approximate calculation of the amount
       of the different elements contained in the earth's crust (to a
       depth of 15 kilometres), and found that the chief mass (over 50
       per cent.) is composed of oxygen; then comes silicon, &c.; while
       the amount of hydrogen is less than 1 per cent., carbon scarcely
       0·25 per cent., nitrogen even less than 0·03 per cent. The
       relative masses of such metals as Cu, Ni, Au is minute. Judging
       from the density (see Chapter VIII.) of the earth, a large
       proportion of its mass must be composed of iron.

The capacity of each element to combine with one or another element,
and to form compounds with them which are in a greater or less degree
prone to give new and yet more complex substances, forms the fundamental
character of each element. Thus sulphur easily combines with the
metals, oxygen, chlorine, or carbon, whilst gold and silver enter
into combinations with difficulty, and form unstable compounds, which
are easily decomposed by heat. The cause or force which induces the
elements to enter into chemical change must be considered, as also the
cause which holds different substances in combination--that is, which
endues the substances formed with their particular degree of stability.
This cause or force is called _affinity_ (_affinitas_, _affinité_,
_Verwandtschaft_), or chemical affinity.[29] Since this force must
be regarded as exclusively an attractive force, like gravity, many
writers (for instance, Bergmann at the end of the last, and Berthollet
at the beginning of this, century) supposed affinity to be essentially
similar to the universal force of gravity, from which it only differs
in that the latter acts at observable distances whilst affinity only
evinces itself at the smallest possible distances. But chemical
affinity cannot be entirely identified with the universal attraction
of gravity, which acts at appreciable distances and is dependent
only on mass and distance, and not on the quality of the material on
which it acts, whilst it is by the quality of matter that affinity is
most forcibly influenced. Neither can it be entirely identified with
cohesion, which gives to homogeneous solid substances their crystalline
form, elasticity, hardness, ductility, and other properties, and to
liquids their surface tension, drop formation, capillarity, and other
properties, because affinity acts between the component parts of a
substance and cohesion on a substance in its homogeneity, although both
act at imperceptible distances (by contact) and have much in common.
Chemical force, which makes one substance penetrate into another,
cannot be entirely identified with even those attracting forces which
make different substances adhere to each other, or hold together (as
when two plane-polished surfaces of solid substances are brought into
close contact), or which cause liquids to soak into solids, or adhere
to their surfaces, or gases and vapours to condense on the surfaces
of solids. These forces must not be confounded with chemical forces,
which cause one substance to penetrate into the substance of another
and to form a new substance, which is never the case with cohesion.
But it is evident that the forces which determine cohesion form a
connecting-link between mechanical and chemical forces, because they
only act by intimate contact. For a long time, and especially during
the first half of this century, chemical attraction and chemical forces
were identified with electrical forces. There is certainly an intimate
relation between them, for electricity is evolved in chemical reactions,
and has also a powerful influence on chemical processes--for instance,
compounds are decomposed by the action of an electrical current. And the
exactly similar relation which exists between chemical phenomena and the
phenomena of heat (heat being developed by chemical phenomena, and heat
being able to decompose compounds) only proves the unity of the forces
of nature, the capability of one force to produce and to be transformed
into others. For this reason the identification of chemical force
with electricity will not bear experimental proof.[30] As of all the
(molecular) phenomena of nature which act on substances at immeasurably
small distances, the phenomena of heat are at present the best
(comparatively) known, having been reduced to the simplest fundamental
principles of mechanics (of energy, equilibrium, and movement), which,
since Newton, have been subjected to strict mathematical analysis, it
is quite natural that an effort, which has been particularly pronounced
during recent years, should have been made to bring chemical phenomena
into strict correlation with the already investigated phenomena of
heat, without, however, aiming at any identification of chemical with
heat phenomena. The true nature of chemical force is still a secret
to us, just as is the nature of the universal force of gravity, and
yet without knowing what gravity really is, by applying mechanical
conceptions, astronomical phenomena have been subjected not only to
exact generalisation but to the detailed prediction of a number of
particular facts; and so, also, although the true nature of chemical
affinity may be unknown, there is reason to hope for considerable
progress in chemical science by applying the laws of mechanics to
chemical phenomena by means of the mechanical theory of heat. As yet
this portion of chemistry has been but little worked at, and therefore,
while forming a current problem of the science, it is treated more
fully in that particular field which is termed either 'theoretical'
or 'physical' chemistry, or, more correctly, _chemical mechanics_. As
this province of chemistry requires a knowledge not only of the various
homogeneous substances which have yet been obtained and of the chemical
transformations which they undergo, but also of the phenomena (of heat
and other kinds) by which these transformations are accompanied, it
is only possible to enter on the study of chemical mechanics after an
acquaintance with the fundamental chemical conceptions and substances
which form the subject of this book.[31]

  [29] This word, first introduced, if I mistake not, into chemistry
       by Glauber, is based on the idea of the ancient philosophers that
       combination can only take place when the substances combining
       have something in common--a medium. As is generally the case,
       another idea evolved itself in antiquity, and has lived until
       now, side by side with the first, to which it is exactly
       contradictory; this considers union as dependent on contrast, on
       polar difference, on an effort to fill up a want.

  [30] Especially conclusive are those cases of so-called metalepsis
       (Dumas, Laurent). Chlorine, in combining with hydrogen, forms
       a very stable substance called 'hydrochloric acid,' which is
       split up by the action of an electrical current into chlorine
       and hydrogen, the chlorine appearing at the positive and
       the hydrogen at the negative pole. Hence electro-chemists
       considered hydrogen to be an electro-positive and chlorine
       an electro-negative element, and that they are held together
       in virtue of their opposite electrical charges. It appears,
       however, from metalepsis, that chlorine can replace hydrogen
       (and, inversely, hydrogen can replace chlorine) in its compounds
       without in any way changing the grouping of the other elements,
       or altering their chief chemical properties. For instance, acetic
       acid in which hydrogen has been replaced by chlorine is still
       capable of forming salts. It must be observed, whilst considering
       this subject, that the explanation suggesting electricity as the
       origin of chemical phenomena is unsound, since it attempts to
       explain one class of phenomena whose nature is almost unknown by
       another class which is no better known. It is most instructive
       to remark that together with the electrical theory of chemical
       attraction there arose and survives a view which explains the
       galvanic current as being a transference of chemical action
       through the circuit--_i.e._, regards the origin of electricity as
       being a chemical one. It is evident that the connection is very
       intimate, although both phenomena are independent and represent
       different forms of molecular (atomic) motion, whose real nature
       is not yet understood. Nevertheless, the connection between
       the phenomena of both categories is not only in itself very
       instructive, but it extends the applicability of the general idea
       of the unity of the forces of nature, conviction of the truth of
       which has held so important a place in the science of the last
       ten years.

  [31] I consider that in an elementary text-book of chemistry, like the
       present, it is only possible and advisable to mention, in
       reference to chemical mechanics, a few general ideas and some
       particular examples referring more especially to gases, whose
       mechanical theory must be regarded as the most complete. The
       molecular mechanics of liquids and solids is as yet in embryo,
       and contains much that is disputable; for this reason, chemical
       mechanics has made less progress in relation to these substances.
       It may not be superfluous here to remark, with respect to the
       conception of chemical affinity, that up to the present time
       gravity, electricity, and heat have all been applied to its
       elucidation. Efforts have also been made to introduce the
       luminiferous ether into theoretical chemistry, and should that
       connection between the phenomena of light and electricity
       which was established by Maxwell be worked out more in detail,
       doubtless these efforts to elucidate all or a great deal by
       the aid of luminiferous ether will again appear in theoretical
       chemistry. An independent chemical mechanics of the material
       particles of matter, and of their internal (atomic) changes,
       would, in my opinion, arise as the result of these efforts.
       Two hundred years ago Newton laid the foundation of a truly
       scientific theoretical mechanics of external visible motion, and
       on this foundation erected the edifice of celestial mechanics.
       One hundred years ago Lavoisier arrived at the first fundamental
       law of the internal mechanics of invisible particles of matter.
       This subject is far from having been developed into a harmonious
       whole, because it is much more difficult, and, although many
       details have been completely investigated, it does not possess
       any starting points. Newton only came after Copernicus and
       Kepler, who had discovered empirically the exterior simplicity
       of celestial phenomena. Lavoisier and Dalton may, in respect to
       the chemical mechanics of the molecular world, be compared to
       Copernicus and Kepler. But a Newton has not yet appeared in the
       molecular world; when he does, I think that he will find the
       fundamental laws of the mechanics of the invisible motions of
       matter more easily and more quickly in the chemical structure
       of matter than in physical phenomena (of electricity, heat,
       and light); for these latter are accomplished by particles of
       matter already arranged, whilst it is now clear that the problem
       of chemical mechanics mainly lies in the apprehension of those
       motions which are invisibly accomplished by the smallest atoms of
       matter.

As the chemical changes to which substances are liable proceed from
internal forces proper to these substances, as chemical phenomena
certainly consist of motions of material parts (from the laws
of the indestructibility of matter and of elements), and as the
investigation of mechanical and physical phenomena proves the law of
the _indestructibility of forces_, or the conservation of energy--that
is, the possibility of the transformation of one kind of motion into
another (of visible or mechanical into invisible or physical)--we are
inevitably obliged to acknowledge the presence in substances (and
especially in the elements of which all others are composed) of a store
of _chemical energy_ or invisible motion inducing them to enter into
combinations. If heat be evolved in a reaction, it means that a portion
of chemical energy is transformed into heat;[32] if heat be absorbed
in a reaction,[33] that it is partly transformed (rendered latent) into
chemical energy. The store of force or energy going to the formation
of new compounds may, after several combinations, accomplished with an
absorption of heat, at last diminish to such a degree that indifferent
compounds will be obtained, although these sometimes, by combining with
energetic elements or compounds, give more complex compounds, which
may be capable of entering into chemical combination. Among elements,
gold, platinum, and nitrogen have but little energy, whilst potassium,
oxygen, and chlorine have a very marked degree of energy. When
dissimilar substances enter into combination they often form substances
of diminished energy. Thus sulphur and potassium when heated easily burn
in air, but when combined together their compound is neither inflammable
nor burns in air like its component parts. Part of the energy of the
potassium and of the sulphur was evolved in their combination in the
form of heat. Just as in the passage of substances from one physical
state into another a portion of their store of heat is absorbed or
evolved, so in combinations or decompositions and in every chemical
process, there occurs a change in the store of chemical energy, and at
the same time an evolution or absorption of heat.[34]

  [32] The theory of heat gave the idea of a store of internal motion
       or energy, and therefore with it, it became necessary to
       acknowledge chemical energy, but there is no foundation whatever
       for identifying heat energy with chemical energy. It may be
       supposed, but not positively affirmed, that heat motion is
       proper to molecules and chemical motion to atoms, but that as
       molecules are made up of atoms, the motion of the one passes to
       the other, and that for this reason heat strongly influences
       reaction and appears or disappears (is absorbed) in reactions.
       These relations, which are apparent and hardly subject to doubt
       on general lines, still present much that is doubtful in detail,
       because all forms of molecular and atomic motion are able to pass
       into each other.

  [33] The reactions which take place (at the ordinary or at a high
       temperature) directly between substances may be clearly divided
       into exothermal, which are accompanied by an evolution of heat,
       and endothermal, which are accompanied by an absorption of heat.
       It is evident that the latter require a source of heat. They are
       determined either by the directly surrounding medium (as in the
       formation of carbon bisulphide from charcoal and sulphur, or in
       decompositions which take place at high temperatures), or else by
       a secondary reaction proceeding simultaneously, or by some other
       form of energy (light, electricity). So, for instance, hydrogen
       sulphide is decomposed by iodine in the presence of water at the
       expense of the heat which is evolved by the solution in water
       of the hydrogen iodide produced. This is the reason why this
       reaction, as exothermal, only takes place in the presence of
       water; otherwise it would be accompanied by a cooling effect. As
       in the combination of dissimilar substances, the bonds existing
       between the molecules and atoms of the homogeneous substances
       have to be broken asunder, whilst in reactions of rearrangement
       the formation of any one substance proceeds simultaneously with
       the formation of another, and, as in reactions, a series of
       physical and mechanical changes take place, it is impossible to
       separate the heat directly depending on a given reaction from
       the total sum of the observed heat effect. For this reason,
       thermochemical data are very complex, and cannot by themselves
       give the key to many chemical problems, as it was at first
       supposed they might. They ought to form a part of chemical
       mechanics, but alone they do not constitute it.

  [34] As chemical reactions are effected by heating, so the heat
       absorbed by substances before decomposition or change of
       state, and called 'specific heat,' goes in many cases to the
       preparation, if it may be so expressed, of reaction, even when
       the limit of the temperature of reaction is not attained. The
       molecules of a substance A, which is not able to react on a
       substance B below a temperature _t_, by being heated from a
       somewhat lower temperature to _t_, undergoes that change which
       had to be arrived at for the formation of A B.

For the comprehension of chemical phenomena as mechanical
processes--_i.e._, the study of the _modus operandi_ of chemical
phenomena--it is most important to consider: (1) the facts gathered
from stoïchiometry, or that part of chemistry which treats of the
quantitative relation, by weight or volume, of the reacting substances;
(2) the distinction between the different forms and classes of
chemical reactions; (3) the study of the changes in properties
produced by alteration in composition; (4) the study of the phenomena
which accompany chemical transformation; (5) a generalisation of the
conditions under which reactions occur. As regards stoïchiometry,
this branch of chemistry has been worked out most thoroughly, and
comprises laws (of Dalton, Avogadro-Gerhardt, and others) which bear
so deeply on all parts of chemistry that at the present time the chief
problem of chemical research consists in the application of general
stoïchiometrical laws to concrete examples, _i.e._, the quantitative
(volumetric or gravimetric) composition of substances. All other
branches of chemistry are clearly subordinate to this most important
portion of chemical knowledge. Even the very signification of reactions
of combination, decomposition, and rearrangement, acquired, as we shall
see, a particular and new character under the influence of the progress
of exact ideas concerning the quantitative relations of substances
entering into chemical changes. Furthermore, in this sense there arose
a new--and, till then, unknown--division of compound substances into
_definite_ and _indefinite_ compounds. Even at the beginning of this
century, Berthollet had not made this distinction. But Prout showed
that a number of compounds contain the substances of which they are they
break up, in exact definite proportions by weight, which are unalterable
under any conditions. Thus, for example, red mercury oxide always
contains sixteen parts by weight of oxygen for every 200 parts by weight
of mercury, which is expressed by the formula HgO. But in an alloy of
copper and silver one or the other metal may be added at will, and in an
aqueous solution of sugar, the relative proportion of the sugar and water
may be altered and nevertheless a homogeneous whole with the sum of the
independent properties will be obtained--_i.e._, in these cases there was
indefinite chemical combination. Although in nature and chemical practice
the formation of indefinite compounds (such as alloys and solutions)
plays as essential a part as the formation of definite chemical
compounds, yet, as the stoïchiometrical laws at present apply chiefly to
the latter, all facts concerning indefinite compounds suffer from
inexactitude, and it is only during recent years that the attention of
chemists has been directed to this province of chemistry.

In chemical mechanics it is, from a qualitative point of view, very
important to clearly distinguish at the very beginning between
_reversible_ and _non-reversible reactions_. Substances capable of
reacting on each other at a certain temperature produce substances which
at the same temperature either can or cannot give back the original
substances. For example, salt dissolves in water at the ordinary
temperature, and the solution so obtained is capable of breaking up
at the same temperature, leaving salt and separating the water by
evaporation. Carbon bisulphide is formed from sulphur and carbon at
about the same temperature at which it can be resolved into sulphur
and carbon. Iron, at a certain temperature, separates hydrogen from
water, forming iron oxide, which, in contact with hydrogen at the same
temperature, is able to produce iron and water. It is evident that if
two substances, A and B, give two others C and D, and the reaction be
reversible, then C and D will form A and B, and, consequently, by taking
a definite mass of A and B, or a corresponding mass of C and D, we shall
obtain, in each case, all four substances--that is to say, there will be
a state of _chemical equilibrium_ between the reacting substances. By
increasing the mass of one of the substances we obtain a new condition
of equilibrium, so that reversible reactions present a means of studying
the _influence of mass_ on the _modus operandi_ of chemical changes.
Many of those reactions which occur with very complicated compounds or
mixtures may serve as examples of non-reversible reactions. Thus many of
the compound substances of animal and vegetable organisms are broken up
by heat, but cannot be re-formed from their products of decomposition at
any temperature. Gunpowder, as a mixture of sulphur, nitre, and carbon,
on being exploded, forms gases from which the original substances cannot
be re-formed at any temperature. In order to obtain them, recourse
must be had to an indirect method _of combination at the moment of
separation_. If A does not under any circumstances combine directly
with B, it does not follow that it cannot give a compound A B. For A
can often combine with C and B with D, and if C has a great affinity
for D, then the reaction of A C or B D produces not only C D, but also
A B. As on the formation of C D, the substances A and B (previously in
A C and B D) are left in a peculiar state of separation, it is supposed
that their mutual combination occurs because they meet together in this
_nascent state_ at the moment of separation (_in statu nascendi_). Thus
chlorine does not directly combine with charcoal, graphite, or diamond;
there are, nevertheless, compounds of chlorine with carbon, and many
of them are distinguished by their stability. They are obtained in the
action of chlorine on hydrocarbons, as the separation products from the
direct action of chlorine on hydrogen. Chlorine takes up the hydrogen,
and the carbon liberated at the moment of its separation, enters into
combination with another portion of the chlorine, so that in the end the
chlorine is combined with both the hydrogen and the carbon.[35]

  [35] It is possible to imagine that the cause of a great many of such
       reactions is, that substances taken in a separate state, for
       instance, charcoal, present a complex molecule composed of
       separate atoms of carbon which are fastened together (united, as
       is usually said) by a considerable affinity; for atoms of the
       same kind, just like atoms of different kinds, possess a mutual
       affinity. The affinity of chlorine for carbon, although unable
       to break this bond asunder, may be sufficient to form a stable
       compound with atoms of carbon, which are already separate. Such
       a view of the subject presents a hypothesis which, although
       dominant at the present time, is without sufficiently firm
       foundation. It is evident, however, that not only does chemical
       reaction itself consist of motions, but that in the compound
       formed (in the molecules) the elements (atoms) forming it are in
       harmonious stable motion (like the planets in the solar system),
       and this motion will affect the stability and capacity for
       reaction, and therefore the mechanical side of chemical action
       must be exceedingly complex. Just as there are solid, physically
       constant non-volatile substances like rock, gold, charcoal,
       &c., so are there stable and chemically constant bodies; while
       corresponding to physically volatile substances there are bodies
       like camphor, which are chemically unstable and variable.

As regards those phenomena which accompany chemical action, the most
important circumstance in reference to chemical mechanics is that not
only do chemical processes produce a mechanical displacement (a motion
of particles), heat, light, electrical potential and current; but that
all these agents are themselves capable of changing and governing
chemical transformations. This reciprocity or reversibility naturally
depends on the fact that all the phenomena of nature are only different
kinds and forms of visible and invisible (molecular) motions. First
sound, and then light, was shown to consist of vibratory motions, as
the laws of physics have proved and developed beyond a doubt. The
connection between heat and mechanical motion and work has ceased to
be a supposition, but has become an accepted fact, and the mechanical
equivalent of heat (425 kilogrammetres of mechanical work correspond
with one kilogram unit of heat or Calorie) gives a mechanical measure
for thermal phenomena. Although the mechanical theory of electrical
phenomena cannot be considered so fully developed as the theory of heat,
both statical and dynamical electricity are produced by mechanical
means (in common electrical machines or in Gramme or other dynamos),
and conversely, a current (in electric motors) can produce mechanical
motion. Thus by connecting a current with the poles of a Gramme dynamo
it may be made to revolve, and, conversely, by rotating it an electrical
current is produced, which demonstrates the reversibility of electricity
into mechanical motion. Accordingly chemical mechanics must look for
the fundamental lines of its advancement in the correlation of chemical
with physical and mechanical phenomena. But this subject, owing to its
complexity and comparative novelty, has not yet been expressed by a
harmonious theory, or even by a satisfactory hypothesis, and therefore
we shall avoid lingering over it.

A chemical change in a certain direction is accomplished not only by
reason of the difference of masses, the composition of the substances
concerned, the distribution of their parts, and their affinity or
chemical energy, but also by reason of the _conditions_ under which the
substances occur. In order that a certain chemical reaction may take
place between substances which are capable of reacting on each other, it
is often necessary to have recourse to conditions which are sometimes
very different from those in which the substances usually occur in
nature. For example, not only is the presence of air (oxygen) necessary
for the combustion of charcoal, but the latter must also be heated to
redness. The red-hot portion of the charcoal burns--_i.e._ combines with
the oxygen of the atmosphere--and in so doing evolves heat, which raises
the temperature of the adjacent parts of charcoal, so that they burn.
Just as the combustion of charcoal is dependent on its being heated
to redness, so also every chemical reaction only takes place under
certain physical, mechanical, or other conditions. The following are the
chief conditions which exert an influence on the progress of chemical
reactions.

(_a_) _Temperature._--Chemical reactions of combination only take
place within certain definite limits of temperature, and cannot be
accomplished outside these limits. We may cite as examples not only
that the combustion of charcoal begins at a red heat, but also that
chlorine and salt only combine with water at a temperature below 0°.
These compounds cannot be formed at a higher temperature, for they are
then wholly or partially broken up into their component parts. A certain
rise in temperature is necessary to start combustion. In certain cases
the effect of this rise may be explained as causing one of the reacting
bodies to change from a solid into a liquid or gaseous form. The
transference into a fluid form facilitates the progress of the reaction,
because it aids the intimate contact of the particles reacting on
each other. Another reason, and to this must be ascribed the chief
influence of heat in exciting chemical action, is that the physical
cohesion, or the internal chemical union, of homogeneous particles is
thereby weakened, and in this way the separation of the particles of
the original substances, and their transference into new compounds, is
rendered easier. When a reaction absorbs heat--as in decomposition--the
reason why heat is necessary is self-evident.

At the present day it may be asserted upon the basis of existing data,
respecting the action of high temperature, that all compound bodies are
decomposed at a more or less high temperature. We have already seen
examples of this in describing the decomposition of mercury oxide into
mercury and oxygen, and the decomposition of wood under the influence
of heat. Many substances are decomposed at a very moderate temperature;
for instance, the fulminating salt which is employed in cartridges is
decomposed at a little above 120°. The majority of those compounds
which make up the mass of animal and vegetable tissues are decomposed
at 200°. On the other hand, there is reason to think that at a very
low temperature no reaction whatever can take place. Thus plants cease
to carry on their chemical processes during the winter. Raoul Pictet
(1892), employing the very low temperatures (as low as -200°C.) obtained
by the evaporation of liquefied gases (_see_ Chap. II.), has recently
again proved that at temperatures below -120°, even such reactions as
those between sulphuric acid and caustic soda or metallic sodium do not
take place, and even the coloration of litmus by acids only commences
at temperatures above -80°. If a given reaction does not take place at
a certain low temperature, it will at first only proceed slowly with
a rise of temperature (even if aided by an electric discharge), and
will only proceed rapidly, with the evolution of heat, when a certain
definite temperature has been reached. Every chemical reaction requires
certain limits of temperature for its accomplishment, and, doubtless,
many of the chemical changes observed by us cannot take place in the
sun, where the temperature is very high, or on the moon, where it is
very low.

The influence of heat on reversible reactions is particularly
instructive. If, for instance, a compound which is capable of being
reproduced from its products of decomposition be heated up to the
temperature at which decomposition begins, the decomposition of a mass
of the substance contained in a definite volume is not immediately
completed. Only a certain fraction of the substance is decomposed, the
other portion remaining unchanged, and if the temperature be raised, the
quantity of the substance decomposed increases; furthermore, for a given
volume, the ratio between the part decomposed and the part unaltered
corresponds with each definite rise in temperature until it reaches that
at which the compound is entirely decomposed. This partial decomposition
under the influence of heat is called _dissociation_. It is possible
to distinguish between the temperatures at which dissociation begins
and ends. Should dissociation proceed at a certain temperature, yet
should the product or products of decomposition not remain in contact
with the still undecomposed portion of the compound, then decomposition
will go on to the end. Thus limestone is decomposed in a limekiln
into lime and carbonic anhydride, because the latter is carried off
by the draught of the furnace. But if a certain mass of limestone be
enclosed in a definite volume--for instance, in a gun barrel--which is
then sealed up, and heated to redness, then, as the carbonic anhydride
cannot escape, a certain proportion only of the limestone will be
decomposed for every increment of heat (rise in temperature) higher
than that at which dissociation begins. Decomposition will cease
when the carbonic anhydride evolved presents a maximum _dissociation
pressure_ corresponding with each rise in temperature. If the pressure
be increased by increasing the quantity of gas, then combination begins
afresh; if the pressure be diminished decomposition will recommence.
Decomposition in this case is exactly similar to evaporation; if the
steam given off by evaporation cannot escape, its pressure will reach a
maximum corresponding with the given temperature, and then evaporation
will cease. Should steam be added it will be condensed in the liquid;
if its quantity be diminished--_i.e._ if the pressure be lessened,
the temperature being constant--then evaporation will go on. We shall
afterwards discuss more fully these phenomena of dissociation, which
were first discovered by Henri St. Claire Deville. We will only remark
that the products of decomposition re-combine with greater facility the
nearer their temperature is to that at which dissociation begins, or, in
other words, that the initial temperature of dissociation is near to the
initial temperature of combination.

(_b_) _The influence of an electric current_, and of electricity in
general, on the progress of chemical transformations is very similar
to the influence of heat. The majority of compounds which conduct
electricity are decomposed by the action of a galvanic current, and as
there is great similarity in the conditions under which decomposition
and combination proceed, combination often proceeds under the influence
of electricity. Electricity, like heat, must be regarded as a peculiar
form of molecular motion, and all that refers to the influence of heat
also refers to the phenomena produced by the action of an electrical
current, with this difference, only that a substance can be separated
into its component parts with much greater ease by electricity, since
the process goes on at the ordinary temperature. The most stable
compounds may be decomposed by this means, and a most important fact
is then observed--namely, that the component parts appear at the
different poles of electrodes by which the current passes through the
substance. Those substances which appear at the positive pole (anode)
are called 'electro-negative,' and those which appear at the negative
pole (cathode, that in connection with the zinc of an ordinary galvanic
battery) are called 'electro-positive.' The majority of non-metallic
elements, such as chlorine, oxygen, &c., and also acids and substances
analogous to them, belong to the first group, whilst the metals,
hydrogen, and analogous products of decomposition appear at the negative
pole. Chemistry is indebted to the decomposition of compounds by the
electric current for many most important discoveries. Many elements
have been discovered by this method, the most important being potassium
and sodium. Lavoisier and the chemists of his time were not able to
decompose the oxygen compounds of these metals, but Davy showed that
they might be decomposed by an electric current, the metals sodium and
potassium appearing at the negative pole. Now that the dynamo gives the
possibility of producing an electric current by the combustion of fuel,
this method of Sir H. Davy is advantageously employed for obtaining
metals, &c. on a large scale, for instance, sodium from fused caustic
soda or chlorine from solutions of salt.

(_c_) Certain unstable compounds are also decomposed by _the action of
light_. Photography is based on this property in certain substances
(for instance, in the salts of silver). The mechanical energy of those
vibrations which determine the phenomena of light is very small, and
therefore only certain, and these generally unstable, compounds can be
decomposed by light--at least under ordinary circumstances. But there is
one class of chemical phenomena dependent on the action of light which
forms as yet an unsolved problem in chemistry--these are the processes
accomplished in plants under the influence of light. Here there take
place most unexpected decompositions and combinations, which are often
unattainable by artificial means. For instance, carbonic anhydride,
which is so stable under the influence of heat and electricity, is
decomposed and evolves oxygen in plants under the influence of light. In
other cases, light decomposes unstable compounds, such as are usually
easily decomposed by heat and other agents. Chlorine combines with
hydrogen under the influence of light, which shows that combination, as
well as decomposition, can be determined by its action, as was likewise
the case with heat and electricity.

(_d_) _Mechanical causes_ exert, like the foregoing agents, an action
both on the process of chemical combination and of decomposition. Many
substances are decomposed by friction or by a blow--as, for example,
the compound called iodide of nitrogen (which is composed of iodine,
nitrogen, and hydrogen), and silver fulminate. Mechanical friction
causes sulphur to burn at the expense of the oxygen contained in
potassium chlorate. Pressure affects both the physical and chemical
state of the reacting substances, and, together with the temperature,
determines the state of a substance. This is particularly evident when
the substance occurs in an elastic-gaseous form since the volume, and
hence also the number of points of encounter between the reacting
substances is greatly altered by a change of pressure. Thus, under equal
conditions of temperature, hydrogen when compressed acts more powerfully
upon iodine and on the solutions of many salts.

(_e_) Besides the various conditions which have been enumerated above,
the progress of chemical reactions is accelerated or retarded by the
_condition of contact_ in which the reacting bodies occur. Other
conditions remaining constant, the rate of progress of a chemical
reaction is accelerated by increasing the number of points of contact.
It will be enough to point out the fact that sulphuric acid does not
absorb ethylene under ordinary conditions of contact, but only after
continued shaking, by which means the number of points of contact
is greatly increased. To ensure complete action between solids, it
is necessary to reduce them to very fine powder and to mix them as
thoroughly as possible. M. Spring, the Belgian chemist, has shown
that finely powdered solids which do not react on each other at the
ordinary temperature may do so under an increased pressure. Thus, under
a pressure of 6,000 atmospheres, sulphur combines with many metals at
the ordinary temperature, and mixtures of the powders of many metals
form alloys. It is evident that an increase in the number of points
or surfaces must be regarded as the chief cause producing reaction,
which is doubtless accomplished in solids, as in liquids and gases, in
virtue of an internal motion of the particles, which motion, although
in different degrees and forms, must exist in all the states of matter.
It is very important to direct attention to the fact that the internal
motion or condition of the parts of the particles of matter must be
different on the surface of a substance from what it is inside; because
in the interior of a substance similar particles are acting on all sides
of every particle, whilst at the surface they act on one side only.
Therefore, the condition of a substance at its surfaces of contact with
other substances must be more or less modified by them--it may be in a
manner similar to that caused by an elevation of temperature. These
considerations throw some light on the action in the large class of
_contact reactions_; that is, such as appear to proceed from the mere
presence (contact) of certain special substances. Porous or powdery
substances are very prone to act in this way, especially spongy platinum
and charcoal. For example, sulphurous anhydride does not combine
directly with oxygen, but this reaction takes place in the presence of
spongy platinum.[36]

  [36] Contact phenomena are separately considered in detail in the work
       of Professor Konovaloff (1884). In my opinion, it must be
       held that the state of the internal motions of the atoms in
       molecules is modified at the points of contact of substances,
       and this state determines chemical reactions, and therefore,
       that reactions of combination, decomposition, and rearrangement
       are accomplished by contact. Professor Konovaloff showed that a
       number of substances, under certain conditions of their surfaces,
       act by contact; for instance, finely divided silica (from the
       hydrate) acts just like platinum, decomposing certain compound
       ethers. As reactions are only accomplished under close contact,
       it is probable that those modifications in the distribution of
       the atoms in molecules which come about by contact phenomena
       prepare the way for them. By this the _rôle_ of contact phenomena
       is considerably extended. Such phenomena should explain the fact
       why a mixture of hydrogen and oxygen yields water (explodes) at
       different temperatures, according to the kind of heated substance
       which transmits this temperature. In chemical mechanics,
       phenomena of this kind have great importance, but as yet they
       have been but little studied. It must not be forgotten that
       contact is a necessary condition for every chemical reaction.

The above general and introductory chemical conceptions cannot be
thoroughly grasped in their true sense without a knowledge of the
particular facts of chemistry to which we shall now turn our attention.
It was, however, absolutely necessary to become acquainted on the
very threshold with such fundamental principles as the laws of the
indestructibility of matter and of the conservation of energy, since it
is only by their acceptance, and under their direction and influence,
that the examination of particular facts can give practical and fruitful
results.




                                CHAPTER I

                       ON WATER AND ITS COMPOUNDS


Water is found almost everywhere in nature, and in all three physical
states. As vapour, water occurs in the atmosphere, and in this form
it is distributed over the entire surface of the earth. The vapour of
water in condensing, by cooling, forms snow, rain, hail, dew, and fog.
One cubic metre (or 1,000,000 cubic centimetres, or 1,000 litres, or
35·316 cubic feet) of air can contain at 0° only 4·8 grams of water, at
20° about 17·0 grams, at 40° about 50·7 grams; but ordinary air only
contains about 60 per cent. of this maximum. Air containing less than
40 per cent. is felt to be dry, whilst air which contains more than
80 per cent. of the same maximum is considered as distinctly damp.[1]
Water in the liquid state, in falling as rain and snow, soaks into
the soil and collects together into springs, lakes, rivers, seas, and
oceans. It is absorbed from the soil by the roots of plants, which, when
fresh, contain from 40 to 80 per cent. of water by weight. Animals
contain about the same amount of water. In a solid state, water appears
as snow, ice, or in an intermediate form between these two, which is
seen on mountains covered with perpetual snow. The water of rivers,[2]
springs, oceans and seas, lakes, and wells contains various substances
in solution mostly salt,--that is, substances resembling common table
salt in their physical properties and chief chemical transformations.
Further, the quantity and nature of these salts differ in different
waters.[3] Everybody knows that there are salt, fresh, iron, and
other waters. The presence of about 3-1/2 per cent. of salts renders
sea-water[4] bitter to the taste and increases its specific gravity.
Fresh water also contains salts, but only in a comparatively small
quantity. Their presence may be easily proved by simply evaporating
water in a vessel. On evaporation the water passes away as vapour,
whilst the salts are left behind. This is why a crust (incrustation),
consisting of salts, previously in solution, is deposited on the insides
of kettles or boilers, and other vessels in which water is boiled.
Running water (rivers, &c.) is charged with salts, owing to its being
formed from the collection of rain water percolating through the soil.
While percolating, the water dissolves certain parts of the soil. Thus
water which filters or passes through saline or calcareous soils becomes
charged with salts or contains calcium carbonate (chalk). Rain water and
snow are much purer than river or spring water. Nevertheless, in passing
through the atmosphere, rain and snow succeed in catching the dust held
in it, and dissolve air, which is found in every water. The dissolved
gases of the atmosphere are partly disengaged, as bubbles from water on
heating, and water after long boiling is quite freed from them.

  [1] In practice, the chemist has to continually deal with gases, and
      gases are often collected over water; in which case a certain
      amount of water passes into vapour, and this vapour mixes with
      the gases. It is therefore most important that he should be able
      to calculate the amount of water or of _moisture in air and other
      gases_. Let us imagine a cylinder standing in a mercury bath, and
      filled with a dry gas whose volume equals _v_, temperature _t_°,
      and pressure or tension _h_ mm. (_h_ millimetres of the column
      of mercury at 0°). We will introduce water into the cylinder in
      such a quantity that a small part remains in the liquid state, and
      consequently that the gas will be saturated with aqueous vapour;
      the volume of the gas will then increase (if a larger quantity
      of water be taken some of the gas will he dissolved in it, and
      the volume may therefore he diminished). We will further suppose
      that, after the addition of the water, the temperature remains
      constant; then since the volume increases, the mercury in the
      cylinder falls, and therefore the pressure as well as the volume
      is increased. In order to investigate the phenomenon we will
      artificially increase the pressure, and reduce the volume to the
      original volume _v_. Then the pressure or tension will be greater
      than _h_, namely _h_ + _f_, which means that by the introduction
      of aqueous vapour the pressure of the gas is increased. The
      researches of Dalton, Gay-Lussac, and Regnault showed that this
      increase is equal to the maximum pressure which is proper to
      the aqueous vapour at the temperature at which the observation
      is made. The maximum pressure for all temperatures may be found
      in the tables made from observations on the pressure of aqueous
      vapour. The quantity _f_ will be equal to this maximum pressure
      of aqueous vapour. This may be expressed thus: the maximum
      tension of aqueous vapour (and of all other vapours) saturating
      a space in a vacuum or in any gas is the same. This rule is
      known as _Dalton's law_. Thus we have a volume of dry gas _v_,
      under a pressure _h_, and a volume of moist gas, saturated with
      vapour, under a pressure _h_ + _f_. The volume _v_ of the dry gas
      under a pressure _h_ + _f_ occupies, from Boyle's law, a volume
      _vh_/_h_ + _f_; consequently the volume occupied by the aqueous
      vapour under the pressure _h_ + _f_ equals _v_-_vh_/(_h_ + _f_),
      or _vf_/(_h_ + _f_). Thus the volumes of the dry gas and of the
      moisture which occurs in it, at a pressure _h_ + _f_, are in the
      ratio _f_ : _h_. And, therefore, if the aqueous vapour saturates
      a space at a pressure _n_, the volumes of the dry air and of the
      moisture which is contained in it are in the ratio (_n_-_f_) :
      _f_, where _f_ is the pressure of the vapour according to the
      tables of vapour tension. Thus, if a volume N of a gas saturated
      with moisture be measured at a pressure H, then the volume of
      the gas, when dry, will be equal to N[(H-f)/H]. In fact, the
      entire volume N must be to the volume of dry gas _x_ as H is to
      H-_f_; therefore, N : _x_ = H : H-_f_, from which _x_ = N[(H-f)/H].
      Under any other pressure--for instance, 760 mm.--The volume of
      dry gas will be _x_H/760, or (H-_f_)/760, and we thus obtain
      the following practical rule: If a volume of a gas saturated
      with aqueous vapour be measured at a pressure H mm., then the
      volume of dry gas contained in it will be obtained by finding the
      volume corresponding to the pressure H, less the pressure due
      to the aqueous vapour at the temperature observed. For example,
      37·5 cubic centimetres of air saturated with aqueous vapour were
      measured at a temperature of 15·3°, and under a pressure of 747·3
      mm. of mercury (at 0°). What will be the volume of dry gas at 0°
      and 760 mm.?

      The pressure of aqueous vapour corresponding to 15·3° is equal to
      12·9 mm., and therefore the volume of dry gas at 15·3° and 747·3
      mm. is equal to 37·5 × (747·3-12·9)/747·3; at 760 mm. it will be
      equal to 37·5 × (734·4/760); and at 0° the volume of dry gas will
      be 37·5 × (734·4/760) × 273/(273 + 15·3) = 34·31 c.c.

      From this rule may also be calculated what fraction of a volume
      of gas is occupied by moisture under the ordinary pressure at
      different temperatures; for instance, at 30° C. _f_ = 31·5,
      consequently 100 volumes of a moist gas or air, at 760 mm.,
      contain a volume of aqueous vapour 100 × (31·5/760), or 4·110; it
      is also found that at 0° there is contained 0·61 p.c. by volume,
      at 10° 1·21 p.c., at 20° 2·29 p.c., and at 50° up to 12·11 p.c.
      From this it may be judged how great an error might be made in
      the measurement of gases by volume if the moisture were not taken
      into consideration. From this it is also evident how great are
      the variations in volume of the atmosphere when it loses or gains
      aqueous vapour, which again explains a number of atmospheric
      phenomena (winds, variation of pressure, rainfalls, storms, &c.)

      If a gas is not saturated, then it is indispensable that the
      degree of moisture should be known in order to determine the
      volume of dry gas from the volume of moist gas. The preceding
      ratio gives the maximum quantity of water which can be held in
      a gas, and the degree of moisture shows what fraction of this
      maximum quantity occurs in a given case, when the vapour does
      not saturate the space occupied by the gas. Consequently, if the
      degree of moisture equals 50 p.c.--that is, half the maximum--then
      the volume of dry gas at 760 mm. is equal to the volume of dry
      gas at 760 mm. multiplied by (_h_-0·5_f_)/760, or, in general, by
      (_h_-_rf_)/760 where _r_ is the degree of moisture. Thus, if it
      is required to measure the volume of a moist gas, it must either
      be thoroughly dried or quite saturated with moisture, or else
      the degree of moisture determined. The first and last methods
      are inconvenient, and therefore recourse is usually had to the
      second. For this purpose water is introduced into the cylinder
      holding the gas to be measured; it is left for a certain time so
      that the gas may become saturated, the precaution being taken
      that a portion of the water remains in a liquid state; then the
      volume of the moist gas is determined, from which that of the
      dry gas may be calculated. In order to find the _weight of the
      aqueous vapour_ in a gas it is necessary to know the weight of a
      cubic measure at 0° and 760 mm. Knowing that one cubic centimetre
      of air in these circumstances weighs 0·001293 gram, and that
      the density of aqueous vapour is 0·62, we find that one cubic
      centimetre of aqueous vapour at 0° and 760 mm. weighs 0·0008 gram,
      and at a temperature _t_° and pressure _h_ the weight of one cubic
      centimetre will be 0·0008 × _h_/760 × 273/(273 + _t_). We already
      know that _v_ volumes of a gas at a temperature _t_° pressure _h_
      contain _v_ × _f_/_h_ volumes of aqueous vapour which saturate it,
      therefore the weight of the aqueous vapour held in _v_ volumes of
      a gas will be

                    _v_ x 0·0008 × _f_/760 × 273/(273 + _t_).

      Accordingly, the weight of water which is contained in one volume
      of a gas depends only on the temperature and not on the pressure.
      This also signifies that evaporation proceeds to the same extent
      in air as in a vacuum, or, in general terms (this is _Dalton's
      law_), vapours and gases diffuse into each other as if into a
      vacuum. In a given space, at a given temperature, a constant
      quantity of vapour enters, whatever be the pressure of the gas
      filling that space.

      From this it is clear that if the weight of the vapour contained
      in a given volume of a gas be known, it is easy to determine the
      degree of moisture _r_ = _p_/(_v_ × 0·0008) × 760/_t_ × (273 +
      _t_)/273. On the is founded the very exact determination of the
      degree of moisture of air by the weight of water contained in a
      given volume. It is easy to calculate from the preceding formula
      the number of grams of water contained at any pressure in one
      cubic metre or million cubic centimetres of air saturated with
      vapour at various temperatures; for instance, at 30° _f_ = 31·5,
      hence _p_ = 29·84 grams.

      The laws of Mariotte, Dalton, and Gay-Lussac, which are here
      applied to gases and vapours, are not entirely exact, but are
      approximately true. If they were quite exact, a mixture of several
      liquids, having a certain vapour pressure, would give vapours of
      a very high pressure, which is not the case. In fact the pressure
      of aqueous vapour is slightly less in a gas than in a vacuum,
      and the weight of aqueous vapour held in a gas is slightly less
      than it should be according to Dalton's law, as was shown by the
      experiments of Regnault and others. This means that the tension
      of the vapour is less in air than in a vacuum. The difference
      does not, however, exceed 5 per cent. of the total pressure of
      the vapours. This _decrement in vapour tension_ which occurs in
      the intermixture of vapours and gases, although small, indicates
      that there is then already, so to speak, a beginning of chemical
      change. The essence of the matter is that in this case there
      occurs, as on contact (see preceding footnote), an alteration in
      the motions of the atoms in the molecules, and therefore also a
      change in the motion of the molecules themselves.

      In the uniform intermixture of air and other gases with aqueous
      vapour, and in the capacity of water to pass into vapour and
      form a uniform mixture with air, we may perceive an instance of
      a physical phenomenon which is analogous to chemical phenomena,
      forming indeed a transition from one class of phenomena to
      the other. Between water and dry air there exists a kind of
      affinity which obliges the water to saturate the air. But such a
      homogeneous mixture is formed (almost) independently of the nature
      of the gas in which evaporation takes place; even in a vacuum
      the phenomenon occurs in exactly the same way as in a gas, and
      therefore it is not the property of the gas, nor its relation to
      water, but the property of the water itself, which compels it to
      evaporate, and therefore in this case chemical affinity is not yet
      operative--at least its action is not clearly pronounced. That it
      does, however, play a certain part is seen from the deviation from
      Dalton's law.

  [2] In falling through the atmosphere, water dissolves the gases of
      the atmosphere, nitric acid, ammonia, organic compounds, salts
      of sodium, magnesium, and calcium, and mechanically washes
      out a mixture of dust and microbes which are suspended in the
      atmosphere. The amount of these and certain other constituents
      is very variable. Even in the beginning and end of the same
      rainfall a variation which is often very considerable may be
      remarked. Thus, for example, Bunsen found that rain collected
      at the beginning of a shower contained 3·7 grams of ammonia per
      cubic metre, whilst that collected at the end of the same shower
      contained only O·64 gram. The water of the entire shower contained
      an average of 1·47 gram of ammonia per cubic metre. In the course
      of a year rain supplies an acre of ground with as much as 5-1/2
      kilos of nitrogen in a combined form. Marchand found in one cubic
      metre of snow water 15·63, and in one cubic metre of rain water
      10·07, grams of sodium sulphate. Angus Smith showed that after a
      thirty hours' fall at Manchester the rain still contained 34·3
      grams of salts per cubic metre. A considerable amount of organic
      matter, namely 25 grams per cubic metre, has been found in rain
      water. The total amount of solid matter in rain water reaches 50
      grams per cubic metre. Rain water generally contains very little
      carbonic acid, whilst river water contains a considerable quantity
      of it. In considering the nourishment of plants it is necessary
      to keep in view the substances which are carried into the soil by
      rain.

      _River water_, which is accumulated from springs and sources
      fed by atmospheric water, contains from 50 to 1,600 parts by
      weight of salts in 1,000,000 parts. The amount of solid matter,
      per 1,000,000 parts by weight, contained in the chief rivers is
      as follows:--the Don 124, the Loire 135, the St. Lawrence 170,
      the Rhone 182, the Dnieper 187, the Danube from 117 to 234, the
      Rhine from 158 to 317, the Seine from 190 to 432, the Thames at
      London from 400 to 450, in its upper parts 387, and in its lower
      parts up to 1,617, the Nile 1,580, the Jordan 1,052. The Neva is
      characterised by the remarkably small amount of solid matter it
      contains. From the investigations of Prof. G. K. Trapp, a cubic
      metre of Neva water contains 32 grams of incombustible and 23
      grams of organic matter, or altogether about 55 grams. This is one
      of the purest waters which is known in rivers. The large amount
      of impurities in river water, and especially of organic impurity
      produced by pollution with putrid matter, makes the water of many
      rivers unfit for use.

      The chief part of the soluble substances in river water consists
      of the calcium salts. 100 parts of the solid residues contain
      the following amounts of calcium carbonate--from the water of
      the Loire 53, from the Thames about 50, the Elbe 55, the Vistula
      65, the Danube 65, the Rhine from 55 to 75, the Seine 75, the
      Rhone from 82 to 94. The Neva contains 40 parts of calcium
      carbonate per 100 parts of saline matter. The considerable amount
      of calcium carbonate which river water contains is very easily
      explained from the fact that water which contains carbonic acid
      in solution easily dissolves calcium carbonate, which occurs all
      over the earth. Besides calcium carbonate and sulphate, river
      water contains magnesium, silica, chlorine, sodium, potassium,
      aluminium, nitric acid, iron and manganese. The presence of salts
      of phosphoric acid has not yet been determined with exactitude
      for all rivers, but the presence of nitrates has been proved with
      certainty in almost all kinds of well-investigated river water.
      The quantity of calcium phosphate does not exceed 0·4 gram in the
      water of the Dnieper, and the Don does not contain more than 5
      grams. The water of the Seine contains about 15 grams of nitrates,
      and that of the Rhone about 8 grams. The amount of ammonia is
      much less; thus in the water of the Rhine about 0·5 gram in June,
      and 0·2 gram in October; the water of the Seine contains the same
      amount. This is less than in rain water. Notwithstanding this
      insignificant quantity, the water of the Rhine alone, which is not
      so very large a river, carries 16,245 kilograms of ammonia into
      the ocean every day. The difference between the amount of ammonia
      in rain and river water depends on the fact that the soil through
      which the rain water passes is able to retain the ammonia. (Soil
      can also absorb many other substances, such as phosphoric acid,
      potassium salts, &c.)

      The waters of springs, rivers, wells, and in general of those
      localities from which it is taken for drinking purposes, may
      be injurious to health if it contains much organic pollution,
      the more so as in such water the lower organisms (bacteria) may
      rapidly develop, and these organisms often serve as the carriers
      or causes of infectious diseases. For instance, certain pathogenic
      (disease-producing) bacteria are known to produce typhoid, the
      Siberian plague, and cholera. Thanks to the work of Pasteur,
      Metchnikoff, Koch, and many others, this province of research has
      made considerable progress. It is possible to investigate the
      number and properties of the germs in water. In bacteriological
      researches a gelatinous medium in which the germs can develop and
      multiply is prepared with gelatin and water, which has previously
      been heated several times, at intervals, to 100° (it is thus
      rendered sterile--that is to say, all the germs in it are killed).
      The water to be investigated is added to this prepared medium in
      a definite and small quantity (sometimes diluted with sterilised
      water to facilitate the calculation of the number of germs), it
      is protected from dust (which contains germs), and is left at
      rest until whole families of lower organisms are developed from
      each germ. These families (colonies) are visible to the naked eye
      (as spots), they may be counted, and by examining them under the
      microscope and observing the number of organisms they produce,
      their significance may be determined. The majority of bacteria
      are harmless, but there are decidedly pathogenic bacteria, whose
      presence is one of the causes of malady and of the spread of
      certain diseases. The number of bacteria in one cubic centimetre
      of water sometimes attains the immense figures of hundreds of
      thousands and millions. Certain well, spring, and river waters
      contain very few bacteria, and are free from disease-producing
      bacteria under ordinary circumstances. By boiling water, the
      bacteria in it are killed, but the organic matter necessary for
      their nourishment remains in the water. The best kinds of water
      for drinking purposes do not contain more than 300 bacteria in a
      cubic centimetre.

      The amount of gases dissolved in river water is much more constant
      than that of its solid constituents. One litre, or 1,000 c.c., of
      water contains 40 to 55 c.c. of gas measured at normal temperature
      and pressure. In winter the amount of gas is greater than in
      summer or autumn. Assuming that a litre contains 50 c.c. of
      gases, it may be admitted that these consist, on an average, of
      20 vols. of nitrogen, 20 vols of carbonic anhydride (proceeding
      in all likelihood from the soil and not from the atmosphere), and
      of 10 vols. of oxygen. If the total amount of gases be less, the
      constituent gases are still in about the same proportion; in many
      cases, however, carbonic anhydride predominates. The water of many
      deep and rapid rivers contains less carbonic anhydride, which
      shows their rapid formation from atmospheric water, and that they
      have not succeeded, during a long and slow course, in absorbing a
      greater quantity of carbonic anhydride. Thus, for instance, the
      water of the Rhine, near Strasburg, according to Deville, contains
      8 c.c. of carbonic anhydride, 16 c.c. of nitrogen, and 7 c.c. of
      oxygen per litre. From the researches of Prof. M. R. Kapoustin and
      his pupils, it appears that in determining the quality of a water
      for drinking purposes, it is most important to investigate the
      composition of the dissolved gases, more especially oxygen.

  [3] _Spring water_ is formed from rain water percolating through the
      soil. Naturally a part of the rain water is evaporated directly
      from the surface of the earth and from the vegetation on it. It
      has been shown that out of 100 parts of water falling on the earth
      only 36 parts flow to the ocean; the remaining 64 are evaporated,
      or percolate far underground. After flowing underground along
      some impervious strata, water comes out at the surface in many
      places as springs, whose temperature is determined by the depth
      from which the water has flowed. Springs penetrating to a great
      depth may become considerably heated, and this is why hot mineral
      springs, with a temperature of up to 30° and higher, are often met
      with. When a spring water contains substances which endow it with
      a peculiar taste, and especially if these substances are such as
      are only found in minute quantities in river and other flowing
      waters, then the spring water is termed a _mineral water_. Many
      such waters are employed for medicinal purposes. Mineral waters
      are classed according to their composition into--(_a_) saline
      waters, which often contain a large amount of common salt; (_b_)
      alkaline waters, which contain sodium carbonate; (_c_) bitter
      waters, which contain magnesia; (_d_) chalybeate waters, which
      hold iron carbonate in solution; (_e_) aërated waters, which are
      rich in carbonic anhydride; (_f_) sulphuretted waters, which
      contain hydrogen sulphide. Sulphuretted waters may be recognised
      by their smell of rotten eggs, and by their giving a black
      precipitate with lead salts, and also by their tarnishing silver
      objects. Aërated waters, which contain an excess of carbonic
      anhydride, effervesce in the air, have a sharp taste, and redden
      litmus paper. Saline waters leave a large residue of soluble
      solid matter on evaporation, and have a salt taste. Chalybeate
      waters have an inky taste, and are  black by an infusion
      of galls; on being exposed to the air they usually give a brown
      precipitate. Generally, the character of mineral waters is mixed.
      In the table below the analyses are given of certain mineral
      springs which are valued for their medicinal properties. The
      quantity of the substances is expressed in millionths by weight.

          Column Headings:
          A: Calcium salts
          B: Sodium chloride
          C: Sodium sulphate
          D: Sodium carbonate
          E: Potassium iodide and bromide

          +-------+-------+--------+-------+-------+-----+------+
          |       |       |        |       |       |     |      |
          |       |  [A]  |   [B]  |  [C]  |  [D]  | [E] | [F]  |
          |       |       |        |       |       |     |      |
          +-------+-------+--------+-------+-------+-----+------+
          |       |       |        |       |       |     |      |
          | I.    | 1,928 |    --  |   152 |   --  | --  |  24  |
          | II.   |   816 |    386 | 1,239 |    26 | --  |  43  |
          | III.  | 1,085 |  1,430 | 1,105 |   --  |  4  |  90  |
          | IV.   |   343 |  3,783 |    16 | 3,431 | --  |  14  |
          | V.    | 3,406 | 15,049 |   --  |   --  |  2  |  --  |
          | VI.   |   352 |  3,145 |   --  |    95 | 35  |  50  |
          | VII.  |   308 |  1,036 | 2,583 | 1,261 |  4  | 178  |
          | VIII. | 1,726 |  9,480 |   --  |   --  | 40  | 120  |
          | IX.   |   551 |  2,040 | 1,150 |   999 | --  |   1  |
          | X.    |   285 |    558 |   279 | 3,813 | --  |  --  |
          |       |       |        |       |       |     |      |
          | XI.   |   340 |    910 | Iron and aluminium  {1,020 |
          |       |       |        |          sulphates: {1,660 |
          +-------+-------+--------+----------------------------+

          Column Headings:
          G: Iron carbonate
          H: Magnesium salts
          I: Silica
          J: Carbonic anhydride
          K: Sulphuretted hydrogen
          L: Total solid contents

          +-------+------+-------+-----+-------+-----+-----------+
          |       |      |       |     |       |     |           |
          |       | [G]  |  [H]  | [I] |  [J]  | [K] |     [L]   |
          +-------+------+-------+-----+-------+-----+-----------+
          |       |      |       |     |       |     |           |
          | I.    |  --  |   448 | 152 | 1,300 | 80  |     2,609 |
          | II.   |   9  |   257 |  46 | 1,485 | --  |     2,812 |
          | III.  |  --  |   187 |  65 | 1,326 | 11  |     3,950 |
          | IV.   |  --  |   251 | 112 | 2,883 | --  |     7,950 |
          | V.    |  17  | 1,587 | 229 |   --  | 76  |    20,290 |
          | VI.   |   1  |   260 |  11 |    20 | --  |     3,970 |
          | VII.  |   4  |   178 |  75 |   --  | --  |     5,451 |
          | VIII. |  26  |   208 |  40 |   --  | --  |    11,790 |
          | IX.   |  30  |   209 |  50 | 2,740 | --  |     4,070 |
          | X.    |   7  |    45 |  45 | 2,268 | --  |     5,031 |
          |       |      |       |     |       {Sulphuric        |
          | XI.   |      |   940 | 190 | 2,550 {and hydrochloric |
          |       |      |       |     |   330 {acids            |
          +-------+------+-------+-----+-------------------------+


      I. Sergieffsky, a sulphur water, Gov. of Samara (temp. 8° C.),
      analysis by Clause. II. Geléznovodskya water source No. 10,
      near Patigorsk, Caucasus (temp. 22·5°), analysis by Fritzsche.
      III. Aleksandroffsky, alkaline-sulphur source, Patigorsk (temp.
      46·5°), average of analyses by Herman, Zinin and Fritzsche.
      IV. Bougountouksky, alkaline source, No. 17, Essentoukah,
      Caucasus (temp. 21·6°), analysis by Fritzsche. V. Saline water,
      Staro-Russi, Gov. of Novgorod, analysis by Nelubin. VI. Water from
      artesian well at the factory of state papers, St. Petersburg,
      analysis by Struve. VII. Sprüdel, Carlsbad (temp. 83·7°), analysis
      by Berzelius. VIII. Kreuznach spring (Elisenquelle), Prussia
      (temp. 8·8°), analysis by Bauer. IX. Eau de Seltz, Nassau,
      analysis by Henry. X. Vichy water, France, analysis by Berthier
      and Puvy. XI. Paramo de Ruiz, New Granada, analysis by Levy; it is
      distinguished by the amount of free acids.

  [4] _Sea water_ contains more non-volatile saline constituents than
      the usual kinds of fresh water. This is explained by the fact that
      the waters flowing into the sea supply it with salts, and whilst a
      large quantity of vapour is given off from the surface of the sea,
      the salts remain behind. Even the specific gravity of sea water
      differs considerably from that of pure water. It is generally
      about 1·02, but in this and also in respect of the amount of salts
      contained, samples of sea water from different localities and from
      different depths offer rather remarkable variations. It will be
      sufficient to point out that one cubic metre of water from the
      undermentioned localities contains the following quantity in grams
      of solid constituents:--Gulf of Venice, 19,122; Leghorn Harbour
      24,312; Mediterranean, near Cetta, 37,665; the Atlantic Ocean
      from 32,585 to 35,695,; the Pacific Ocean from 35,233 to 34,708.
      In closed seas which do not communicate, or are in very distant
      communication, with the ocean, the difference is often still
      greater. Thus the Caspian Sea contains 6,300 grams; the Black
      Sea and Baltic 17,700. Common salt forms the chief constituent
      of the saline matter of sea or ocean water; thus in one cubic
      metre of sea water there are 25,000-31,000 grams of common salt,
      2,600-6,000 grams of magnesium chloride, 1,200-7,000 grams of
      magnesium sulphate, 1,500-6,000 grams of calcium sulphate, and
      10-700 grams of potassium chloride. The small amount of organic
      matter and of the salts of phosphoric acid in sea water is very
      remarkable. Sea water (the composition of which is partially
      discussed in Chapter X.) contains, in addition to salts of common
      occurrence, a certain and sometimes minute amount of the most
      varied elements, even gold and silver, and as the mass of water of
      the oceans is so enormous these 'traces' of rare substances amount
      to large quantities, so that it may be hoped that in time methods
      will be found to extract even gold from sea water, which by means
      of the rivers forms a vast reservoir for the numerous products
      of the changes taking place on the earth's surface. The works of
      English, American, German, Russian, Swedish, and other navigators
      and observers prove that a study of the composition of sea water
      not only explains much in the history of the earth's life, but
      also gives the possibility (especially since the researches of
      C. O. Makaroff of the St. Petersburg Academy) of fixing one's
      position in the ocean in the absence of other means, for instance,
      in a fog, or in the dark.

In general terms water is called pure when it is clear and free from
insoluble particles held in suspension and visible to the naked eye,
from which it may be freed by filtration through charcoal, sand, or
porous (natural or artificial) stones, and when it possesses a clean
fresh taste. It depends on the absence of any taste, decomposing organic
matter, on the quantity of air[5] and atmospheric gases in solution, and
on the presence of mineral substances to the amount of about 300 grams
per ton (or 1000 kilograms per cubic metre, or, what is the same, 300
milligrams to a kilogram or a litre of water), and of not more than 100
grams of organic matter.[6] Such water is suitable for drinking and
every practical application, but evidently it is not pure in a chemical
sense. A _chemically pure water_ is necessary not only for scientific
purposes, as an independent substance having constant and definite
properties, but also for many practical purposes--for instance, in
photography and in the preparation of medicines--because many properties
of substances in solution are changed by the impurities of natural
waters. Water is usually purified by distillation, because the solid
substances in solution are not transformed into vapours in this process.
Such _distilled_ water is prepared by chemists and in laboratories by
boiling water in closed metallic boilers or stills, and causing the
steam produced to pass into a condenser--that is, through tubes (which
should be made of tin, or, at all events, tinned, as water and its
impurities do not act on tin) surrounded by cold water, and in which
the steam, being cooled, condenses into water which is collected[7] in
a receiver. By standing exposed to the atmosphere, however, the water
in time absorbs air, and dust carried in the air. Nevertheless, in
distillation, water retains, besides air, a certain quantity of volatile
impurities (especially organic) and the walls of the distillation
apparatus are partly corroded by the water, and a portion, although
small, of their substance renders the water not entirely pure, and a
residue is left on evaporation.[8]

  [5] The taste of water is greatly dependent on the quantity of
      dissolved gases it contains. These gases are given off on
      boiling, and it is well known that, even when cooled, boiled
      water has, until it has absorbed gaseous substances from the
      atmosphere, quite a different taste from fresh water containing
      a considerable amount of gas. The dissolved gases, especially
      oxygen and carbonic anhydride, have an important influence on
      the health. The following instance is very instructive in this
      respect. The Grenelle artesian well at Paris, when first opened,
      supplied a water which had an injurious effect on men and animals.
      It appeared that this water did not contain oxygen, and was in
      general very poor in gases. As soon as it was made to fall in
      a cascade, by which it absorbed air, it proved quite fit for
      consumption. In long sea voyages fresh water is sometimes not
      taken at all, or only taken in a small quantity, because it spoils
      by keeping, and becomes putrid from the organic matter it contains
      undergoing decomposition. Fresh water may he obtained directly
      from sea-water by distillation. The distilled water no longer
      contains sea salts, and is therefore fit for consumption, but
      it is very tasteless and has the properties of boiled water. In
      order to render it palatable certain salts, which are usually held
      in fresh water, are added to it, and it is made to flow in thin
      streams exposed to the air in order that it may become saturated
      with the component parts of the atmosphere--that is, absorb gases.

  [6] _Hard water_ is such as contains much mineral matter, and
      especially a large proportion of calcium salts. Such water,
      owing to the amount of lime it contains, does not form a lather
      with soap, prevents vegetables boiled in it from softening
      properly, and forms a large amount of incrustation on vessels
      in which it is boiled. When of a high degree of hardness, it
      is injurious for drinking purposes, which is evident from the
      fact that in several large cities the death-rate has been found
      to decrease after introducing a soft water in the place of a
      hard water. _Putrid water_ contains a considerable quantity of
      decomposing organic matter, chiefly vegetable, but in populated
      districts, especially in towns, chiefly animal remains. Such
      water acquires an unpleasant smell and taste, by which stagnant
      bog water and the water of certain wells in inhabited districts
      are particularly characterised. Water of this kind is especially
      injurious at a period of epidemic. It may be partially purified
      by being passed through charcoal, which retains the putrid and
      certain organic substances, and also certain mineral substances.
      Turbid water may be purified to a certain extent by the addition
      of alum, which aids, after standing some time, the formation of
      a sediment. Condy's fluid (potassium permanganate) is another
      means of purifying putrid water. A solution of this substance,
      even if very dilute, is of a red colour; on adding it to a putrid
      water, the permanganate oxidises and destroys the organic matter.
      When added to water in such a quantity as to impart to it an
      almost imperceptible rose colour it destroys much of the organic
      substances it contains. It is especially salutary to add a small
      quantity of Condy's fluid to impure water in times of epidemic.

      The presence in water of one gram per litre, or 1,000 grams per
      cubic metre, of any substance whatsoever, renders it unfit and
      even injurious for consumption by animals, and this whether
      organic or mineral matter predominates. The presence of 1 p.c. of
      chlorides makes water quite salt, and produces thirst instead of
      assuaging it. The presence of magnesium salts is most unpleasant;
      they have a disagreeable bitter taste, and, in fact, impart to sea
      water its peculiar taste. A large amount of nitrates is only found
      in impure water, and is usually injurious, as they may indicate
      the presence of decomposing organic matter.

  [7] [Illustration: FIG. 4.--Distillation by means of a metallic still.
      The liquid in C is heated by the fire F. The vapours rise through
      the head A and pass by the tube T to the worm S placed in a vessel
      R, through which a current of cold water flows by means of the
      tubes D and P.]

      Distilled water may be prepared, or distillation in general
      carried on, either in a metal still with worm condenser (fig. 4)
      or on a small scale in the laboratory in a glass retort (fig.
      5) heated by a lamp. Fig. 5 illustrates the main parts of the
      usual glass laboratory apparatus used for distillation. The steam
      issuing from the retort (on the right-hand side) passes through a
      glass tube surrounded by a larger tube, through which a stream of
      cold water passes, by which the steam is condensed, and runs into
      a receiver (on the left-hand side).

      [Illustration: FIG. 5.--Distillation from a glass retort. The neck
      of the retort fits into the inner tube of the Liebig's condenser.
      The space between the inner and outer tube of the condenser is
      filled with cold water, which enters by the tube _g_ and flows out
      at _f_.]


  [8] One of Lavoisier's first memoirs (1770) referred to this question.
      He investigated the formation of the earthy residue in the
      distillation of water in order to prove whether it was possible,
      as was affirmed, to convert water into earth, and he found that
      the residue was produced by the action of water on the sides of
      the vessel containing it, and not from the water itself. He proved
      this to be the case by direct weighing.

For certain physical and chemical researches, however, it is necessary
to have perfectly pure water. To obtain it, a solution of potassium
permanganate is added to distilled water until the whole is a light
rose colour. By this means the organic matter in the water is destroyed
(converted into gases or non-volatile substances). An excess of
potassium permanganate does no harm, because in the next distillation it
is left behind in the distillation apparatus. The second distillation
should take place in a platinum retort with a platinum receiver.
Platinum is a metal which is not acted on either by air or water, and
therefore nothing passes from it into the water. The water obtained
in the receiver still contains air. It must then be boiled for a long
time, and afterwards cooled in a vacuum under the receiver of an air
pump. Pure water does not leave any residue on evaporation; does not
in the least change, however long it be kept; does not decompose like
water only once distilled or impure; and it does not give bubbles of
gas on heating, nor does it change the colour of a solution of potassium
permanganate.

Water, purified as above described, has constant _physical_ and
_chemical properties_. For instance, it is of such water only that
one cubic centimetre weighs one gram at 4° C.--_i.e._ it is only such
pure water whose specific gravity equals 1 at 4° C.[9] Water in a
solid state forms crystals of the hexagonal system[10] which are seen
in snow, which generally consists of star-like clusters of several
crystals, and also in the half-melted scattered ice floating on rivers
in spring time. At this time of the year the ice splits up into spars
or prisms, bounded by angles proper to substances crystallising in the
hexagonal system.

  [9] Taking the generally-accepted specific gravity of water at its
      greatest density--_i.e._ at 4° as one--it has been shown by
      experiment that the specific gravity of water at different
      temperatures is as follows:

                   At   0°  0·99987 | At  30°  0·99574
                   "  +10°  0·99974 | "   40°  0·99233
                   "   15°  0·99915 | "   50°  0·98817
                   "   20°  0·99827 | "  100°  0·95859

      A comparison of all the data at present known shows that the
      variation of the specific gravity S_{t} with the temperature
      _t_ (determined by the mercurial thermometer) maybe expressed
      (Mendeléeff 1891) by the formula

        S_{t} = 1 - (_t_-4)^{2}/(94·1 + _t_) (703·51-_t_) 1·9

      +-----------+-------------+---------------------------+-----------+
      |   t° C.   |             | Variation of sp. gr. with |           |
      | according |Sp. gr. S_{t}|          a rise of        |  Volume   |
      |  to the   |(at 4° =     +--------------+------------+taking vol.|
      | mercurial |  1,000,000) |Temp. per 1°C.|Pressure per| at 4° = 1 |
      |thermometer|             |   or ds/dt   |1 atmosphere|           |
      |           |             |              |  or ds/dp  |           |
      +-----------+-------------+--------------+------------+-----------+
      |   -10     |   998,281   |     +264     |     +54    | 1,001,722 |
      |     0     |   999,873   |      +65     |     +50    | 1,000,127 |
      |    10     |   999,738   |      -85     |     +47    | 1,000,262 |
      |    20     |   998,272   |     -203     |     +45    | 1,001,731 |
      |    30     |   995,743   |     -299     |     +43    | 1,004,276 |
      |    50     |   988,174   |     -450     |     +40    | 1,011,967 |
      |    70     |   977,948   |     -569     |     +39    | 1,022,549 |
      |    90     |   965,537   |     -670     |     +41    | 1,035,692 |
      |   100     |   958,595   |     -718     |     +42    | 1,043,194 |
      |   120     |   943,814   |     -810     |     +43    | 1,060,093 |
      |   160     |   907,263   |     -995     |     +55    | 1,102,216 |
      |   200     |   863,473   |   -1,200     |     +73    | 1,158,114 |
      +-----------+-------------+--------------+------------+-----------+

      If the temperature be determined by the hydrogen thermometer,
      whose indications between 0° and 100° are slightly lower than the
      mercurial (for example, about 0·1° C. at 20°), then a slightly
      smaller sp. gr. will be obtained for a given _t_. Thus Chappuis
      (1892) obtained 0·998233 for 20°. Water at 4° is taken as the
      basis for reducing measures of length to measures of weight and
      volume. The _metric, decimal, system_ of measures of weights and
      volumes is generally employed in science. The starting point of
      this system is the metre (39·37 inches) divided into decimetres
      (= 0·1 metre), centimetres (= 0·01 metre), millimetres (= 0·001
      metre), and micrometres (= one millionth of a metre). A cubic
      decimetre is called a _litre_, and is used for the measurement
      of volumes. The weight of a litre of water at 4° in a vacuum is
      called a kilogram. One thousandth part of a kilogram of water
      weighs one _gram_. It is divided into decigrams, centigrams,
      and milligrams (= 0·001 gram). An English pound equals 453·59
      grams. The great advantage of this system is that it is a decimal
      one, and that it is universally adopted in science and in most
      international relations. _All the measures cited in this work are
      metrical._ The units most often used in science are:--Of length,
      the centimetre; of weight, the gram; of time, the second; of
      temperature, the degree Celsius or Centigrade. According to the
      most trustworthy determinations (Kupfer in Russia 1841, and Chaney
      in England 1892), the weight of a c. dcm. of water at 4° in vacuo
      is about 999·9 grms. For ordinary purposes the weight of a c.
      dcg. may be taken as equal to a kg. Hence the litre (determined
      by the weight of water it holds) is slightly greater than a cubic
      decimetre.

  [10] As solid substances appear in independent, regular, crystalline
       forms which are dependent, judging from their cleavage or
       lamination (in virtue of which mica breaks, up into laminae,
       and Iceland spar, &c., into pieces bounded by faces inclined to
       each other at angles which are definite for each substance), on
       an inequality of attraction (cohesion, hardness) in different
       directions which intersect at definite angles the determination
       of crystalline form therefore affords one of the most important
       characteristics for identifying definite chemical compounds. The
       elements of crystallography which comprise a special science
       should therefore he familiar to all who desire to work in
       scientific chemistry. In this work we shall only have occasion
       to speak of a few crystalline forms, some of which are shown in
       figs. 6 to 12.

       [Illustration: FIG. 6.--Example of the form belonging to the
       regular system. Combination of an octahedron and a cube. The
       former predominates. Alum, fluor spar, suboxide of copper, and
       others.]

       [Illustration: FIG. 7.--Rhombic Dodecahedron of the regular
       system. Garnet.]

       [Illustration: FIG. 8.--Hexagonal prism terminated by hexagonal
       pyramids. Quartz, &c.]

       [Illustration: FIG. 9.--Rhombohedron. Calc spar, &c.]

       [Illustration: FIG. 10.--Rhombic system. Desmine.]

       [Illustration: FIG. 11.--Triclinic pyramid.]

       [Illustration: FIG. 12.--Triclinic system. Albite, &c.]

The temperatures at which water passes from one state to another are
taken as fixed points on the thermometer scale; namely, the zero
corresponds with the temperature of melting ice, and the temperature
of the steam disengaged from water boiling at the normal barometric
pressure (that is 760 millimetres measured at 0°, at the latitude of
45°, at the sea level) is taken as 100° of the Celsius scale. Thus, the
fact that water liquefies at 0° and boils at 100° is taken as one of
its properties as a definite chemical compound. The weight of a litre
of water at 4° is 1,000 grams, at 0° it is 999·8 grams. The weight of a
litre of ice at 0° is less--namely, 917 grams; the weight of the same
cubic measure of water vapour at 760 mm. pressure and 100° is only 0·60
gram; the density of the vapour compared with air = 0·62, and compared
with hydrogen = 9.

These data briefly characterise the physical properties of water as a
separate substance. To this may be added that water is a mobile liquid,
colourless, transparent, without taste or smell, &c. Its latent heat
of vaporisation is 534 units, of liquefaction 79 units of heat.[11]
The large amount of heat stored up in water vapour and also in liquid
water (for its specific heat is greater than that of other liquids)
renders it available in both forms for heating purposes. The chemical
reactions which water undergoes, and by means of which it is formed,
are so numerous, and so closely allied to the reactions of many other
substances, that it is impossible to describe the majority of them at
this early stage of chemical exposition. We shall become acquainted
with many of them afterwards, but at present we shall only cite certain
compounds formed by water. In order to see clearly the nature of the
various kinds of compounds formed by water we will begin with the most
feeble, which are determined by purely mechanical superficial properties
of the reacting substances.[12]

  [11] Of all known liquids, water exhibits the greatest _cohesion_ of
       particles. Indeed, it ascends to a greater height in capillary
       tubes than other liquids; for instance, two and a half times as
       high as alcohol, nearly three times as high as ether, and to a
       much greater height than oil of vitriol, &c. In a tube one mm. in
       diameter, water at 0° ascends 15·3 mm., measuring from the height
       of the liquid to two-thirds of the height of the meniscus, and at
       100° it rises 12·5 mm. The cohesion varies very uniformly with
       the temperature; thus at 50° the height of the capillary column
       equals 13·9 mm.--that is, the mean between the columns at 0°
       and 100°. This uniformity is not destroyed even at temperatures
       near the freezing point, and hence it may be assumed that at
       high temperatures cohesion will vary as uniformly as at ordinary
       temperatures; that is, the difference between the columns at 0°
       and 100° being 2·8 mm., the height of the column at 500° should
       be 15·2-(5 × 2·8) = 1·2 mm.; or, in other words, at these high
       temperatures the cohesion between the particles of water would he
       almost _nil_. Only certain solutions (sal ammoniac and lithium
       chloride), and these only with a great excess of water, rise
       higher than pure water in capillary tubes. The great cohesion of
       water doubtless determines many of both its physical and chemical
       properties.

       The quantity of heat required to raise the temperature of one
       part by weight of water from 0° to 1°, _i.e._ by 1° C., is called
       the _unit of heat_ or calorie; the _specific heat of liquid
       water_ at 0° is taken as equal to unity. The variation of this
       specific heat with a rise in temperature is inconsiderable in
       comparison with the variation exhibited by the specific heats of
       other liquids. According to Ettinger, the specific heat of water
       at 20° = 1·016, at 50° = 1·039, and at 100° = 1·073. The specific
       heat of water is greater than that of any other known liquid;
       for example, the specific heat of alcohol at 0° is 0·55--_i.e._
       the quantity of heat which raises 55 parts of water 1° raises
       100 parts of alcohol 1°. The specific heat of oil of turpentine
       at 0° is 0·41, of ether 0·53, of acetic acid 0·5274, of mercury
       0·033. Hence water is the best condenser or absorber of heat.
       This property of water has an important significance in practice
       and in nature. Water prevents rapid cooling or heating, and thus
       tempers cold and heat. The specific heats of ice and aqueous
       vapour are much less than that of water; namely, that of ice is
       0·504, and of steam 0·48.

       With an increase in pressure equal to one atmosphere, the
       compressibility of water (_see_ Note 9) is 0·000047, of mercury
       0·00000352, of ether 0·00012 at 0°, of alcohol at 13° 0·000095.
       The addition of various substances to water generally decreases
       both its compressibility and cohesion. The compressibility of
       other liquids increases with a rise of temperature, but for water
       it decreases up to 53° and then increases like other liquids.

       The _expansion of water_ by heat (Note 9) also exhibits many
       peculiarities which are not found in other liquids. The expansion
       of water at low temperatures is very small compared with other
       liquids; at 4° it is almost zero, and at 100° it is equal to
       0·0008; below 4° it is negative--_i.e._ water on cooling then
       expands, and does not decrease in volume. In passing into a solid
       state, the specific gravity of water decreases; at 0° one c.c.
       of water weighs 0·999887 gram, and one c.c. of ice at the same
       temperature weighs only 0·9175 gram. The ice formed, however,
       contracts on cooling like the majority of other substances. Thus
       100 volumes of ice are produced from 92 volumes of water--that
       is, water expands considerably on freezing, which fact determines
       a number of natural phenomena. The freezing point of water falls
       with an increase in pressure (0·007° per atmosphere), because in
       freezing water expands (Thomson), whilst with substances which
       contract in solidifying the melting point rises with an increase
       in pressure; thus, paraffin under one atmosphere melts at 46°,
       and under 100 atmospheres at 49°.

       When liquid water passes into vapour, the cohesion of its
       particles must be destroyed, as the particles are removed to
       such a distance from each other that their mutual attraction
       no longer exhibits any influence. As the cohesion of aqueous
       particles varies at different temperatures, the quantity of heat
       which is expended in overcoming this cohesion--or the _latent
       heat of evaporation_--will for this reason alone be different at
       different temperatures. The quantity of heat which is consumed in
       the transformation of one part by weight of water, at different
       temperatures, into vapour was determined by Regnault with great
       accuracy. His researches showed that one part by weight of water
       at 0°, in passing into vapour having a temperature _t_°, consumes
       606·5 + 0·305_t_ units of heat, at 50° 621·7, at 100° 637·0, at
       150° 652·2, and at 200° 667·5. But this quantity includes also
       the quantity of heat required for heating the water from 0° to
       _t_°--_i.e._ besides the latent heat of evaporation, also that
       heat which is used in heating the water in a liquid state to a
       temperature _t_°. On deducting this amount of heat, we obtain the
       latent heat of evaporation of water as 606·5 at 0°, 571 at 50°,
       534 at 100°, 494 at 150°, and only 453 at 200°, which shows that
       the conversion of water at different temperatures into vapour
       at a constant temperature requires very different quantities of
       heat. This is chiefly dependent on the difference of the cohesion
       of water at different temperatures; the cohesion is greater at
       low than at high temperatures, and therefore at low temperatures
       a greater quantity of heat is required to overcome the cohesion.
       On comparing these quantities of heat, it will be observed that
       they decrease rather uniformly, namely their difference between
       0° and 100° is 72, and between 100° and 200° is 81 units of
       heat. From this we may conclude that this variation will be
       approximately the same for high temperatures also, and therefore
       that no heat would be required for the conversion of water into
       vapour at a temperature of about 400°. At this temperature, water
       passes into vapour whatever be the pressure (see Chap. II. The
       absolute boiling point of water, according to Dewar, is 370°,
       the critical pressure 196 atmospheres). It must here be remarked
       that water, in presenting a greater cohesion, requires a larger
       quantity of heat for its conversion into vapour than other
       liquids. Thus alcohol consumes 208, ether 90, turpentine 70,
       units of heat in their conversion into vapour.

       The whole amount of heat which is consumed in the conversion of
       water into vapour is not used in overcoming the cohesion--that
       is, in internal accomplished in the liquid. A part of this heat
       is employed in moving the aqueous particles; in fact, aqueous
       vapour at 100° occupies a volume 1,659 times greater than that
       of water (at the ordinary pressure), consequently a portion of
       the heat or work is employed in lifting the aqueous particles, in
       overcoming pressure, or in external work, which may be usefully
       employed, and which is so employed in steam engines. In order
       to determine this work, let us consider the variation of the
       maximum _pressure_ or _vapour tension of steam_ at different
       temperatures. The observations of Regnault in this respect,
       as on those preceding, deserve special attention from their
       comprehensiveness and accuracy. The pressure or tension of
       aqueous vapour at various temperatures is given in the adjoining
       table, and is expressed in millimetres of the barometric column
       reduced to 0°.

            +------------+---------+-------------+----------+
            |Temperature | Tension | Temperature | Tension  |
            +------------+---------+-------------+----------+
            |   -20°     |    0·9  |     70°     |   233·3  |
            |   -10°     |    2·1  |     90°     |   525·4  |
            |     0°     |    4·6  |    100°     |   760·0  |
            |   +10°     |    9·1  |    105°     |   906·4  |
            |    15°     |   12·7  |    110°     |  1075·4  |
            |    20°     |   17·4  |    115°     |  1269·4  |
            |    25°     |   23·5  |    120°     |  1491·3  |
            |    30°     |   31·5  |    150°     |  3581·0  |
            |    50°     |   92·0  |    200°     | 11689·0  |
            +------------+---------+-------------+----------+

       The table shows the boiling points of water at different
       pressures. Thus on the summit of Mont Blanc, where the average
       pressure is about 424 mm., water boils at 84·4°. In a rarefied
       atmosphere water boils even at the ordinary temperature, but in
       evaporating it absorbs heat from the neighbouring parts, and
       therefore it becomes cold and may even freeze if the pressure
       does not exceed 4·6 mm., and especially if the vapour be rapidly
       absorbed as it is formed. Oil of vitriol, which absorbs the
       aqueous vapour, is used for this purpose. Thus ice may be
       obtained artificially at the ordinary temperature with the aid
       of an air-pump. This table of the tension of aqueous vapour also
       shows the temperature of water contained in a closed boiler if
       the pressure of the steam formed be known. Thus at a pressure
       of five atmospheres (a pressure of five times the ordinary
       atmospheric pressure--_i.e._ 5 × 760 = 3,800 mm.) the temperature
       of the water would be 152°. The table also shows the pressure
       produced on a given surface by steam on issuing from a boiler.
       Thus steam having a temperature of 152° exerts a pressure of
       517 kilos on a piston whose surface equals 100 sq. cm., for the
       pressure of one atmosphere on one sq. cm. equals 1,033 kilos,
       and steam at 152° has a pressure of five atmospheres. As a
       column of mercury 1 mm. high exerts a pressure of 1·35959 grams
       on a surface of 1 sq. cm., therefore the pressure of aqueous
       vapour at 0° corresponds with a pressure of 6·25 grams per
       square centimetre. The pressures for all temperatures may be
       calculated in a similar way, and it will be found that at 100°
       it is equal to 1,033·28 grams. This means that if a cylinder
       be taken whose sectional area equals 1 sq. cm., and if water
       be poured into it and it be closed by a piston weighing 1,033
       grams, then on heating it in a vacuum to 100° no steam will be
       formed, because the steam cannot overcome the pressure of the
       piston; and if at 100° 534 units of heat be transmitted to each
       unit of weight of water, then the whole of the water will be
       converted into vapour having the same temperature; and so also
       for every other temperature. The question now arises, to what
       height does the piston rise under these circumstances? that is,
       in other words, What is the volume occupied by the steam under
       a known pressure? For this we must know the weight of a cubic
       centimetre of steam at various temperatures. It has been shown by
       experiment that the density of steam, which does not saturate a
       space, varies very inconsiderably at all possible pressures, and
       is nine times the density of hydrogen under similar conditions.
       Steam which saturates a space varies in density at different
       temperatures, but this difference is very small, and its
       average density with reference to air is 0·64. We will employ
       this number in our calculation, and will calculate what volume
       the steam occupies at 100°. One cubic centimetre of air at 0°
       and 760 mm. weighs 0·001293 gram, at 100° and under the same
       pressure it will weigh 0·001293/1·368 or about 0·000946 gram,
       and consequently one cubic centimetre of steam whose density
       is 0·64 will weigh 0·000605 gram at 100°, and therefore one
       gram of aqueous vapour will occupy a volume of about 1·653 c.c.
       Consequently, the piston in the cylinder of 1 sq. cm. sectional
       area, and in which the water occupied a height of 1 cm., will
       be raised 1,653 cm. on the conversion of this water into steam.
       This piston, as has been mentioned, weighs 1,033 grams, therefore
       the _external work of the steam_--that is, that work which the
       water does in its conversion into steam at 100°--is equal to
       lifting a piston weighing 1,033 grams to a height of 1,653 cm.,
       or 17·07 kilogram-metres of work--_i.e._ is capable of lifting
       17 kilograms 1 metre, or 1 kilogram 17 metres. One gram of water
       requires for its conversion into steam 534 gram units of heat or
       0·534 kilogram unit of heat--_i.e._ the quantity of heat absorbed
       in the evaporation of one gram of water is equal to the quantity
       of heat which is capable of heating 1 kilogram of water 0·534°.
       Each unit of heat, as has been shown by accurate experiment,
       is capable of doing 424 kilogram-metres of work. Hence, in
       evaporating, one gram of water expends 424 × 0·534 = (almost) 227
       kilogram-metres of work. The external work was found to be only
       17 kilogram-metres, therefore 210 kilogram-metres are expended in
       overcoming the internal cohesion of the aqueous particles, and
       consequently about 92 p.c. of the total heat or work is consumed
       in overcoming the internal cohesion. The following figures are
       thus calculated approximately:--

       +------------+----------------+-----------------+--------------+
       |            | Total work of  |External work of |              |
       |Temperature | evaporation in |   vapour in     |   Internal   |
       |            |kilogram-metres |kilogram-metres  |work of vapour|
       +------------+----------------+-----------------+--------------+
       |     0°     |      255       |       13        |     242      |
       |    50°     |      242       |       15        |     227      |
       |   100°     |      226       |       17        |     209      |
       |   150°     |      209       |       19        |     190      |
       |   200°     |      192       |       20        |     172      |
       +------------+----------------+-----------------+--------------+

       The work necessary for overcoming the internal cohesion of
       water in its passage into vapour decreases with the rise in
       temperature--that is, corresponds with the decrease of cohesion;
       and, in fact, the variations which take place in this case are
       very similar to those which are observed in the heights to which
       water rises in capillary tubes at different temperatures. It is
       evident, therefore, that the amount of external--or, as it is
       termed, useful--work which water can supply by its evaporation
       is very small compared with the amount which it expends in its
       conversion into vapour.

       In considering certain physico-mechanical properties of water, I
       had in view not only their importance for theory and practice,
       but also their purely chemical significance; for it is evident
       from the above considerations that even in a physical change
       of state the greatest part of the work done is employed in
       overcoming cohesion, and that an enormous amount of internal
       energy must be expended in overcoming chemical cohesion or
       affinity.

  [12] When it is necessary to heat a considerable mass of liquid in
       different vessels, it would be very uneconomical to make use of
       metallic vessels and to construct a separate furnace for each;
       such cases are continually met with in practice. Steam from a
       boiler is introduced into the liquid, or, in general, into the
       vessel which it is required to heat. The steam, in condensing and
       passing into a liquid state, parts with its latent heat, and as
       this is very considerable a small quantity of steam will produce
       a considerable heating effect. If it be required, for instance,
       to heat 1,000 kilos of water from 20° to 50°, which requires
       approximately 30,000 units of heat, steam at 100° is passed into
       the water from a boiler. Each kilogram of water at 50° contains
       about 50 units of heat, and each kilogram of steam at 100°
       contains 637 units of heat; therefore, each kilogram of steam in
       cooling to 50° gives up 587 units of heat, and consequently 52
       kilos of steam are capable of heating 1,000 kilos of water from
       20° to 50°. Water is very often applied for heating in chemical
       practice. For this purpose metallic vessels or pans, called
       'water-baths,' are made use of. They are closed by a cover formed
       of concentric rings lying on each other. The vessels--such as
       beakers, evaporating basins, retorts, &c.--containing liquids,
       are placed on these rings, and the water in the bath is heated.
       The steam given off heats the bottom of the vessels to be heated,
       and thus effects the evaporation or distillation.

Water is mechanically attracted by many substances; it adheres to their
surfaces just as dust adheres to objects, or one piece of polished glass
adheres to another. Such attraction is termed 'moistening,' 'soaking,'
or 'absorption of water.' Thus water moistens clean glass and adheres
to its surface, is absorbed by the soil, sand, and clay, and does
not flow away from them, but lodges itself between their particles.
Similarly, water soaks into a sponge, cloth, hair, or paper, &c., but
fat and greasy substances in general are not moistened. Attraction
of this kind does not alter the physical or chemical properties of
water. For instance, under these circumstances water, as is known from
everyday experience, may be expelled from objects by drying. Water
which is in any way held mechanically may be dislodged by mechanical
means, by friction, pressure, centrifugal force, &c. Thus water is
squeezed from wet cloth by pressure or centrifugal machines. But objects
which in practice are called dry (because they do not feel wet) often
still contain moisture, as may be proved by heating the object in a
glass tube closed at one end. By placing a piece of paper, dry earth,
or any similar object (especially porous substances) in such a glass
tube, and heating that part of the tube where the object is situated,
it will be remarked that water condenses on the cooler portions of
the tube. The presence of such absorbed, or 'hygroscopic,' water is
generally best detected in non-volatile substances by drying them at
100°, or under the receiver of an air-pump and over substances which
attract water chemically. By weighing a substance before and after
drying, it is easy to determine the amount of hygroscopic water from
the loss in weight.[13] Only in this case the amount of water must be
judged with care, because the loss in weight may sometimes proceed
from the decomposition of the substance itself, with disengagement of
gases or vapour. In making exact weighings the hygroscopic capacity
of substances--that is, their capacity to absorb moisture--must be
continually kept in view, as otherwise the weight will be untrue from
the presence of moisture. The quantity of moisture absorbed depends on
the degree of moisture of the atmosphere (that is, on the tension of the
aqueous vapour in it) in which a substance is situated. In an entirely
dry atmosphere, or in a vacuum, the hygroscopic water is expelled, being
converted into vapour; therefore, substances containing hygroscopic
water may be completely dried by placing them in a dry atmosphere or
in a vacuum. The process is aided by heat, as it increases the tension
of the aqueous vapour. Phosphoric anhydride (a white powder), liquid
sulphuric acid, solid and porous calcium chloride, or the white powder
of ignited copper sulphate, are most generally employed in drying
gases. They absorb the moisture contained in air and all gases to a
considerable, but not unlimited, extent. Phosphoric anhydride and
calcium chloride deliquesce, become damp, sulphuric acid changes from an
oily thick liquid into a more mobile liquid, and ignited copper sulphate
becomes blue; after which changes these substances partly lose their
capacity of holding water, and can, if it be in excess, even give up
their water to the atmosphere. We may remark that the order in which
these substances are placed above corresponds with the order in which
they stand in respect to their capacity for absorbing moisture. Air
dried by calcium chloride still contains a certain amount of moisture,
which it can give up to sulphuric acid. The most complete desiccation
takes place with phosphoric anhydride. Water is also removed from
many substances by placing them in a dish over a vessel containing a
substance absorbing water under a glass bell jar.[14] The bell jar,
like the receiver of an air pump, should be hermetically closed. In this
case desiccation takes place; because sulphuric acid, for instance,
first dries the air in the bell jar by absorbing its moisture, the
substance to be dried then parts with its moisture to the dry air,
from which it is again absorbed by the sulphuric acid, &c. Desiccation
proceeds still better under the receiver of an air pump, for then the
aqueous vapour is formed more quickly than in a bell jar full of air.

  [13] [Illustration: FIG. 13.--Drying oven, composed of brazed copper.
       It is heated by a lamp. The object to be dried is placed on the
       gauze inside the oven. The thermometer indicates the temperature.]

       In order to dry any substance at about 100°--that is, at the
       boiling point of water (hygroscopic water passes off at this
       temperature)--an apparatus called a 'drying-oven' is employed.
       It consists of a double copper box; water is poured into the
       space between the internal and external boxes, and the oven is
       then heated over a stove or by any other means, or else steam
       from a boiler is passed between the walls of the two boxes.
       When the water boils, the temperature inside the inner box will
       be approximately 100° C. The substance to be dried is placed
       inside the oven, and the door is closed. Several holes are cut
       in the door to allow the free passage of air, which carries
       off the aqueous vapour by the chimney on the top of the oven.
       Often, however, desiccation is carried on in copper ovens
       heated directly over a lamp (fig. 13). In this case any desired
       temperature may be obtained, which is determined by a thermometer
       fixed in a special orifice. There are substances which only
       part with their water at a much higher temperature than 100°,
       and then such air baths are very useful. In order to determine
       directly the amount of water in a substance which does not part
       with anything except water at a red heat, the substance is placed
       in a bulb tube. By first weighing the tube empty and then with
       the substance to be dried in it, the weight of the substance
       taken may be found. The tube is then connected on one side with
       a gas-holder full of air, which, on opening a stop-cock, passes
       first through a flask containing sulphuric acid, and then into a
       vessel containing lumps of pumice stone moistened with sulphuric
       acid. In passing through these vessels the air is thoroughly
       dried, having given up all its moisture to the sulphuric acid.
       Thus dry air will pass into the bulb tube, and as hygroscopic
       water is entirely given up from a substance in dry air even at
       the ordinary temperature, and still more rapidly on heating,
       the moisture given up by the substance in the tube will be
       carried off by the air passing through it. This damp air then
       passes through a U-shaped tube full of pieces of pumice stone
       moistened with sulphuric acid, which absorbs all the moisture
       given off from the substance in the bulb tube. Thus all the water
       expelled from the substance will collect in the [U] tube, and so,
       if this be weighed before and after, the difference will show the
       quantity of water expelled from the substance. If only water (and
       not any gases) come over, the increase of the weight of the [U]
       tube will be equal to the decrease in the weight of the bulb tube.

  [14] Instead of under a glass bell jar, drying over sulphuric acid is
       often carried on in a desiccator consisting of a shallow
       wide-mouthed glass vessel, closed by a well-fitting ground-glass
       cover. Sulphuric acid is poured over the bottom of the
       desiccator, and the substance to be dried is placed on a glass
       stand above the acid. A lateral glass tube with a stop-cock is
       often fused into the desiccator in order to connect it with an
       air pump, and so allow drying under a diminished pressure, when
       the moisture evaporates more rapidly. The fact that in the usual
       form of desiccator the desiccating substance (sulphuric acid) is
       placed beneath the substance to be dried has the disadvantage
       that the moist air being lighter than dry air distributes itself
       in the upper portion of the desiccator and not below. Hempel,
       in his desiccator (1891), avoids this by placing the absorbent
       above the substance to be dried. The process of desiccation
       can be further accelerated by cooling the upper portion of the
       desiccator, and so inducing ascending and descending currents of
       air within the apparatus.

From what has been said above, it is evident that the transference of
moisture to gases and the absorption of hygroscopic moisture present
great resemblance to, but still are not, chemical combinations with
water. Water, when combined as hygroscopic water, does not lose its
properties and does not form new substances.[15]

  [15] Chappuis, however, determined that in wetting 1 gram of charcoal
       with water 7 units of heat are evolved, and on pouring carbon
       bisulphide over 1 gram of charcoal as much as 24 units of heat
       are evolved. Alumina (1 gram), when moistened with water, evolves
       2-1/2 calories. This indicates that in respect to evolution of
       heat moistening already presents a transition towards exothermal
       combinations (those evolving heat in their formation).

The attraction of water for substances which dissolve in it is of a
different character. In the solution of substances in water there
proceeds a peculiar kind of indefinite combination; a new homogeneous
substance is formed from the two substances taken. But here also the
bond connecting the substances is very unstable. Water containing
different substances in solution boils at a temperature near to its
usual boiling point. From the solution of substances which are lighter
than water itself, there are obtained solutions of a less density than
water--as, for example, in the solution of alcohol in water; whilst a
heavier substance in dissolving in water gives it a higher specific
gravity. Thus salt water is heavier than fresh.[16]

  [16] Strong acetic acid (C_{2}H_{4}O_{2}), whose specific gravity at
       15° is 1·055, does not become lighter on the addition of water
       (a lighter substance, sp. gr. = 0·999), but heavier, so that a
       solution of 80 parts of acetic acid and 20 parts of water has a
       specific gravity of 1·074, and even a solution of equal parts of
       acetic acid and water (50 p.c.) has a sp. gr. of 1·065, which is
       still greater than that of acetic acid itself. This shows the
       high degree of contraction which takes place on solution. In
       fact, solutions--and, in general, liquids--on mixing with water,
       decrease in volume.

We will consider _aqueous solutions_ somewhat fully, because, among
other reasons, solutions are constantly being formed on the earth and
in the waters of the earth, in plants and in animals, in chemical
processes and in the arts, and these solutions play an important part
in the chemical transformations which are everywhere taking place,
not only because water is everywhere met with, but chiefly because a
substance in solution presents the most favourable conditions for the
process of chemical changes, which require a mobility of parts and a
possible distension of parts. In dissolving, a solid substance acquires
a mobility of parts, and a gas loses its elasticity, and therefore
reactions often take place in solutions which do not proceed in the
undissolved substances. Further, a substance, distributed in water,
evidently breaks up--that is, becomes more like a gas and acquires a
greater mobility of parts. All these considerations require that in
describing the properties of substances, particular attention should be
paid to their relation to water as a solvent.

[Illustration: FIG. 14.--Method of transferring a gas into a cylinder
filled with mercury and whose open end is immersed under the mercury
in a bath having two glass sides. The apparatus containing the gas is
represented on the right. Its upper extremity is furnished with a tube
extending under the cylinder. The lower part of the vessel communicates
with a vertical tube. If mercury be poured into this tube, the pressure
of the gas in the apparatus is increased, and it passes through the
gas-conducting tube into the cylinder, where it displaces the mercury,
and can be measured or subjected to the action of absorbing agents, such
as water.]

It is well known that water dissolves many substances. Salt, sugar,
alcohol, and a number of other substances, dissolve in water and form
homogeneous liquids with it. To demonstrate the solubility of gases
in water, a gas should be taken which has a high co-efficient of
solubility--for instance, ammonia. This is introduced into a bell jar
(or cylinder, as in fig. 14), which is previously filled with mercury
and stands in a mercury bath. If water be then introduced into the
cylinder, the mercury will rise, owing to the water dissolving the
ammonia gas. If the column of mercury be less than the barometric
column, and if there be sufficient water to dissolve the gas, all the
ammonia will be absorbed by the water. The water is introduced into
the cylinder by a glass pipette, with a bent end. The bent end is put
into water, and the air is sucked out from the upper end. When full of
water, its upper end is closed with the finger, and the bent end placed
in the mercury bath under the orifice of the cylinder. On blowing into
the pipette the water will rise to the surface of the mercury in the
cylinder owing to its lightness. The solubility of a gas like ammonia
may be demonstrated by taking a flask full of the gas, and closed
by a cork with a tube passing through it. On placing the tube under
water, the water will rise into the flask (this may be accelerated by
previously warming the flask), and begin to play like a fountain inside
it. Both the rising of the mercury and the fountain clearly show the
considerable affinity of water for ammonia gas, and the force acting
in this dissolution is rendered evident. A certain period of time is
required both for the homogeneous intermixture of gases (diffusion) and
the process of solution, which depends, not only on the surface of the
participating substances, but also on their nature. This is seen from
experiment. Solutions of different substances heavier than water, such
as salt or sugar, are poured into tall jars. Pure water is then very
carefully poured into these jars (through a funnel) on to the top of the
solutions, so as not to disturb the lower stratum, and the jars are then
left undisturbed. The line of demarcation between the solution and the
pure water will be visible, owing to their different co-efficients of
refraction. Notwithstanding that the solutions taken are heavier than
water, after some time complete intermixture will ensue. Gay Lussac
convinced himself of this fact by this particular experiment, which
he conducted in the cellars under the Paris Astronomical Observatory.
These cellars are well known as the locality where numerous interesting
researches have been conducted, because, owing to their depth under
ground, they have a uniform temperature during the whole year; the
temperature does not change during the day, and this was indispensable
for the experiments on the diffusion of solutions, in order that no
doubt as to the results should arise from a daily change of temperature
(the experiment lasted several months), which would set up currents in
the liquids and intermix their strata. Notwithstanding the uniformity
of the temperature, the substance in solution in time ascended into the
water and distributed itself uniformly through it, proving that there
exists between water and a substance dissolved in it a particular kind
of attraction or striving for mutual interpenetration in opposition to
the force of gravity. Further, this effort, or rate of diffusion, is
different for salt or sugar or for various other substances.[16 bis] It
follows therefore that a peculiar force acts in solution, as in actual
chemical combinations, and solution is determined by a particular kind
of motion (by the chemical energy of a substance) which is proper to the
substance dissolved and to the solvent.

  [16 bis] Graham, in the jelly formed by gelatine, and De Vries in
       gelatinous silica (Chapter XVIII.) most frequently employed
        (tinted) substances, for instance, K_{2}Cr_{2}O_{7},
       which showed the rate of diffusion with very great clearness.
       Prof. Oumoff employed the method described in Chapter X., Note
       17, for this purpose.

Graham made a series of experiments similar to those above described,
and showed that the _rate of diffusion_[17] in water is very
variable--that is, a uniform distribution of a substance in the water
dissolving it is attained in different periods of time with different
solutions. Graham compared diffusive capacity with volatility. There are
substances which diffuse easily, and there are others which diffuse with
difficulty, just as there are more or less volatile substances. Seven
hundred cubic centimetres of water were poured into a jar, and by means
of a syphon (or a pipette) 100 cub. centimetres of a solution containing
10 grams of a substance were cautiously poured in so as to occupy the
lower portion of the jar. After a lapse of several days, successive
layers of 50 cubic centimetres were taken from the top downwards, and
the quantity of substance dissolved in the different layers determined.
Thus, common table salt, after fourteen days, gave the following amounts
(in milligrams) in the respective layers, beginning from the top: 104,
120, 126, 198, 267, 340, 429, 535, 654, 766, 881, 991, 1,090, 1,187,
and 2,266 in the remainder; whilst albumin in the same time gave, in
the first seven layers, a very small amount, and beginning from the
eighth layer, 10, 15, 47, 113, 343, 855, 1,892, and in the remainder
6,725 milligrams. Thus, the diffusive power of a solution depends on
time and on the nature of the substance dissolved, which fact may
serve, not only for explaining the process of solution, but also for
distinguishing one substance from another. Graham showed that substances
which rapidly diffuse through liquids are able to rapidly pass through
membranes and crystallise, whilst substances which diffuse slowly and do
not crystallise are _colloids_, that is, resemble glue, and penetrate
through a membrane slowly, and form jellies; that is, occur in insoluble
forms,[18] as will be explained in speaking of silica.

  [17] The researches of Graham, Fick, Nernst, and others showed that
       the quantity of a dissolved substance which is transmitted
       (rises) from one stratum of liquid to another in a vertical
       cylindrical vessel is not only proportional to the time and to
       the sectional area of the cylinder, but also to the amount and
       nature of the substance dissolved in a stratum of liquid, so that
       each substance has its corresponding co-efficient of diffusion.
       The cause of the diffusion of solutions must be considered
       as essentially the same as the cause of the diffusion of
       gases--that is, as dependent on motions which are proper to their
       molecules; but here most probably those purely chemical, although
       feebly-developed, forces, which incline the substances dissolved
       to the formation of definite compounds, also play their part.

  [18] [Illustration: FIG. 15.--Dialyser. Apparatus for the separation
       of substances which pass through a membrane from those which do
       not. Description in text.]

       The rate of diffusion--like the rate of transmission--through
       membranes, or _dialysis_ (which plays an important part in the
       vital processes of organisms and also in technical processes),
       presents, according to Graham's researches, a sharply defined
       change in passing from such crystallisable substances as the
       majority of salts and acids to substances which are capable
       of giving jellies (gum, gelatin, &c.) The former diffuse
       into solutions and pass through membranes much more rapidly
       than the latter, and Graham therefore distinguishes between
       _crystalloids_, which diffuse rapidly, and _colloids_, which
       diffuse slowly. On breaking solid colloids into pieces, a total
       absence of cleavage is remarked. The fracture of such substances
       is like that of glue or glass. It is termed a 'conchoidal'
       fracture. Almost all the substances of which animal and vegetable
       bodies consist are colloids, and this is, at all events, partly
       the reason why animals and plants have such varied forms,
       which have no resemblance to the crystalline forms of the
       majority of mineral substances. The colloid solid substances in
       organisms--that is, in animals and plants--almost always contain
       water, and take most peculiar forms, of networks, of granules,
       of hairs, of mucous, shapeless masses, &c., which are quite
       different from the forms taken by crystalline substances. When
       colloids separate out from solutions, or from a molten state,
       they present a form which is similar to that of the liquid from
       which they were formed. Glass may he taken as the best example of
       this. Colloids are distinguishable from crystalloids, not only
       by the absence of crystalline form, but by many other properties
       which admit of clearly distinguishing both these classes of
       solids, as Graham showed. Nearly all colloids are capable of
       passing, under certain circumstances, from a soluble into an
       insoluble state. The best example is shown by white of eggs
       (albumin) in the raw and soluble form, and in the hard-boiled
       and insoluble form. The majority of colloids, on passing into
       an insoluble form in the presence of water, give substances
       having a gelatinous appearance, which is familiar to every one
       in starch, solidified glue, jelly, &c. Thus gelatin, or common
       carpenter's glue, when soaked in water, swells up into an
       insoluble jelly. If this jelly be heated, it melts, and is then
       soluble in water, but on cooling it again forms a jelly which
       is insoluble in water. One of the properties which distinguish
       colloids from crystalloids is that the former pass very slowly
       through a membrane, whilst the latter penetrate very rapidly.
       This may be shown by taking a cylinder, open at both ends, and
       by covering its lower end with a bladder or with vegetable
       parchment (unsized paper immersed for two or three minutes in a
       mixture of sulphuric acid and half its volume of water, and then
       washed), or any other membranous substance (all such substances
       are themselves colloids in an insoluble form). The membrane must
       be firmly tied to the cylinder, so as not to leave any opening.
       Such an apparatus is called a _dialyser_ (fig. 15), and the
       process of separation of crystalloids from colloids by means of
       such a membrane is termed _dialysis_. An aqueous solution of a
       crystalloid or colloid, or a mixture of both, is poured into the
       dialyser, which is then placed in a vessel containing water, so
       that the bottom of the membrane is covered with water. Then,
       after a certain period of time, the crystalloid passes through
       the membrane, whilst the colloid, if it does pass through at all,
       does so at an incomparably slower rate. The crystalloid naturally
       passes through into the water until the solution attains the same
       strength on both sides of the membrane. By replacing the outside
       water with fresh water, a fresh quantity of the crystalloid may
       be separated from the dialyser. While a crystalloid is passing
       through the membrane, a colloid remains almost entirely in the
       dialyser, and therefore a mixed solution of these two kinds of
       substances may be separated from each other by a dialyser. The
       study of the properties of colloids, and of the phenomena of
       their passage through membranes, should elucidate much respecting
       the phenomena which are accomplished in organisms.

Hence, if it be desired to increase the rate of solution, recourse must
be had to stirring, shaking, or some such mechanical motion. But if once
a uniform solution is formed, it will remain uniform, no matter how
heavy the dissolved substance is, or how long the solution be left at
rest, which fact again shows the presence of a force holding together
the particles of the body dissolved and of the solvent.[19]

  [19] The formation of solutions may be considered in two aspects,
       from a physical and from a chemical point of view, and it is
       more evident in solutions than in any other department of
       chemistry how closely these provinces of natural science are
       allied together. On the one hand solutions form a particular
       case of a physico-mechanical interpenetration of homogeneous
       substances, and a juxtaposition of the molecules of the substance
       dissolved and of the solvent, similar to the juxtaposition which
       is exhibited in homogeneous substances. From this point of view
       this diffusion of solutions is exactly similar to the diffusion
       of gases, with only this difference, that the nature and store of
       energy are different in gases from what they are in liquids, and
       that in liquids there is considerable friction, whilst in gases
       there is comparatively little. The penetration of a dissolved
       substance into water is likened to evaporation, and solution to
       the formation of vapour. This resemblance was clearly expressed
       even by Graham. In recent years the Dutch chemist, Van't Hoff,
       has developed this view of solutions in great detail, having
       shown (in a memoir in the _Transactions of the Swedish Academy
       of Science_, Part 21, No. 17, 'Lois de l'équilibre chimique
       dans l'état dilué, gazeux ou dissous,' 1886), that for dilute
       solutions the _osmotic pressure_ follows the same laws of Boyle,
       Mariotte, Gay-Lussac, and Avogadro-Gerhardt as for gases. The
       osmotic pressure of a substance dissolved in water is determined
       by means of membranes which allow the passage of water, but not
       of a substance dissolved in it, through them. This property is
       found in animal protoplasmic membranes and in porous substances
       covered with an amorphous precipitate, such as is obtained by the
       action of copper sulphate on potassium ferrocyanide (Pfeffer,
       Traube). If, for instance, a one p.c. solution of sugar he placed
       in such a vessel, which is then closed and placed in water, the
       water passes through the walls of the vessel and increases the
       pressure by 50 mm. of the barometric column. If the pressure be
       artificially increased inside the vessel, then the water will be
       expelled through the walls. De Vries found a convenient means
       of determining _isotonic_ solutions (those presenting a similar
       osmotic pressure) in the cells of plants. For this purpose a
       portion of the soft part of the leaves of the _Tradescantis
       discolor_, for instance, is cut away and moistened with the
       solution of a given salt and of a given strength. If the osmotic
       pressure of the solution taken be less than that of the sap
       contained in the cells they will change their form or shrink; if,
       on the other hand, the osmotic pressure be greater than that of
       the sap, then the cells will expand, as can easily be seen under
       the microscope. By altering the amount of the different salts in
       solution it is possible to find for each salt the strength of
       solution at which the cells begin to swell, and at which they
       will consequently have an equal osmotic pressure. As it increases
       in proportion to the amount of a substance dissolved per 100
       parts of water, it is possible, knowing the osmotic pressure
       of a given substance--for instance, sugar at various degrees
       of concentration of solution--and knowing the composition of
       isotonic solutions compared with sugar, to determine the osmotic
       pressure of all the salts investigated. The osmotic pressure of
       dilute solutions determined in this manner directly or indirectly
       (from observations made by Pfeffer and De Vries) was shown to
       follow the same laws as those of the pressure of gases; for
       instance, by doubling or increasing the quantity of a salt (in
       a given volume) _n_ times, the pressure is doubled or increases
       _n_ times. So, for example, in a solution containing one part
       of sugar per 100 parts of water the osmotic pressure (according
       to Pfeffer) = 58·5 cm. of mercury, if 2 parts of sugar = 101·6,
       if 4 parts = 208·2 and so on, which proves that the ratio is
       true within the limits of experimental error. (2) Different
       substances for equal strengths of solutions, show very different
       osmotic pressures, just as gases for equal parts by weight in
       equal volumes show different tensions. (3) If, for a given dilute
       solution at 0°, the osmotic pressure equal _p_°, then at _t_°
       it will be greater and equal to _p_°(1 + 0·00367_t_), _i.e._ it
       increases with the temperature in exactly the same manner as the
       tension of gases increases. (4) If in dilute solutions of such
       substances as do not conduct an electric current (for instance,
       sugar, acetone, and many other organic bodies) the substances
       be taken in the ratio of their molecular weights (expressed by
       their formulæ, see Chapter VII.), then not only will the osmotic
       pressure be equal, but its magnitude will be determined by
       that tension which would be proper to the vapours of the given
       substances when they would be contained in the space occupied by
       the solution, just as the tension of the vapours of molecular
       quantities of the given substances will be equal, and determined
       by the laws of Gay-Lussac, Mariotte, and Avogadro-Gerhardt. Those
       formulæ (Chapter VII., Notes 23 and 24) by which the gaseous
       state of matter is determined, may also be applied in the present
       case. So, for example, the osmotic pressure _p_, in centimetres
       of mercury, of a one per cent. solution of sugar, may be
       calculated according to the formula for gases:

                       M_p_ = 6200_s_(273 + _t_),

       where M is the molecular weight, _s_ the weight in grams of a
       cubic centimetre of vapour, and _t_ its temperature. For sugar M
       = 342 (because its molecular composition is C_{12}H_{22}O_{11}).
       The specific gravity of the solution of sugar is 1·003, hence the
       weight of sugar _s_ contained in a 1 per cent. solution = 0·01003
       gram. The observation was made at _t_ = 14°. Hence, according
       to the formula, we find _p_ = 52·2 centimetres. And experiments
       carried on at 14° gave 53·5 centimetres, which is very near to
       the above. (5) For the solutions of salts, acids, and similar
       substances, which conduct an electric current, the calculated
       pressure is usually (but not always in a definite or multiple
       number of times) less than the observed by _i_ times, and this
       _i_ for dilute solutions of MgSO_{4} is nearly 1, for CO_{2} = 1,
       for KCl, NaCl, KI, KNO_{3} greater than 1, and approximates to 2,
       for BaCl_{2}, MgCl_{2}, K_{2}CO_{3}, and others between 2 and 3,
       for HCl, H_{2}SO_{4}, NaNO_{3}, CaN_{2}O_{6}, and others nearly
       2 and so on. It should be remarked that the above deductions are
       only applicable (and with a certain degree of accuracy) to dilute
       solutions, and in this respect resemble the generalisations of
       Michel and Kraft (see Note 44). Nevertheless, the arithmetical
       relation found by Van't Hoff between the formation of vapours
       and the transition into dilute solutions forms an important
       scientific discovery, which should facilitate the explanation of
       the nature of solutions, while the osmotic pressure of solutions
       already forms a very important aspect of the study of solutions.
       In this respect it is necessary to mention that Prof. Konovaloff
       (1891, and subsequently others also) discovered the dependence
       (and it may be a sufficient explanation) of the osmotic pressure
       upon the differences of the tensions of aqueous vapours and
       aqueous solutions; this, however, already enters into a special
       province of physical chemistry (certain data are given in Note
       49 and following), and to this physical side of the question
       also belongs one of the extreme consequences of the resemblance
       of osmotic pressure to gaseous pressure, which is that the
       concentration of a uniform solution varies in parts which are
       heated or cooled. Soret (1881) indeed observed that a solution
       of copper sulphate containing 17 parts of the salt at 20° only
       contained 14 parts after heating the upper portion of the tube
       to 80° for a long period of time. This aspect of solution, which
       is now being very carefully and fully worked out, may be called
       the _physical_ side. Its other aspect is purely _chemical_, for
       solution does not take place between any two substances, but
       requires a special and particular attraction or affinity between
       them. A vapour or gas permeates any other vapour or gas, but a
       salt which dissolves in water may not be in the least soluble
       in alcohol, and is quite insoluble in mercury. In considering
       solutions as a manifestation of chemical force (and of chemical
       energy), it must be acknowledged that they are here developed to
       so feeble an extent that the definite compounds (that is, those
       formed according to the law of multiple proportions) formed
       between water and a soluble substance dissociate even at the
       ordinary temperature, forming a homogeneous system--that is,
       one in which both the compound and the products into which it
       decomposes (water and the aqueous compound) occur in a liquid
       state. The chief difficulty in the comprehension of solutions
       depends on the fact that the mechanical theory of the structure
       of liquids has not yet been so fully developed as the theory of
       gases, and solutions are liquids. The conception of solutions
       as liquid dissociated definite chemical compounds is based on
       the following considerations: (1) that there exist certain
       undoubtedly definite chemical crystallised compounds (such as
       H_{2}SO_{4},H_{2}O; or NaCl,2H_{2}O; or CaCl_{2},6H_{2}O; &c.)
       which melt on a certain rise of temperature, and then form true
       solutions; (2) that metallic alloys in a molten condition are
       real solutions, but on cooling they often give entirely distinct
       and definite crystallised compounds, which are recognised by
       the properties of alloys; (3) that between the solvent and the
       substance dissolved there are formed, in a number of cases,
       many undoubtedly definite compounds, such as compounds with
       water of crystallisation; (4) that the physical properties of
       solutions, and especially their specific gravities (a property
       which can be very accurately determined), vary with a change in
       composition, and in such a manner as would be required by the
       formation of one or more definite but dissociating compounds.
       Thus, for example, on adding water to fuming sulphuric acid its
       density is observed to decrease until it attains the definite
       composition H_{2}SO_{4}, or SO_{3} + H_{2}O, when the specific
       gravity increases, although on further diluting with water it
       again falls. Moreover (Mendeléeff, _The Investigation of Aqueous
       Solutions from their Specific Gravities_, 1887), the increase
       in specific gravity (_ds_), varies in all well-known solutions
       with the proportion of the substance dissolved (_dp_), and this
       dependence can be expressed by a formula (_ds_/_dp_ = A + B_p_)
       between the limits of definite compounds whose existence in
       solutions must be admitted, and this is in complete accordance
       with the dissociation hypothesis. Thus, for instance, from
       H_{2}SO_{4} to H_{2}SO_{4} + H_{2}O (both these substances exist
       as definite compounds in a free state), the fraction _ds_/_dp_
       = 0·0729-0·000749_p_ (where _p_ is the percentage amount of
       H_{2}SO_{4}). For alcohol C_{2}H_{6}O, whose aqueous solutions
       have been more accurately investigated than all others, the
       definite compound C_{2}H_{6}O + 3H_{2}O, and others must be
       acknowledged in its solutions.

       The two aspects of solution above mentioned, and the hypotheses
       which have as yet been applied to the examination of solutions,
       although they have somewhat different starting points, will
       doubtless in time lead to a general theory of solutions,
       because the same common laws govern both physical and chemical
       phenomena, inasmuch as the properties and motions of molecules,
       which determine physical properties, depend on the motions and
       properties of atoms, which determine chemical reactions. For
       details of the questions dealing with theories of solution,
       recourse must now be had to special memoirs and to works on
       theoretical (physical) chemistry; for this subject forms one of
       special interest at the present epoch of the development of our
       science. In working out chiefly the chemical side of solutions, I
       consider it to be necessary to reconcile the two aspects of the
       question; this seems to me to be all the more possible, as the
       physical side is limited to dilute solutions only, whilst the
       chemical side deals mainly with strong solutions.

In the consideration of the process of solution, besides the conception
of diffusion, another fundamental conception is necessary--namely, that
of the _saturation of solutions_.

Just as moist air may be diluted with any desired quantity of dry air,
so also an indefinitely large quantity of a liquid solvent may be
taken, and yet a uniform solution will be obtained. But more than a
definite quantity of aqueous vapour cannot be introduced into a certain
volume of air at a certain temperature. The excess above the point of
saturation will remain in the liquid state.[20] The relation between
water and substances dissolved in it is similar. More than a definite
quantity of a substance cannot, at a certain temperature, dissolve in
a given quantity of water; the excess does not unite with the water.
Just as air or a gas becomes saturated with vapour, so water becomes
saturated with a substance dissolved in it. If an excess of a substance
be added to water which is already saturated with it, it will remain in
its original state, and will not diffuse through the water. The quantity
of a substance (either by volume with gases, or by weight with solids
and liquids) which is capable of saturating 100 parts of water is called
the _co-efficient of solubility_ or the _solubility_. In 100 grams of
water at 15°, there can be dissolved not more than 35·86 grams of
common salt. Consequently, its solubility at 15° is equal to 35·86.[21]
It is most important to turn attention to the _existence of the solid
insoluble substances of nature_, because on them depends the shape
of the substances of the earth's surface, and of plants and animals.
There is so much water on the earth's surface, that were the surface of
substances formed of soluble matters it would constantly change, and
however substantial their forms might be, mountains, river banks and sea
shores, plants and animals, or the habitations and coverings of men,
could not exist for any length of time.[22]

  [20] A system of (chemically or physically) re-acting substances in
       different states of aggregation--for instance, some solid, others
       liquid or gaseous--is termed a heterogeneous system. Up to now it
       is only systems of this kind which can be subjected to detailed
       examination in the sense of the mechanical theory of matter.
       Solutions (_i.e._ unsaturated ones) form fluid homogeneous
       systems, which at the present time can only be investigated with
       difficulty.

       In the case of limited solution of liquids in liquids, _the
       difference between the solvent and the substance dissolved_ is
       clearly seen. The former (that is, the solvent) may be added in
       an unlimited quantity, and yet the solution obtained will always
       be uniform, whilst only a definite saturating proportion of the
       substance dissolved can be taken, We will take water and common
       (sulphuric) ether. On shaking the ether with the water, it will
       be remarked that a portion of it dissolves in the water. If the
       ether be taken in such a quantity that it saturates the water
       and a portion of it remains undissolved, then this remaining
       portion will act as a solvent, and water will diffuse through it
       and also form a saturated solution of water in the ether taken.
       Thus two saturated solutions will be obtained. One solution
       will contain ether dissolved in water, and the other solution
       will contain water dissolved in ether. These two solutions will
       arrange themselves in two layers, according to their density;
       the ethereal solution of water will be on the top. If the upper
       ethereal solution be poured off from the aqueous solution,
       any quantity of ether may be added to it; this shows that the
       dissolving substance is ether. If water be added to it, it is
       no longer dissolved in it; this shows that water saturates the
       ether--here water is the substance dissolved. If we act in the
       same manner with the lower layer, we shall find that water is the
       solvent and ether the substance dissolved. By taking different
       amounts of ether and water, the degree of solubility of ether in
       water, and of water in ether, may be easily determined. Water
       approximately dissolves 1/10 of its volume of ether, and ether
       dissolves a very small quantity of water. Let us now imagine that
       the liquid poured in dissolves a considerable amount of water,
       and that water dissolves a considerable amount of the liquid. Two
       layers could not be formed, because the saturated solutions would
       resemble each other, and therefore they would intermix in all
       proportions. This is, consequently, a case of a phenomenon where
       two liquids present considerable co-efficients of solubility
       in each other, but where it is impossible to say what these
       co-efficients are, because it is impossible to obtain a saturated
       solution.

  [21] The solubility, or co-efficient of solubility, of a substance is
       determined by various methods. Either a solution is expressly
       prepared with a clear excess of the soluble substance and
       saturated at a given temperature, and the quantity of water and
       of the substance dissolved in it determined by evaporation,
       desiccation, or other means; or else, as is done with gases,
       definite quantities of water and of the soluble substance are
       taken and the amount remaining undissolved is determined.

       [Illustration: FIG. 16.--Bunsen's absorptiometer. Apparatus for
       determining the solubility of gases in liquids.]

       The solubility of a gas in water is determined by means of an
       apparatus called an _absorptiometer_ (fig. 16). It consists of an
       iron stand _f_, on which an india-rubber ring rests. A wide
       glass tube is placed on this ring, and is pressed down on it by
       the ring _h_ and the screws _i i_. The tube is thus firmly fixed
       on the stand. A cock _r_, communicating with a funnel _r_, passes
       into the lower part of the stand. Mercury can be poured into the
       wide tube through this funnel, which is therefore made of steel,
       as copper would be affected by the mercury. The upper ring _h_
       is furnished with a cover _p_, which can be firmly pressed down
       on to the wide tube, and hermetically closes it by means of an
       india-rubber ring. The tube _r r_ can be raised at will, and
       so by pouring mercury into the funnel the height of the column
       of mercury, which produces pressure inside the apparatus, can
       be increased. The pressure can also be diminished at will, by
       letting mercury out through the cock _r_. A graduated tube _e_,
       containing the gas and liquid to be experimented on, is placed
       inside the wide tube. This tube is graduated in millimetres
       for determining the pressure, and it is calibrated for volume,
       so that the number of volumes occupied by the gas and liquid
       dissolving it can be easily calculated. This tube can also be
       easily removed from the apparatus. The lower portion of this
       tube when removed from the apparatus is shown to the right of
       the figure. It will be observed that its lower end is furnished
       with a male screw _b_, fitting in a nut _a_. The lower surface of
       the nut _a_ is covered with india-rubber, so that on screwing up
       the tube its lower end presses upon the india-rubber, and thus
       hermetically closes the whole tube, for its upper end is fused
       up. The nut _a_ is furnished with arms _c c_, and in the stand
       _f_ there are corresponding spaces, so that when the screwed-up
       internal tube is fixed into stand _f_, the arms _c c_ fix into
       these spaces cut in _f_. This enables the internal tube to be
       fixed on to the stand _f_. When the internal tube is fixed in the
       stand, the wide tube is put into its right position, and mercury
       and water are poured into the space between the two tubes, and
       communication is opened between the inside of the tube _e_ and
       the mercury between the interior and exterior tubes. This is done
       by either revolving the interior tube _e_, or by a key turning
       the nut about the bottom part of _f_. The tube _e_ is filled with
       gas and water as follows: the tube is removed from the apparatus,
       filled with mercury, and the gas to be experimented on is passed
       into it. The volume of the gas is measured, the temperature and
       pressure determined, and the volume it would occupy at 0° and
       760 mm. calculated. A known volume of water is then introduced
       into the tube. The water must be previously boiled, so as to be
       quite freed from air in solution. The tube is then closed by
       screwing it down on to the india-rubber on the nut. It is then
       fixed on to the stand _f_, mercury and water are poured into
       the intervening space between it and the exterior tube, which
       is then screwed up and closed by the cover _p_, and the whole
       apparatus is left at rest for some time, so that the tube _e_,
       and the gas in it, may attain the same temperature as that of
       the surrounding water, which is marked by a thermometer _k_
       tied to the tube _e_. The interior tube is then again closed by
       turning it in the nut, the cover _p_ again shut, and the whole
       apparatus is shaken in order that the gas in the tube _e_ may
       entirely saturate the water. After several shakings, the tube
       _e_ is again opened by turning it in the nut, and the apparatus
       is left at rest for a certain time; it is then closed and again
       shaken, and so on until the volume of gas does not diminish after
       a fresh shaking--that is, until saturation ensues. Observations
       are then made of the temperature, the height of the mercury in
       the interior tube, and the level of the water in it, and also
       of the level of the mercury and water in the exterior tube. All
       these data are necessary in order to calculate the pressure under
       which the solution of the gas takes place, and what volume of gas
       remains undissolved, and also the quantity of water which serves
       as the solvent. By varying the temperature of the surrounding
       water, the amount of gas dissolved at various temperatures may
       be determined. Bunsen, Carius, and many others determined the
       solution of various gases in water, alcohol, and certain other
       liquids, by means of this apparatus. If in a determination of
       this kind it is found that _n_ cubic centimetres of water at
       a pressure _h_ dissolve _m_ cubic centimetres of a given gas,
       measured at 0° and 760 mm., when the temperature under which
       solution took place was _t_°, then it follows that at the
       temperature _t the co-efficient of solubility of the gas_ in 1
       volume of the liquid will be equal to _m_/_n_ × 760/_h_.

       This formula is very clearly understood from the fact that the
       co-efficient of solubility of gases is that quantity measured at
       0° and 760 mm., which is absorbed at a pressure of 760 mm. by one
       volume of a liquid. If _n_ cubic centimetres of water absorb _m_
       cubic centimetres of a gas, then one cubic centimetre absorbs
       _m_/_n_. If _m_/_n_ c.c. of a gas are absorbed under a pressure
       of _h_ mm., then, according to the law of the variation of
       solubility of a gas with the pressure, there would he dissolved,
       under a pressure of 760 mm., a quantity varying in the same ratio
       to _m_/_n_ as 760 : _h_. In determining the residual volume of gas
       its moisture (note 1) must be taken into consideration.

       Below are given the number of grams of several substances
       saturating 100 grams of water--that is, their co-efficients of
       solubility by weight at three different temperatures:--

       +----------------------------------------------+--------+---------+
       |                                              |        |         |
       |                                        At 0° | At 20° | At 100° |
       +----------------------------------------------+--------+---------+
       |        {Oxygen, O_{2}                 6/1000 | 4/1000 |    --   |
       |Gases   {Carbonic anhydride, CO_{2}    35/100 | 18/100 |    --   |
       |        {Ammonia, NH_{3}                 90·0 |   51·8 |    7·3  |
       |        {Phenol, C_{6}H_{6}O              4·9 |    5·2 |   [oo]  |
       |Liquids {Amyl alcohol, C_{5}H_{12}O       4·4 |    2·9 |    --   |
       |        {Sulphuric acid, H_{2}SO_{4}     [oo] |   [oo] |   [oo]  |
       |        {Gypsum, CaSO_{4},2H_{2}O         1/5 |    1/4 |    1/5  |
       |        {Alum, AlKS_{2}O_{8},12H_{2}O     3·3 |   15·4 |  357·5  |
       |Solids  {Anhydrous sodium sulphate,       4·5 |   20   |   43    |
       |        {  Na_{2}SO_{4}                       |        |         |
       |        {Common Salt, NaCl               35·7 |   36·0 |   39·7  |
       |        {Nitre, KNO_{3}                  13·3 |   31·7 |  246·0  |
       +----------------------------------------------+--------+---------+

       Sometimes a substance is so slightly soluble that it may be
       considered as insoluble. Many such substances are met with both
       in solids and liquids, and such a gas as oxygen, although it
       does dissolve, does so in so small a proportion by weight that
       it might be considered as zero did not the solubility of even
       so little oxygen play an important part in nature (as in the
       respiration of fishes) and were not an infinitesimal quantity
       of a gas by weight so easily measured by volume. The sign [oo],
       which stands on a line with sulphuric acid in the above table,
       indicates that it intermixes with water in all proportions.
       There are many such cases among liquids, and everybody knows,
       for instance, that spirit (absolute alcohol) can be mixed in any
       proportion with water.

  [22] Just as the existence must he admitted of substances which are
       completely undecomposable (chemically) at the ordinary
       temperature--and of substances which are entirely non-volatile
       at such a temperature (as wood and gold), although capable
       of decomposing (wood) or volatilising (gold) at a higher
       temperature--so also the existence must be admitted of substances
       which are totally insoluble in water without some degree of
       change in their state. Although mercury is partially volatile at
       the ordinary temperature, there is no reason to think that it
       and other metals are soluble in water, alcohol, or other similar
       liquids. However, mercury forms solutions, as it dissolves
       other metals. On the other hand, there are many substances
       found in nature which are so very slightly soluble in water,
       that in ordinary practice they may be considered as insoluble
       (for example, barium sulphate). For the comprehension of that
       general plan according to which a change of state of substances
       (combined or dissolved, solid, liquid, or gaseous) takes place,
       it is very important to make a distinction at this boundary
       line (on approaching zero of decomposition, volatility, or
       solubility) between an insignificant amount and zero, but the
       present methods of research and the data at our disposal at the
       present time only just touch such questions (by studying the
       electrical conductivity of dilute solutions and the development
       of micro-organisms in them). It must be remarked, besides, that
       water in a number of cases does not dissolve a substance as
       such, but acts on it chemically and forms a soluble substance.
       Thus glass and many rocks, especially if taken as powder, are
       chemically changed by water, but are not directly soluble in it.

Substances which are easily soluble in water bear a certain resemblance
to it. Thus sugar and salt in many of their superficial features remind
one of ice. Metals, which are not soluble in water, have no points in
common with it, whilst on the other hand they dissolve each other in
a molten state, forming alloys, just as oily substances dissolve each
other; for example, tallow is soluble in petroleum and in olive oil,
although they are all insoluble in water. From this it is evident that
the _analogy of substances forming a solution_ plays an important part,
and as aqueous and all other solutions are liquids, there is good reason
to believe that in the process of solution solid and gaseous substances
change in a physical sense, passing into a liquid state. These
considerations elucidate many points of solution--as, for instance, the
variation of the co-efficient of solubility with the temperature and the
evolution or absorption of heat in the formation of solutions.

The solubility--that is, the quantity of a substance necessary for
saturation--_varies with the temperature_, and, further, with an
increase in temperature the solubility of solid substances generally
increases, and that of gases decreases; this might be expected, as solid
substances by heating, and gases by cooling, approach to a liquid or
dissolved state.[23] A graphic method is often employed to express the
variation of solubility with temperature. On the axis of abscissæ or on
a horizontal line, temperatures are marked out and perpendiculars are
raised corresponding with each temperature, whose length is determined
by the solubility of the salt at that temperature--expressing, for
instance, one part by weight of a salt in 100 parts of water by one
unit of length, such as a millimetre. By joining the summits of the
perpendiculars, a curve is obtained which expresses the degree of
solubility at different temperatures. For solids, the curve is generally
an ascending one--_i.e._ recedes from the horizontal line with the
rise in temperature. These curves clearly show by their inclination
the degree of rapidity of increase in solubility with the temperature.
Having determined several points of a curve--that is, having made a
determination of the solubility for several temperatures--the solubility
at intermediary temperatures may be determined from the form of the
curve so obtained; in this way the empirical law of solubility may be
examined.[24] The results of research have shown that the solubility of
certain salts--as, for example, common table salt--varies comparatively
little with the temperature; whilst for other substances the solubility
increases by equal amounts for equal increments of temperature. Thus,
for example, for the saturation of 100 parts of water by potassium
chloride there is required at 0°, 29·2 parts, at 20°, 34·7, at 40°,
40·2, at 60°, 45·7; and so on, for every 10° the solubility increases
by 2·75 parts by weight of the salt. Therefore the solubility of the
potassium chloride in water may be expressed by a direct equation: _a_ =
29·2 + 0·275_t_, where _a_ represents the solubility at _t_°. For other
salts, more complicated equations are required. For example, for nitre:
_a_ = 13·3 + 0·574_t_ + 0·01717_t_^2 + 0·0000036_t_^3, which shows that
when _t_ = 0° _a_ = 13·3, when _t_ = 10° _a_ = 20·8, and when _t_ = 100°
_a_ = 246·0.

  [23] Beilby (1883) experimented on paraffin, and found that one litre
       of solid paraffin at 21° weighed 874 grams, and when liquid, at
       its melting-point 38°, 783 grams, at 49°, 775 grams, and at 60°,
       767 grams, from which the weight of a litre of liquefied paraffin
       would be 795·4 grams at 21° if it could remain liquid at that
       temperature. By dissolving solid paraffin in lubricating oil at
       21° Beilby found that 795·6 grams occupy one cubic decimetre,
       from which he concluded that the solution contained liquefied
       paraffin.

  [24] Gay-Lussac was the first to have recourse to such a graphic
       method of expressing solubility, and he considered, in accordance
       with the general opinion, that by joining up the summits of the
       ordinates in one harmonious curve it is possible to express the
       entire change of solubility with the temperature. Now, there are
       many reasons for doubting the accuracy of such an admission, for
       there are undoubtedly critical points in curves of solubility
       (for example, of sodium sulphate, as shown further on), and it
       may be that definite compounds of dissolved substances with
       water, in decomposing within known limits of temperature, give
       critical points more often than would be imagined; it may even
       be, indeed, that instead of a continuous curve, solubility should
       be expressed--if not always, then not unfrequently--by straight
       or broken lines. According to Ditte, the solubility of sodium
       nitrate, NaNO_{3}, is expressed by the following figures per 100
       parts of water:--

          0°    4°   10°   15°   21°   29°   36°    51°    68°
         66·7  71·0  76·3  80·6  85·7  92·9  99·4  113·6  125·1

       In my opinion (1881) these data should be expressed with
       exactitude by a straight line, 67·5 + 0·87_t_, which entirely
       agrees with the results of experiment. According to this
       the figure expressing the solubility of salt at 0° exactly
       coincides with the composition of a definite chemical
       compound--NaNO_{3},7H_{2}O. The experiments made by Ditte
       showed that all saturated solutions between 0° and -15·7° have
       such a composition, and that at the latter temperature the
       solution completely solidifies into one homogeneous whole.
       Between 0° and -15·7° the solution NaNO_{3},7H_{2}O does not
       deposit either salt or ice. Thus the solubility of sodium
       nitrate is expressed by a broken straight line. In recent times
       (1888) Étard discovered a similar phenomenon in many of the
       sulphates. Brandes, in 1830, shows a diminution in solubility
       below 100° for manganese sulphate. The percentage by weight
       (_i.e._ per 100 parts of the solution, and not of water) of
       saturation for ferrous sulphate, FeSO_{4}, from -2° to +65° =
       13·5 + 0·3784_t_--that is, the solubility of the salt increases.
       The solubility remains constant from 65° to 98° (according to
       Brandes the solubility then increases; this divergence of opinion
       requires proof), and from 98° to 150° it falls as = 104·35 -
       0·6685_t_. Hence, at about +156° the solubility should = 0, and
       this has been confirmed by experiment. I observe, on my part,
       that Étard's formula gives 38·1 p.c. of salt at 65° and 38·8
       p.c. at 92°, and this maximum amount of salt in the solution
       very nearly corresponds with the composition FeSO_{4},14H_{2}O,
       which requires 37·6 p.c. From what has been said, it is evident
       that the data concerning solubility require a new method of
       investigation, which should have in view the entire scale of
       solubility--from the formation of completely solidified solutions
       (cryohydrates, which we shall speak of presently) to the
       separation of salts from their solutions, if this is accomplished
       at a higher temperature (for manganese and cadmium sulphates
       there is an entire separation, according to Étard), or to the
       formation of a constant solubility (for potassium sulphate the
       solubility, according to Étard, remains constant from 163° to
       220° and equals 24·9 p.c.) (See Chapter XIV., note 50, solubility
       of CaCl_{2}.)

Curves of solubility give the means of estimating the _amount of salt
separated_ by the cooling to a known extent of a solution saturated
at a given temperature. For instance, if 200 parts of a solution of
potassium chloride in water saturated at a temperature of 60° be taken,
and it be asked how much of the salt will be separated by cooling the
solution to 0°, if its solubility at 60° = 45·7 and at 0° = 29·2? The
answer is obtained in the following manner: At 60° a saturated solution
contains 45·7 parts of potassium chloride per 100 parts by weight of
water, consequently 145·7 parts by weight of the solution contain 45·7
parts, or, by proportion, 200 parts by weight of the solution contain
62·7 parts of the salt. The amount of salt remaining in solution at 0°
is calculated as follows; In 200 grams taken there will be 137·3 grams
of water; consequently, this amount of water is capable of holding only
40·1 grams of the salt, and therefore in lowering the temperature from
60° to 0° there should separate from the solution 62·7-40·1 = 22·6 grams
of the dissolved salt.

The difference in the solubility of salts, &c., with a rise or fall
of temperature is often taken advantage of, especially in technical
work, for the separation of salts, in intermixture from each other.
Thus a mixture of potassium and sodium chlorides (this mixture is met
with in nature at Stassfurt) is separated from a saturated solution by
subjecting it alternately to boiling (evaporation) and cooling. The
sodium chloride separates out in proportion to the amount of water
expelled from the solution by boiling, and is removed, whilst the
potassium chloride separates out on cooling, as the solubility of this
salt rapidly decreases with a lowering in temperature. Nitre, sugar, and
many other soluble substances are purified (refined) in a similar manner.

Although in the majority of cases the solubility of solids increases
with the temperature, yet there are some solid substances whose
solubilities decrease on heating. Glauber's salt, or sodium sulphate,
forms a particularly instructive example of the case in question.
If this salt be taken in an ignited state (deprived of its water of
crystallisation), then its solubility in 100 parts of water varies with
the temperature in the following manner: at 0°, 5 parts of the salt
form a saturated solution; at 20°, 20 parts of the salt, at 33° more
than 50 parts. The solubility, as will be seen, increases with the
temperature, as is the case with nearly all salts; but starting from 33°
it suddenly diminishes, and at a temperature of 40°, less than 50 parts
of the salt dissolve, at 60° only 45 parts of the salt, and at 100°
about 43 parts of the salt in 100 parts of water. This phenomenon may
be traced to the following facts: Firstly, that this salt forms various
compounds with water, as will be afterwards explained; secondly, that
at 33° the compound Na_{2}SO_{4} + 10H_{2}O formed from the solution
at lower temperatures, melts; and thirdly, that on evaporation at a
temperature above 33° an anhydrous salt, Na_{2}SO_{4} separates out. It
will be seen from this example how complicated such an apparently simple
phenomenon as solution really is; and all data concerning solutions lead
to the same conclusion. This complexity becomes evident in investigating
the _heat of solution_. If solution consisted of a physical change
only, then in the solution of gases there would be evolved--and in the
solution of solids, there would be absorbed--just that amount of heat
corresponding to the change of state; but in reality a large amount of
heat is always evolved in solution, depending on the fact that in the
process of solution chemical combination takes place accompanied by an
evolution of heat. Seventeen grams of ammonia (this weight corresponds
with its formula NH_{3}), in passing from a gaseous into a liquid state,
evolve 4,400 units of heat (latent heat); that is, the quantity of heat
necessary to raise the temperature of 4,400 grams of water 1°. The
same quantity of ammonia, in dissolving in an excess of water, evolves
twice as much heat--namely 8,800 units--showing that the combination
with water is accompanied by the evolution of 4,400 units of heat.
Further, the chief part of this heat is separated in dissolving in small
quantities of water, so that 17 grams of ammonia, in dissolving in 18
grams of water (this weight corresponds with its composition H_{2}O),
evolve 7,535 units of heat, and therefore the formation of the solution
NH_{3} + H_{2}O evolves 3,135 units of heat beyond that due to the
change of state. As in the solution of gases, the heat of liquefaction
(of physical change of state) and of chemical combination with water
are both positive (+), therefore in the _solution of gases_ in water a
_heat effect_ is always observed. This phenomenon is different in the
solution of solid substances, because their passage from a solid to a
liquid state is accompanied by an absorption of heat (negative,-heat),
whilst their chemical combination with water is accompanied by an
evolution of heat (+ heat); consequently, their sum may either be a
cooling effect, when the positive (chemical) portion of heat is less
than the negative (physical), or it may be, on the contrary, a heating
effect. This is actually the case. 124 grams of sodium thiosulphate
(employed in photography) Na_{2}S_{2}O_{3},5H_{2}O in melting (at 48°)
absorbs 9,700 units of heat, but in dissolving in a large quantity of
water at the ordinary temperature it absorbs 5,700 units of heat, which
shows the evolution of heat (about + 4,000 units), notwithstanding the
cooling effect observed in the process of solution, in the act of the
chemical combination of the salt with water.[25] But in most cases
solid substances in dissolving in water evolve heat, notwithstanding
the passage into a liquid state, which indicates so considerable an
evolution of (+) heat in the act of combination with water that it
exceeds the absorption of (-) heat dependent on the passage into a
liquid state, Thus, for instance, calcium chloride, CaCl_{2}, magnesium
sulphate, MgSO_{4}, and many other salts evolve heat in dissolving; for
example, 60 grams of magnesium sulphate evolve about 10,000 units of
heat. Therefore, _in the solution of solid bodies_ either a cooling[26]
or a heating[27] effect is produced, according to the difference of the
reacting affinities. When they are considerable--that is, when water is
with difficulty separated from the resultant solution, and only with a
rise of temperature (such substances absorb water vapour)--then much
heat is evolved in the process of solution, just as in many reactions of
direct combination, and therefore a considerable heating of the solution
is observed. Of such a kind, for instance, is the solution of sulphuric
acid (oil of vitriol H_{2}SO_{4}), and of caustic soda (NaHO), &c., in
water.[28]

  [25] The latent heat of fusion is determined at the temperature of
       fusion, whilst solution takes place at the ordinary temperature,
       and one must think that at this temperature the latent heat
       would be different, just as the latent heat of evaporation
       varies with the temperature (see Note 11). Besides which, in
       dissolving, disintegration of the particles of both the solvent
       and the substance dissolved takes place, a process which in its
       mechanical aspect resembles evaporation, and therefore must
       consume much heat. The heat emitted in the solution of a solid
       must therefore be considered (Personne) as composed of three
       factors--(1) positive, the effect of combination; (2) negative,
       the effect of transference into a liquid state; and (3) negative,
       the effect of disintegration. In the solution of a liquid by
       a liquid the second factor is removed; and therefore, if the
       heat evolved in combination is greater than that absorbed in
       disintegration a heating effect is observed, and in the reverse
       case a cooling effect; and, indeed, sulphuric acid, alcohol, and
       many liquids evolve heat in dissolving in each other. But the
       solution of chloroform in carbon bisulphide (Bussy and Binget),
       or of phenol (or aniline) in water (Alexéeff), produces cold.
       In the solution of a small quantity of water in acetic acid
       (Abasheff), or hydrocyanic acid (Bussy and Binget), or amyl
       alcohol (Alexéeff), cold is produced, whilst in the solution of
       these substances in an excess of water heat is evolved.

       The relation existing between the solubility of solid bodies and
       the heat and temperature of fusion and solution has been studied
       by many investigators, and more recently (1893) by Schröder,
       who states that in the solution of a solid body in a solvent
       which does not act chemically upon it, a very simple process
       takes place, which differs but little from the intermixture of
       two gases which do not react chemically upon each other. The
       following relation between the heat of solution _Q_ and the
       heat of fusion _p_ may then be taken: _P_/_T__{0} = _Q_/_T_ =
       constant, where _T__{0} and _T_ are the absolute (from -273°)
       temperatures of fusion and saturation. Thus, for instance, in the
       case of naphthalene the calculated and observed magnitudes of the
       heat of solution differ but slightly from each other.

       The fullest information concerning the solution of liquids
       in liquids has been gathered by W. T. Alexéeff (1883-1885);
       these data are, however, far from being sufficient to solve
       the mass of problems respecting this subject. He showed that
       two liquids which dissolve in each other, intermix together in
       all proportions at a certain temperature. Thus the solubility
       of phenol, C_{6}H_{6}O, in water, and the converse, is limited
       up to 70°, whilst above this temperature they intermix in all
       proportions. This is seen from the following figures, where p
       is the percentage amount of phenol and _t_ the temperature at
       which the solution becomes turbid--that is, that at which it is
       saturated:--

       _p_ = 7·12  10·20  15·31  26·15  28·55  36·70  48·86  61·15  71·97
       _t_ = 1°    45°    60°    67°    67°    67°    65°    53°    20°

       It is exactly the same with the solution of benzene, aniline,
       and other substances in molten sulphur. Alexéeff discovered a
       similar complete intermixture for solutions of secondary butyl
       alcohol in water at about 107°; at lower temperatures the
       solubility is not only limited, but between 50° and 70° it is at
       its minimum, both for solutions of the alcohol in water and for
       water in the alcohol; and at a temperature of 5° both solutions
       exhibit a fresh change in their scale of solubility, so that a
       solution of the alcohol in water which is saturated between 5°
       and 40° will become turbid when heated to 60°. In the solution of
       liquids in liquids, Alexéeff observed a lowering in temperature
       (an absorption of heat) and an absence of change in specific
       heat (calculated for the mixture) much more frequently than had
       been done by previous observers. As regards his hypothesis (in
       the sense of a mechanical and not a chemical representation of
       solutions) that substances in solution preserve their physical
       states (as gases, liquids, or solids), it is very doubtful, for
       it would necessitate admitting the presence of ice in water or
       its vapour.

       From what has been said above, it will be clear that even in so
       very simple a case as solution, it is impossible to calculate
       the heat emitted by chemical action alone, and that the chemical
       process cannot be separated from the physical and mechanical.

  [26] The cooling effect produced in the solution of solids (and also in
       the expansion of gases and in evaporation) is applied to the
       _production of low temperatures_. Ammonium nitrate is very often
       used for this purpose; in dissolving in water it absorbs 77 units
       of heat per each part by weight. On evaporating the solution thus
       formed, the solid salt is re-obtained. The application of the
       various _freezing mixtures_ is based on the same principle. Snow
       or broken ice frequently enters into the composition of these
       _mixtures_, advantage being taken of its latent heat of fusion in
       order to obtain the lowest possible temperature (without altering
       the pressure or employing heat, as in other methods of obtaining
       a low temperature). For laboratory work recourse is most often
       had to a mixture of three parts of snow and one part of common
       salt, which causes the temperature to fall from 0° to -21° C.
       Potassium thiocyanate, KCNS, mixed with water (3/4 by weight of
       the salt) gives a still lower temperature. By mixing ten parts of
       crystallised calcium chloride, CaCl_{2},6H_{2}O, with seven parts
       of snow, the temperature may even fall from 0° to -55°.

  [27] The heat which is evolved in solution, or even in the dilution of
       solutions, is also sometimes made use of in practice. Thus
       caustic soda (NaHO), in dissolving or on the addition of water to
       a strong solution of it, evolves so much heat that it can replace
       fuel. In a steam boiler, which has been previously heated to
       the boiling point, another boiler is placed containing caustic
       soda, and the exhaust steam is made to pass through the latter;
       the formation of steam then goes on for a somewhat long period
       of time without any other heating. Norton makes use of this for
       smokeless street locomotives.

  [28] [Illustration: FIG. 17.--Curves expressing the contraction,
       quantity of heat, and rises of temperature produced by mixing
       sulphuric acid with water. Percentage of H_{2}SO_{4} is given
       along the axis of abscissae.]

       The temperatures obtained by mixing monohydrated sulphuric acid,
       H_{2}SO_{4}, with different quantities of water, are shown on
       the lowest curve in fig. 17, the relative proportions of both
       substances being expressed in percentages by weight along the
       horizontal axis. The greatest rise of temperature is 149°. It
       corresponds with the greatest evolution of heat (given on the
       middle curve) corresponding with a definite volume (100 c.c.)
       of the solution produced. The top curve expresses the degree
       of contraction, which also corresponds with 100 volumes of
       the solution produced. The greatest contraction, as also the
       greatest rise of temperature, corresponds with the formation of a
       trihydrate, H_{2}SO_{4},2H_{2}O (= 73·1 p.c. H_{2}SO_{4}), which
       very likely repeats itself in a similar form in other solutions,
       although all the phenomena (of contraction, evolution of heat,
       and rise of temperature) are very complex and are dependent
       on many circumstances. One would think, however, judging from
       the above examples, that all other influences are feebler in
       their action than chemical attraction, especially when it is so
       considerable as between sulphuric acid and water.

Solution is a reversible reaction; for, if the water be expelled from a
solution, the substance originally taken is obtained again. But it must
be borne in mind that the expulsion of the water taken for solution is
not always accomplished with equal facility, because water has different
degrees of chemical affinity for the substance dissolved. Thus, if a
solution of sulphuric acid, which mixes with water in all proportions,
be heated, it will be found that very different degrees of heat are
required to expel the water. When it is in a large excess, water is
given off at a temperature slightly above 100°, but if it be in but a
small proportion there is such an affinity between it and the sulphuric
acid that at 120°, 150°, 200°, and even at 300°, water is still retained
by the sulphuric acid. The bond between the remaining quantity of water
and the sulphuric acid is evidently stronger than the bond between the
sulphuric acid and the excess of water. The force acting in solutions
is consequently of different intensity, starting from so feeble an
attraction that the properties of water--as, for instance, its power
of evaporation--are but very little changed, and ending with cases of
strong attraction between the water and the substance dissolved in or
chemically combined with it. In consideration of the very important
significance of the phenomena, and of the cases of the breaking up of
solutions with separation of water or of the substance dissolved from
them, we shall further discuss them separately, after having acquainted
ourselves with certain peculiarities of the solution of gases and of
solid bodies.

The solubility of gases, which is usually measured by the volume of
gas[29] (at 0° and 760 mm. pressure) per 100 volumes of water, varies
not only with the nature of the gas (and also of the solvent), and with
the temperature, but also with the pressure, because gases themselves
change their volume considerably with the pressure. As might be
expected, (1) gases which are easily liquefied (by pressure and cold)
are more soluble than those which are liquefied with difficulty. Thus,
in 100 volumes of water only two volumes of hydrogen dissolve at 0°
and 760 mm., three volumes of carbonic oxide, four volumes of oxygen,
&c., for these are gases which are liquefied with difficulty; whilst
there dissolve 180 volumes of carbonic anhydride, 130 of nitrous oxide,
and 437 of sulphurous anhydride, for these are gases which are rather
easily liquefied. (2) The solubility of a gas is diminished by heating,
which is easily intelligible from what has been said previously--the
elasticity of a gas becomes greater, it is removed further from a liquid
state. Thus 100 volumes of water at 0° dissolve 2·5 volumes of air, and
at 20° only 1·7 volume. For this reason cold water, when brought into a
warm room, parts with a portion of the gas dissolved in it.[30] (3) The
quantity of the gas dissolved varies directly with the pressure. This
rule is called the _law of Henry and Dalton_, and is applicable to those
gases which are little soluble in water. Therefore a gas is separated
from its solution in water in a vacuum, and water saturated with a gas
under great pressure parts with it if the pressure be diminished. Thus
many mineral springs are saturated underground with carbonic anhydride
under the great pressure of the column of water above them. On coming
to the surface, the water of these springs boils and foams on giving
up the excess of dissolved gas. Sparkling wines and aërated waters are
saturated under pressure with the same gas. They hold the gas so long
as they are in a well-corked vessel. When the cork is removed and the
liquid comes in contact with air at a lower pressure, part of the gas,
unable to remain in solution at a lower pressure, is separated as froth
with the hissing sound familiar to all. It must be remarked that the law
of Henry and Dalton belongs to the class of _approximate laws_, like
the laws of gases (Gay-Lussac's and Mariotte's) and many others--that
is, it expresses only a portion of a complex phenomenon, the limit
towards which the phenomenon aims. The matter is rendered complicated
from the influence of the degree of solubility and of affinity of the
dissolved gas for water. Gases which are little soluble--for instance,
hydrogen, oxygen, and nitrogen--follow the law of Henry and Dalton the
most closely. Carbonic anhydride exhibits a decided deviation from the
law, as is seen from the determinations of Wroblewski (1882). He showed
that at 0° a cubic centimetre of water absorbs 1·8 cubic centimetre of
the gas under a pressure of one atmosphere; under 10 atmospheres, 16
cubic centimetres (and not 18, as it should be according to the law);
under 20 atmospheres, 26·6 cubic centimetres (instead of 36), and under
30 atmospheres, 33·7 cubic centimetres.[31] However, as the researches
of Sechenoff show, the absorption of carbonic anhydride within certain
limits of change of pressure, and at the ordinary temperature, by
water--and even by solutions of salts which are not chemically changed
by it, or do not form compounds with it--very closely follows the law of
Henry and Dalton, so that the chemical bond between this gas and water
is so feeble that the breaking up of the solution with separation of
the gas is accomplished by a decrease of pressure alone.[32] The case
is different if a considerable affinity exists between the dissolved
gas and water. Then it might even be expected that the gas would not be
entirely separated from water in a vacuum, as should be the case with
gases according to the law of Henry and Dalton. Such gases--and, in
general, all those which are very soluble--exhibit a distinct deviation
from the law of Henry and Dalton. As examples, ammonia and hydrochloric
acid gas may be taken. The former is separated by boiling and decrease
of pressure, while the latter is not, but they both deviate distinctly
from the law.

        +---------------+-----------------+--------------------+
        |Pressure in mm.|Ammonia dissolved| Hydrochloric acid  |
        |  of mercury   | in 100 grams of |gas dissolved in 100|
        |               |   water at 0°   |grams of water at 0°|
        +---------------+-----------------+--------------------+
        |               |      Grams      |       Grams        |
        |       100     |       28·0      |        65·7        |
        |       500     |       69·2      |        78·2        |
        |     1,000     |      112·6      |        85·6        |
        |     1,500     |      165·6      |         --         |
        +---------------+-----------------+--------------------+

  [29] If a volume of gas _v_ be measured under a pressure of _h_ mm. of
       mercury (at 0°) and at a temperature _t_° Centigrade, then,
       according to the combined laws of Boyle, Mariotte, and of
       Gay-Lussac, its volume at 0° and 760 mm. will equal the product
       of _v_ into 760 divided by the product of _h_ into 1 + _a__t_°,
       where _a_ is the co-efficient of expansion of gases, which is
       equal to 0·00367. The weight of the gas will be equal to its
       volume at 0° and 760 mm. multiplied by its density referred to
       air and by the weight of one volume of air at 0° and 760 mm.
       The weight of one litre of air under these conditions being =
       1·293 gram. If the density of the gas be given in relation to
       hydrogen this must be divided by 14·4 to bring it in relation to
       air. If the gas be measured when saturated with aqueous vapour,
       then it must be reduced to the volume and weight of the gas when
       dry, according to the rules given in Note 1. If the pressure
       be determined by a column of mercury having a temperature _t_,
       then by dividing the height of the column by 1 + 0·00018_t_ the
       corresponding height at 0° is obtained. If the gas be enclosed in
       a tube in which a liquid stands above the level of the mercury,
       the height of the column of the liquid being = H and its density
       = D, then the gas will be under a pressure which is equal to the
       barometric pressure less HD/13·59, where 13·59 is the density of
       mercury. By these methods the _quantity of a gas_ is determined,
       and its observed volume reduced to normal conditions or to parts
       by weight. The physical data concerning vapours and gases must be
       continually kept in sight in dealing with and measuring gases.
       The student must become perfectly familiar with the calculations
       relating to gases.

  [30] According to Bunsen, Winkler, Timofeeff, and others, 100 vols. of
       water under a pressure of one atmosphere absorb the following
       volumes of gas (measured at 0° and 760 mm.):--

            1    2    3     4    5     6     7     8    9    10    11
        0° 4·82 2·35 2·15 179·7 3·54 130·5 437·1 688·6 5·4 104960 7·38
       20° 3·10 1·54 1·83  90·1 2·32  67·0 290·5 362·2 3·5  65400 4·71

       1, oxygen; 2, nitrogen; 3, hydrogen; 4, carbonic anhydride;
       5, carbonic oxide; 6, nitrous oxide; 7, hydrogen sulphide; 8,
       sulphurous anhydride; 9, marsh gas; 10, ammonia; 11, nitric
       oxide. The decrease of solubility with a rise of temperature
       varies for different gases; it is greater, the greater the
       molecular weight of the gas. It is shown by calculation that
       this decrease varies (Winkler) as the cube root of the molecular
       weight of the gas. This is seen from the following table:

             +--------------+-------------+---------------+
             | Decrease of  | Cube root of| Ratio between |
             | solubility   |  molecular  | decrease and  |
             | per 20° in   |   weight.   | cube root of  |
             |  per cent.   |             |   mol. wt.    |
             +--------------+-------------+---------------+
             | H_{2}  15·32 |    1·259    |     12·17     |
             | N_{2}  34·33 |    3·037    |     11·30     |
             | CO     34·44 |    3·037    |     11·34     |
             | NO     36·24 |    3·107    |     11·66     |
             | O_{2}  36·55 |    3·175    |     11·51     |
             +--------------+-------------+---------------+

       The decrease in the coefficient of absorption with the
       temperature must be connected with a change in the physical
       properties of the water. Winkler (1891) remarked a certain
       relation between the internal friction and the coefficient of
       absorption at various temperatures.

  [31] These figures show that the co-efficient of solubility decreases
       with an increase of pressure, notwithstanding that the carbonic
       anhydride approaches a liquid state. As a matter of fact,
       liquefied carbonic anhydride does not intermix with water,
       and does not exhibit a rapid increase in solubility at its
       temperature of liquefaction. This indicates, in the first place,
       that solution does not consist in liquefaction, and in the
       second place that the solubility of a substance is determined
       by a peculiar attraction of water for the substance dissolving.
       Wroblewski even considered it possible to admit that a dissolved
       gas retains its properties as a gas. This he deduced from
       experiments, which showed that the rate of diffusion of gases
       in a solvent is, for gases of different densities, inversely
       proportional to the square roots of their densities, just as
       the velocities of gaseous molecules (see Note 34). Wroblewski
       showed the affinity of water, H_{2}O, for carbonic anhydride,
       CO_{2}, from the fact that on expanding moist compressed carbonic
       anhydride (compressed at 0° under a pressure of 10 atmospheres)
       he obtained (a fall in temperature takes place from the
       expansion) a very unstable definite crystalline compound, CO_{2}
       + 8H_{2}O.

  [32] As, according to the researches of Roscoe and his collaborators,
       ammonia exhibits a considerable deviation at low temperatures
       from the law of Henry and Dalton, whilst at 100° the deviation
       is small, it would appear that the dissociating influence of
       temperature affects all gaseous solutions; that is, at high
       temperatures, the solutions of all gases will follow the law, and
       at lower temperatures there will in all cases be a deviation from
       it.

It will be remarked, for instance, from this table that whilst the
pressure increased 10 times, the solubility of ammonia only increased
4-1/2 times.

A number of examples of such cases of the absorption of gases by liquids
might be cited which do not in any way, even approximately, agree with
the laws of solubility. Thus, for instance, carbonic anhydride is
absorbed by a solution of caustic potash in water, and if sufficient
caustic potash be present it is not separated from the solution by
a decrease of pressure. This is a case of more intimate chemical
combination. A correlation less completely studied, but similar and
clearly chemical, appears in certain cases of the solution of gases in
water, and we shall afterwards find an example of this in the solution
of hydrogen iodide; but we will first stop to consider a remarkable
application of the law of Henry and Dalton[33] in the case of the
solution of a mixture of two gases, and this we must do all the more
because the phenomena which there take place cannot be foreseen without
a clear theoretical representation of the nature of gases.[34]

  [33] The ratio between the pressure and the amount of gas dissolved
       was discovered by Henry in 1805, and Dalton in 1807 pointed
       out the adaptability of this law to cases of gaseous mixtures,
       introducing the conception of partial pressures which is
       absolutely necessary for a right comprehension of Dalton's law.
       The conception of partial pressures essentially enters into
       that of the diffusion of vapours in gases (footnote 1); for
       the pressure of damp air is equal to the sum of the pressures
       of dry air and of the aqueous vapour in it, and it is admitted
       as a corollary to Dalton's law that evaporation in dry air
       takes place as in a vacuum. It is, however, necessary to remark
       that the volume of a mixture of two gases (or vapours) is only
       approximately equal to the sum of the volumes of its constituents
       (the same, naturally, also refers to their pressures)--that is to
       say, in mixing gases a change of volume occurs, which, although
       small, is quite apparent when carefully measured. For instance,
       in 1888 Brown showed that on mixing various volumes of sulphurous
       anhydride (SO_{2}) with carbonic anhydride (at equal pressures
       of 760 mm. and equal temperatures) a decrease of pressure of
       3·9 millimetres of mercury was observed. The possibility of a
       chemical action in similar mixtures is evident from the fact that
       equal volumes of sulphurous and carbonic anhydrides at -19° form,
       according to Pictet's researches in 1888, a liquid which may be
       regarded as an unstable chemical compound, or a solution similar
       to that given when sulphurous anhydride and water combine to an
       unstable chemical whole.

  [34] The origin of the kinetic theory of gases now generally accepted,
       according to which they are animated by a rapid progressive
       motion, is very ancient (Bernouilli and others in the last
       century had already developed a similar representation), but
       it was only generally accepted after the mechanical theory of
       heat had been established, and after the work of Krönig (1855),
       and especially after its mathematical side had been worked out
       by Clausius and Maxwell. The pressure, elasticity, diffusion,
       and internal friction of gases, the laws of Boyle, Mariotte,
       and of Gay-Lussac and Avogadro-Gerhardt are not only explained
       (deduced) by the kinetic theory of gases, but also expressed
       with perfect exactitude; thus, for example, the magnitude of
       the internal friction of different gases was foretold with
       exactitude by Maxwell, by applying the theory of probabilities
       to the impact of gaseous particles. The kinetic theory of gases
       must therefore be considered as one of the most brilliant
       acquisitions of the latter half of the present century. The
       velocity of the progressive motion of the particles of a gas, one
       cubic centimetre of which weighs _d_ grams, is found, according
       to the theory, to be equal to the square root of the product of
       3_pDq_ divided by _d_, where _p_ is the pressure under which _d_
       is determined expressed in centimetres of the mercury column,
       _D_ the weight of a cubic centimetre of mercury in grams (_D_
       = 13·59, _p_ = 76, consequently the normal pressure = 1,033
       grams on a sq. cm.), and _g_ the acceleration of gravity in
       centimetres (_g_ = 980·5, at the sea level and long. 45° = 981·92
       at St. Petersburg; in general it varies with the longitude and
       altitude of the locality). Therefore, at 0° the velocity of
       hydrogen is 1,843, and of oxygen 461, metres per second. This
       is the average velocity, and (according to Maxwell and others)
       it is probable that the velocities of individual particles
       are different; that is, they occur in, as it were, different
       conditions of temperature, which it is very important to take
       into consideration in investigating many phenomena proper to
       matter. It is evident from the above determination of the
       velocity of gases, that different gases at the same temperature
       and pressure have average velocities, which are inversely
       proportional to the square roots of their densities; this is also
       shown by direct experiment on the flow of gases through a fine
       orifice, or through a porous wall. This _dissimilar velocity of
       flow_ for different gases is frequently taken advantage of in
       chemical researches (see Chap. II. and also Chap. VII.) in order
       to separate two gases having different densities and velocities.
       The difference of the velocity of flow of gases also determines
       the phenomenon cited in the following footnote for demonstrating
       the existence of an internal motion in gases.

       If for a certain mass of a gas which fully and exactly follows
       the laws of Mariotte and Gay-Lussac the temperature _t_ and the
       pressure _p_ be changed simultaneously, then the entire change
       would be expressed by the equation _pv_ = _C_(1 + _at_), or,
       what is the same, _pv_ = _RT_, where _T_ = _t_ + 273 and _C_
       and _R_ are constants which vary not only with the units taken
       but with the nature of the gas and its mass. But as there are
       discrepancies from both the fundamental laws of gases (which
       will be discussed in the following chapter), and as, on the
       one hand, a certain attraction between the gaseous molecules
       must be admitted, while on the other hand the molecules of
       gases themselves must occupy a portion of a space, hence for
       ordinary gases, within any considerable variation of pressure and
       temperature, recourse should be had to Van der Waal's formula--

                (_p_ + _a_/_v_^2)(_v_-_p_) = R(1 + _at_)

       where _a_ is the true co-efficient of expansion of gases.

       The formula of Van der Waals has an especially important
       significance in the case of the passage of a gas into a liquid
       state, because the fundamental properties of both gases and
       liquids are equally well expressed by it, although only in their
       general features.

       The further development of the questions referring to the
       subjects here touched on, which are of especial interest for the
       theory of solutions, must be looked for in special memoirs and
       works on theoretical and physical chemistry. A small part of this
       subject will be partially considered in the footnotes of the
       following chapter.

_The law of partial pressures_ is as follows:--The solubility of gases
in intermixture with each other does not depend on the influence of
the total pressure acting on the mixture, but on the influence of that
portion of the total pressure which is due to the volume of each given
gas in the mixture. Thus, for instance, if oxygen and carbonic anhydride
were mixed in equal volumes and exerted a pressure of 760 millimetres,
then water would dissolve so much of each of these gases as would be
dissolved if each separately exerted a pressure of half an atmosphere,
and in this case, at 0° one cubic centimetre of water would dissolve
0·02 cubic centimetre of oxygen and 0·90 cubic centimetre of carbonic
anhydride. If the pressure of a gaseous mixture equals _h_, and in _n_
volumes of the mixture there be _a_ volumes of a given gas, then its
solution will proceed as though this gas were dissolved under a pressure
(_h_ × _a_)/_n_. That portion of the pressure under influence of which
the solution proceeds is termed the 'partial' pressure.

In order to clearly understand the cause of the law of partial
pressures, an explanation must be given of the fundamental properties
of gases. Gases are elastic and disperse in all directions. We are led
from what we know of gases to the assumption that these fundamental
properties of gases are due to a rapid progressive motion, in all
directions, which is proper to their smallest particles (molecules).[35]
These molecules in impinging against an obstacle produce a pressure.
The greater the number of molecules impinging against an obstacle in
a given time, the greater the pressure. The pressure of a separate
gas or of a gaseous mixture depends on the sum of the pressures of
all the molecules, on the number of blows in a unit of time on a unit
of surface, and on the mass and velocity (or the _vis viva_) of the
impinging molecules. The nature of the different molecules is of no
account; the obstacle is acted on by a pressure due to the sum of their
_vis viva_. But, in a chemical action such as the solution of gases,
the nature of the impinging molecules plays, on the contrary, the
most important part. In impinging against a liquid, a portion of the
gas enters into the liquid itself, and is held by it so long as other
gaseous molecules impinge against the liquid--exert a pressure on it. As
regards the solubility of a given gas, for the number of blows it makes
on the surface of a liquid, it is immaterial whether other molecules
of gases impinge side by side with it or not. Hence, the solubility
of a given gas will be proportional, not to the total pressure of a
gaseous mixture, but to that portion of it which is due to the given
gas separately. Moreover, the saturation of a liquid by a gas depends
on the fact that the molecules of gases that have entered into a liquid
do not remain at rest in it, although they enter in a harmonious kind
of motion with the molecules of the liquid, and therefore they throw
themselves off from the surface of the liquid (just like its vapour
if the liquid be volatile). If in a unit of time an equal number of
molecules penetrate into (leap into) a liquid and leave (or leap out
of) a liquid, it is saturated. It is a case of mobile equilibrium, and
not of rest. Therefore, if the pressure be diminished, the number of
molecules departing from the liquid will exceed the number of molecules
entering into the liquid, and a fresh state of mobile equilibrium only
takes place under a fresh equality of the number of molecules departing
from and entering into the liquid. In this manner the main features of
the solution are explained, and furthermore of that special (chemical)
attraction (penetration and harmonious motion) of a gas for a liquid,
which determines both the measure of solubility and the degree of
stability of the solution produced.

  [35] Although the actual motion of gaseous molecules, which is accepted
       by the kinetic theory of gases, cannot be seen, yet its existence
       may be rendered evident by taking advantage of the difference in
       the velocities undoubtedly belonging to different gases which
       are of different densities under equal pressures. The molecules
       of a light gas must move more rapidly than the molecules of
       a heavier gas in order to produce the same pressure. Let us
       take, therefore, two gases--hydrogen and air; the former is
       14·4 times lighter than the latter, and hence the molecules of
       hydrogen must move almost four times more quickly than air (more
       exactly 3·8, according to the formula given in the preceding
       footnote). Consequently, if a porous cylinder containing air is
       introduced into an atmosphere of hydrogen, then in a given time
       the volume of hydrogen which succeeds in entering the cylinder
       will be greater than the volume of air leaving the cylinder, and
       therefore the pressure inside the cylinder will rise until the
       gaseous mixture (of air and hydrogen) attains an equal density
       both inside and outside the cylinder. If now the experiment
       be reversed and air surround the cylinder, and hydrogen be
       inside the cylinder, then more gas will leave the cylinder than
       enters it, and hence the pressure inside the cylinder will be
       diminished. In these considerations we have replaced the idea of
       the number of molecules by the idea of volumes. We shall learn
       subsequently that equal volumes of different gases contain an
       equal number of molecules (the law of Avogadro-Gerhardt), and
       therefore instead of speaking of the number of molecules we can
       speak of the number of volumes. If the cylinder be partially
       immersed in water the rise and fall of the pressure can be
       observed directly, and the experiment consequently rendered
       self-evident.

The consequences of the law of partial pressures are exceedingly
numerous and important. All liquids in nature are in contact with the
atmosphere, which, as we shall afterwards see more fully, consists of
an intermixture of gases, chiefly four in number--oxygen, nitrogen,
carbonic anhydride, and aqueous vapour. 100 volumes of air contain,
approximately, 78 volumes of nitrogen, and about 21 volumes of oxygen;
the quantity of carbonic anhydride, by volume, does not exceed 0·05.
Under ordinary circumstances, the quantity of aqueous vapour is much
greater than this, but it varies of course with climatic conditions. We
conclude from these numbers that the solution of nitrogen in a liquid
in contact with the atmosphere will proceed under a partial pressure of
(78/100) × 760 mm. if the atmospheric pressure equal 760 mm.; similarly,
under a pressure of 600 mm. of mercury, the solution of oxygen will
proceed under a partial pressure of about 160 mm., and the solution of
carbonic anhydride only under the very small pressure of 0·4 mm. As,
however, the solubility of oxygen in water is twice that of nitrogen,
the ratio of O to N dissolved in water will be greater than the ratio in
air. It is easy to calculate what quantity of each of the gases will be
contained in water, and taking the simplest case we will calculate what
quantity of oxygen, nitrogen, and carbonic anhydride will be dissolved
from air having the above composition at 0° and 760 mm. pressure. Under
a pressure of 760 mm. 1 cubic centimetre of water dissolves 0·0203
cubic centimetre of nitrogen or under the partial pressure of 600 mm. it
will dissolve 0·0203 × 600/760, or 0·0160 cubic centimetre; of oxygen
0·0411 × 160/760, or 0·0086 cubic centimetre; of carbonic anhydride 1·8
× 0·4/760 or 0·00095 cubic centimetre: hence, 100 cubic centimetres
of water will contain at 0° altogether 2·55 cubic centimetres of
atmospheric gases, and 100 volumes of air dissolved in water will
contain about 62 p.c. of nitrogen, 34 p.c. of oxygen, and 4 p.c. of
carbonic anhydride. The water of rivers, wells, &c. usually contains
more carbonic anhydride. This proceeds from the oxidation of organic
substances falling into the water. The amount of oxygen, however,
dissolved in water appears to be actually about 1/3 the dissolved gases,
whilst air contains only 1/5 of it by volume.

According to the law of partial pressures, whatever gas be dissolved in
water will be expelled from the solution in an atmosphere of another
gas. This depends on the fact that gases dissolved in water escape
from it in a vacuum, because the pressure is nil. An atmosphere of
another gas acts like a vacuum on a gas dissolved in water. Separation
then proceeds, because the molecules of the dissolved gas no longer
impinge upon the liquid, are not dissolved in it, and those previously
held in solution leave the liquid in virtue of their elasticity.[36]
For the same reason a gas may be entirely expelled from a gaseous
solution by boiling--at least, in many cases when it does not form
particularly stable compounds with water. In fact the surface of the
boiling liquid will be occupied by aqueous vapour, and therefore all
the pressure acting on the gas will be due to the aqueous vapour. On
this account, the partial pressure of the dissolved gas will be very
inconsiderable, and this is the sole reason why _a gas separates from a
solution on boiling the liquid containing it_. At the boiling point of
water the solubility of gases in water is still sufficiently great for a
considerable quantity of a gas to remain in solution. The gas dissolved
in the liquid is carried away, together with the aqueous vapour; if
boiling be continued for a long time, all the gas will finally be
separated.[37]

  [36] Here two cases occur; either the atmosphere surrounding the
       solution may be limited, or it may be proportionally so vast
       as to be unlimited, like the earth's atmosphere. If a gaseous
       solution be brought into an atmosphere of another gas which is
       limited--for instance, as in a closed vessel--then a portion of
       the gas held in solution will be expelled, and thus pass over
       into the atmosphere surrounding the solution, and will produce
       its partial pressure. Let us imagine that water saturated with
       carbonic anhydride at 0° and under the ordinary pressure is
       brought into an atmosphere of a gas which is not absorbed by
       water; for instance, that 10 c.c. of an aqueous solution of
       carbonic anhydride is introduced into a vessel holding 10 c.c.
       of such a gas. The solution will contain 18 c.c. of carbonic
       anhydride. The expulsion of this gas proceeds until a state
       of equilibrium is arrived at. The liquid will then contain a
       certain amount of carbonic anhydride, which is retained under the
       partial pressure of that gas which has been expelled. Now, how
       much gas will remain in the liquid and how much will pass over
       into the surrounding atmosphere? In order to solve this problem,
       let us suppose that _x_ cubic centimetres of carbonic anhydride
       are retained in the solution. It is evident that the amount
       of carbonic anhydride which passed over into the surrounding
       atmosphere will be 18-_x_, and the total volume of gas will be
       10 + 18-_x_ or 28-_x_ cubic centimetres. The partial pressure
       under which the carbonic anhydride is then dissolved will be
       (supposing that the common pressure remains constant the whole
       time) equal to (18-_x_)/(28-_x_), hence there is not in solution
       18 c.c. of carbonic anhydride (as would be the case were the
       partial pressure equal to the atmospheric pressure), but only
       18(18-_x_)/(28-_x_), which is equal to _x_, and we therefore
       obtain the equation 18(18-_x_)/(28-_x_) = _x_, hence _x_ = 8·69.
       Again, where the atmosphere into which the gaseous solution is
       introduced is not only that of another gas but also unlimited,
       then the gas dissolved will, on passing over from the solution,
       diffuse into this atmosphere, and produce an infinitely small
       pressure in the unlimited atmosphere. Consequently, no gas can
       be retained in solution under this infinitely small pressure,
       and it will be entirely expelled from the solution. For this
       reason water saturated with a gas which is not contained in air,
       will be entirely deprived of the dissolved gas if left exposed
       to the air. Water also passes off from a solution into the
       atmosphere, and it is evident that there might be such a case as
       a constant proportion between the quantity of water vaporised and
       the quantity of a gas expelled from a solution, so that not the
       gas alone, but the entire gaseous solution, would pass off. A
       similar case is exhibited in solutions which are not decomposed
       by heat (such as those of hydrogen chloride and iodide), as will
       afterwards be considered.

  [37] However, in those cases when the variation of the co-efficient of
       solubility with the temperature is not sufficiently great, and
       when a known quantity of aqueous vapour and of the gas passes
       off from a solution at the boiling point, an atmosphere may be
       obtained having the same composition as the liquid itself. In
       this case the amount of gas passing over into such an atmosphere
       will not be greater than that held by the liquid, and therefore
       such a gaseous solution will distil over unchanged. The solution
       will then represent, like a solution of hydriodic acid in water,
       a liquid which is not altered by distillation, while the pressure
       under which this distillation takes place remains constant. Thus
       in all its aspects solution presents gradations from the most
       feeble affinities to examples of intimate chemical combination.
       The _amount of heat_ evolved in the solution of equal volumes of
       different gases is in distinct relation with these variations of
       stability and solubility of different gases. 22·3 litres of the
       following gases (at 760 mm. pressure) evolve the following number
       of (gram) units of heat in dissolving in a large mass of water;
       carbonic anhydride 5,600, sulphurous anhydride 7,700, ammonia
       8,800, hydrochloric acid 17,400, and hydriodic acid 19,400. The
       two last-named gases, which are not expelled from their solution
       by boiling, evolve approximately twice as much heat as gases like
       ammonia, which are separated from their solutions by boiling,
       whilst gases which are only slightly soluble evolve very much
       less heat.

It is evident that the conception of the partial pressures of gases
should be applied not only to the formations of solutions, but also
to all cases of chemical action of gases. Especially numerous are its
applications to the physiology of respiration, for in these cases it is
only the oxygen of the atmosphere that acts.[38]

  [38] Among the numerous researches concerning this subject, certain
       results obtained by Paul Bert are cited in Chapter III., and we
       will here point out that Prof. Sechenoff, in his researches on
       the absorption of gases by liquids, very fully investigated the
       phenomena of the solution of carbonic anhydride in solutions
       of various salts, and arrived at many important results, which
       showed that, on the one hand, in the solution of carbonic
       anhydride in solutions of salts on which it is capable of acting
       chemically (for example, sodium carbonate, borax, ordinary sodium
       phosphate), there is not only an increase of solubility, but also
       a distinct deviation from the law of Henry and Dalton; whilst, on
       the other hand, that solutions of salts which are not acted on
       by carbonic anhydride (for example, the chlorides, nitrates, and
       sulphates) absorb less of it, owing to the 'competition' of the
       salt already dissolved, and follow the law of Henry and Dalton,
       but at the same time show undoubted signs of a chemical action
       between the salt, water, and carbonic anhydride. Sulphuric acid
       (whose co-efficient of absorption is 92 vols. per 100), when
       diluted with water, absorbs less and less carbonic anhydride,
       until the hydrate H_{2}SO_{4},H_{2}O (co-eff. of absorption then
       equals 66 vols.) is formed; then on further addition of water the
       solubility again rises until a solution of 100 p.c. of water is
       obtained.

The solution of _solids_, whilst depending only in a small measure
on the pressure under which solution takes place (because solids and
liquids are almost incompressible), is very clearly dependent on the
temperature. In the great majority of cases the solubility of solids
in water increases with the temperature; and further, the rapidity of
solution increases also. The latter is determined by the rapidity of
diffusion of the solution formed into the remainder of the water. The
solution of a solid in water, although it is as with gases, a physical
passage into a liquid state, is determined, however, by its chemical
affinity for water; this is clearly shown from the fact that in solution
there occurs a diminution in volume, a change in the boiling point of
water, a change in the tension of its vapour, in the freezing point,
and in many similar properties. If solution were a physical, and not a
chemical, phenomenon, it would naturally be accompanied by an increase
and not by a diminution of volume, because generally in melting a solid
increases in volume (its density diminishes). _Contraction_ is the usual
phenomenon accompanying solution and takes place even in the addition
of solutions to water,[39] and in the solution of liquids in water,[40]
just as happens in the combination of substances when evidently new
substances are produced.[41] The contraction which takes place in
solution is, however, very small, a fact which depends on the small
compressibility of solids and liquids, and on the insignificance of the
compressing force acting in solution.[42] The change of volume which
takes place in the solution of solids and liquids, or the alteration in
specific gravity[43] corresponding with it, depends on peculiarities of
the dissolving substances, and of water, and, in the majority of cases,
is not proportional to the quantity of the substance dissolved,[44]
showing the existence of a chemical force between the solvent and the
substance dissolved which is of the same nature as in all other forms of
chemical reaction.[45]

  [39] Kremers made this observation in the following simple form:--He
       took a narrow-necked flask, with a mark on the narrow part (like
       that on a litre flask which is used for accurately measuring
       liquids), poured water into it, and then inserted a funnel,
       having a fine tube which reached to the bottom of the flask.
       Through this funnel he carefully poured a solution of any salt,
       and (having removed the funnel) allowed the liquid to attain a
       definite temperature (in a water bath); he then filled the flask
       up to the mark with water. In this manner two layers of liquid
       were obtained, the heavy saline solution below and water above.
       The flask was then shaken in order to accelerate diffusion, and
       it was observed that the volume became less if the temperature
       remained constant. This can be proved by calculation, if the
       specific gravity of the solutions and water be known. Thus at 15°
       one c.c. of a 20 p.c. solution of common salt weighs 1·1500 gram,
       hence 100 grams occupy a volume of 86·96 c.c. As the sp. gr.
       of water at 15° = 0·99916, therefore 100 grams of water occupy
       a volume of 100·08 c.c. The sum of the volumes is 187·04 c.c.
       After mixing, 200 grams of a 10 p.c. solution are obtained. Its
       specific gravity is 1·0725 (at 15° and referred to water at its
       maximum density), hence the 200 grams will occupy a volume of
       186·48 c.c. The contraction is consequently equal to 0·56 c.c.

  [40] The contractions produced in the case of the solution of sulphuric
       acid in water are shown in the diagram Fig. 17 (page 77). Their
       maximum is 10·1 c.c. per 100 c.c. of the solution formed. A
       maximum contraction of 4·15 at 0°, 3·78 at 15°, and 3·50 at 30°,
       takes place in the solution of 46 parts by weight of anhydrous
       alcohol in 54 parts of water. This signifies that if, at 0°, 46
       parts by weight of alcohol be taken per 54 parts by weight of
       water, then the sum of their separate volumes will he 104·15, and
       after mixing their total volume will be 100.

  [41] This subject will be considered later in this work, and we shall
       then see that the contraction produced in reactions of
       combination (of solids or liquids) is very variable in its
       amount, and that there are, although rarely, reactions of
       combination in which contraction does not take place, or when an
       increase of volume is produced.

  [42] The compressibility of solutions of common salt is less, according
       to Grassi, than that of water. At 18° the compression of water
       per million volumes = 48 vols. for a pressure of one atmosphere;
       for a 15 p.c. solution of common salt it is 32, and for a 24
       p.c. solution 26 vols. Similar determinations were made by Brown
       (1887) for saturated solutions of sal ammoniac (38 vols.), alum
       (46 vols.), common salt (27 vols.), and sodium sulphate at +1°,
       when the compressibility of water = 47 per million volumes.
       This investigator also showed that substances which dissolve
       with an evolution of heat and with an increase in volume (as,
       for instance, sal ammoniac) are partially separated from their
       saturated solutions by an increase of pressure (this experiment
       was particularly conclusive in the case of sal ammoniac), whilst
       the solubility of substances which dissolve with an absorption
       of heat or diminution in volume increases, although very
       slightly, _with an increase of pressure_. Sorby observed the same
       phenomenon with common salt (1863).

  [43] The most trustworthy data relating to the variation of the
       specific gravity of solutions with a change of their composition
       and temperature, are collected and discussed in my work cited
       in footnote 19. The practical (for the amount of a substance in
       solution is determined by the aid of the specific gravities of
       solutions, both in works and in laboratory practice) and the
       theoretical (for specific gravity can be more accurately observed
       than other properties, and because a variation in specific
       gravity governs the variation of many other properties) interest
       of this subject, besides the strict rules and laws to which it
       is liable, make one wish that this province of data concerning
       solutions may soon be enriched by further observations of as
       accurate a nature as possible. Their collection does not present
       any great difficulty, although requiring much time and attention.
       Pickering in London and Tourbaba in Kharkoff must be ranked first
       among those who have pursued problems of this nature during
       recent years.

  [44] Inasmuch as the degree of change exhibited in many properties on
       the formation of solutions is not large, so, owing to the
       insufficient accuracy of observations, a proportionality between
       this change and a change of composition may, in a first rough
       approximation and especially within narrow limits of change
       of composition, easily be imagined in cases where it does not
       even exist. The conclusion of Michel and Kraft is particularly
       instructive in this respect; in 1854, on the basis of their
       incomplete researches, they supposed that the increment of the
       specific gravity of solutions was proportional to the increment
       of a salt in a given volume of a solution, which is only true
       for determinations of specific gravity which are exact to
       the second decimal place--an accuracy insufficient even for
       technical determinations. Accurate measurements do not confirm
       a proportionality either in this case or in many others where
       a ratio has been generally accepted; as, for example, for the
       rotatory power (with respect to the plane of polarisation)
       of solutions, and for their capillarity, &c. Nevertheless,
       such a method is not only still made use of, but even has its
       advantages when applied to solutions within a limited scope--as,
       for instance, very weak solutions, and for a first acquaintance
       with the phenomena accompanying solution, and also as a means
       for facilitating the application of mathematical analysis to
       the investigation of the phenomenon of solution. Judging by the
       results obtained in my researches on the specific gravity of
       solutions, I think that in many cases it would be nearer the
       truth to take the change of properties as proportional, not to
       the amount of a substance dissolved, but to the product of this
       quantity and the amount of water in which it is dissolved; the
       more so since many chemical relations vary in proportion to the
       reacting masses, and a similar ratio has been established for
       many phenomena of attraction studied by mechanics. This product
       is easily arrived at when the quantity of water in the solutions
       to be compared is constant, as is shown in investigating the fall
       of temperature in the formation of ice (_see_ footnote 49, p. 91).

  [45] All the different forms of chemical reaction may be said to take
       place in the process of solution. (1) _Combinations_ between
       the solvent and the substance dissolved, which are more or less
       stable (more or less dissociated). This form of reaction is the
       most probable, and is that most often observed. (2) Reactions
       of _substitution_ or of _double decomposition_ between the
       molecules. Thus it may be supposed that in the solution of
       sal ammoniac, NH_{4}Cl, the action of water produces ammonia,
       NH_{4}HO, and hydrochloric acid, HCl, which are dissolved in
       the water and simultaneously attract each other. As these
       solutions and many others do indeed exhibit signs, which are
       sometimes indisputable, of similar double decompositions (thus
       solutions of sal-ammoniac yield a certain amount of ammonia),
       it is probable that this form of reaction is more often met
       with than is generally thought. (3) Reactions of _isomerism_ or
       _replacement_ are also probably met with in solution, all the
       more as here molecules of different kinds come into intimate
       contact, and it is very likely that the configuration of the
       atoms in the molecules under these influences is somewhat
       different from what it was in its original and isolated state.
       One is led to this supposition especially from observations made
       on solutions of substances which rotate the plane of polarisation
       (and observations of this kind are very sensitive with respect
       to the atomic structure of molecules), because they show, for
       example (according to Schneider, 1881), that strong solutions
       of malic acid rotate the plane of polarisation to the right,
       whilst its ammonium salts in all degrees of concentration
       rotate the plane of polarisation to the left. (4) Reactions of
       _decomposition_ under the influences of solution are not only
       rational in themselves, but have in recent years been recognised
       by Arrhenius, Ostwald, and others, particularly on the basis of
       electrolytic determinations. If a portion of the molecules of a
       solution occur in a condition of decomposition, the other portion
       may occur in a yet more complex state of combination, just as the
       velocity of the motion of different gaseous molecules may be far
       from being the same (_see_ Note 34, p. 81).

       It is, therefore, very probable that the reactions taking place
       in solution vary both quantitatively and qualitatively with
       the mass of water in the solution, and the great difficulty
       in arriving at a definite conclusion as to the nature of the
       chemical relations which take place in the process of solution
       will be understood, and if besides this the existence of a
       physical process, like the sliding between and interpenetration
       of two homogeneous liquids, be also recognised in solution,
       then the complexity of the problem as to the actual nature of
       solutions, which is now to the fore, appears in its true light.
       However, the efforts which are now being applied to the solution
       of this problem are so numerous and of such varied aspect that
       they will afford future investigators a vast mass of material
       towards the construction of a complete theory of solution.

       For my part, I am of opinion that the study of the physical
       properties of solutions (and especially of weak ones) which now
       obtains, cannot give any fundamental and complete solution of
       the problem whatever (although it should add much to both the
       provinces of physics and chemistry), but that, parallel with it,
       should be undertaken the study of the influence of temperature,
       and especially of low temperatures, the application to solutions
       of the mechanical theory of heat, and the comparative study of
       the chemical properties of solutions. The beginning of all this
       is already established, but it is impossible to consider in so
       short an exposition of chemistry the further efforts of this kind
       which have been made up to the present date.

The feeble development of the chemical affinities acting in solutions of
solids becomes evident from those multifarious methods by which _their
solutions are decomposed_, whether they be saturated or not. On heating
(absorption of heat), on cooling, and by internal forces alone, aqueous
solutions in many cases separate into their components or their definite
compounds with water. The water contained in solutions is removed
from them as vapour, or, by freezing, in the form of ice,[46] but the
_tension of the vapour of water_[47] held in solution is less than
that of water in a free state, and the _temperature of the formation
of ice_ from solutions is lower than 0°. Further, both the diminution
of vapour tension and the lowering of the freezing point proceed, in
dilute solutions, almost in proportion to the amount of a substance
dissolved.[48] Thus, if per 100 grams of water there be in solution
1, 5, 10 grams of common salt (NaCl), then at 100° the vapour tension
of the solutions decreases by 4, 21, 43 mm. of the barometric column,
against 760 mm., or the vapour tension of water, whilst the freezing
points are -0·58°, -2·91°, and -6·10° respectively. The above figures[49]
are almost proportional to the amounts of salt in solution (1, 5, and
10 per 100 of water). Furthermore, it has been shown by experiment that
the ratio of the diminution of vapour tension to the vapour tension
of water at different temperatures in a given solution is an almost
constant quantity,[50] and that for every (dilute) solution the ratio
between the diminution of vapour tension and of the freezing point is
also a tolerably constant quantity.[51]

  [46] If solutions are regarded as being in a state of dissociation
       (_see_ footnote 19, p. 64) it would be expected that they would
       contain free molecules of water, which form one of the products
       of the decomposition of those definite compounds whose formation
       is the cause of solution. In separating as ice or vapour, water
       makes, with a solution, a heterogeneous system (made up of
       substances in different physical states) similar, for instance,
       to the formation of a precipitate or volatile substance in
       reactions of double decomposition.

  [47] If the substance dissolved is non-volatile (like salt or sugar),
       or only slightly volatile, then the whole of the tension of
       the vapour given off is due to the water, but if a solution
       of a volatile substance--for instance, a gas or a volatile
       liquid--evaporates, then only a portion of the pressure belongs
       to the water, and the whole pressure observed consists of the sum
       of the pressures of the vapours of the water and of the substance
       dissolved. The majority of researches bear on the first case,
       which will be spoken of presently, and the observations of D. P.
       Konovaloff (1881) refer to the second case. He showed that in the
       case of two volatile liquids, mutually soluble in each other,
       forming two layers of saturated solutions (for example, ether
       and water, Note 20, p. 67), both solutions have an equal vapour
       tension (in the case in point the tension of both is equal to 431
       mm. of mercury at 19·8°). Further, he found that for solutions
       which are formed in all proportions, the tension is either
       greater (solutions of alcohol and water) or less (solutions of
       formic acid) than that which answers to the rectilinear change
       (proportional to the composition) from the tension of water
       to the tension of the substance dissolved; thus, the tension,
       for example, of a 70 p.c. solution of formic acid is less, at
       all temperatures, than the tension of water and of formic acid
       itself. In this case the tension of a solution is never equal to
       the sum of the tensions of the dissolving liquids, as Regnault
       already showed when he distinguished this case from that in
       which a mixture of liquids, which are insoluble in each other,
       evaporates. From this it is evident that a mutual action occurs
       in solution, which diminishes the vapour tensions proper to the
       individual substances, as would be expected on the supposition of
       the formation of compounds in solutions, because the elasticity
       then always diminishes.

  [48] This amount is usually expressed by the weight of the substance
       dissolved per 100 parts by weight of water. Probably it would
       be better to express it by the quantity of the substance in a
       definite volume of the solution--for instance, in a litre--or
       by the ratios of the number of molecules of water and of the
       substance dissolved.

  [49] The variation of the vapour tension of solutions has been
       investigated by many. The best known researches are those of
       Wüllner in Germany (1858-1860) and of Tamman in Russia (1887).
       The researches on the temperature of the formation of ice from
       various solutions are also very numerous; Blagden (1788),
       Rüdorff (1861), and De Coppet (1871) established the beginning,
       but this kind of investigation takes its chief interest from
       the work of Raoult, begun in 1882 on aqueous solutions, and
       afterwards continued for solutions in various other easily frozen
       liquids--for instance, benzene, C_{6}H_{6} (melts at 4·96°),
       acetic acid, C_{2}H_{4}O_{2} (16·75°), and others. An especially
       important interest is attached to these cryoscopic investigations
       of Raoult in France on the depression of the freezing point,
       because he took solutions of many well-known carbon-compounds
       and discovered a simple relation between the molecular weight
       of the substances and the temperature of crystallisation of the
       solvent, which enabled this kind of research to be applied to
       the investigation of the nature of substances. We shall meet
       with the application of this method later on (_see also_ Chapter
       VII.), and at present will only cite the deduction arrived at
       from these results. The solution of one-hundredth part of that
       molecular gram weight which corresponds with the formula of a
       substance dissolved (for example, NaCl = 58·5, C_{2}H_{6}O = 46,
       &c.) in 100 parts of a solvent lowers the freezing point of its
       solution in water 0·185°, in benzene 0·49°, and in acetic acid
       O·39°, or twice as much as with water. And as in weak solutions
       the depression or fall of freezing point is proportional to the
       amount of the substance dissolved, it follows that the fall of
       freezing point for all other solutions may be calculated from
       this rule. So, for instance, the weight which corresponds with
       the formula of acetone, C_{3}H_{6}O is 58; a solution containing
       2·42, 6·22, and 12·35 grams of acetone per 100 grams of water,
       forms ice (according to the determinations of Beckmann) at
       0·770°, 1·930°, and 3·820°, and these figures show that with
       a solution containing 0·58 gram of acetone per 100 of water
       the fall of the temperature of the formation of ice will be
       0·185°, 0·180°, and 0·179°. It must be remarked that the law of
       proportionality between the fall of temperature of the formation
       of ice, and the composition of a solution, is in general only
       approximate, and is only applicable to weak solutions (Pickering
       and others).

       We will here remark that the theoretical interest of this subject
       was strengthened on the discovery of the connection existing
       between the fall of tension, the fall of the temperature of the
       formation of ice, of osmotic pressure (Van't Hoff, Note 19),
       and of the electrical conductivity of solutions, and we will
       therefore supplement what we have already said on the subject by
       some short remarks on the method of cryoscopic investigations,
       although the details of the subject form the subject of more
       special works on physical chemistry (such as Ostwald's _Lehrbuch
       der allgemeinen Chemie_, 1891-1894, 2 vols.)

       In order to determine the _temperature of the formation of
       ice_ (or of crystallisation of other solvents), a solution
       of known strength is prepared and poured into a cylindrical
       vessel surrounded by a second similar vessel, leaving a layer
       of air between the two, which, being a bad conductor, prevents
       any rapid change of temperature. The bulb of a sensitive and
       corrected thermometer is immersed in the solution, and also a
       bent platinum wire for stirring the solution; the whole is then
       cooled (by immersing the apparatus in a freezing mixture), and
       the temperature at which ice begins to separate observed. If
       the temperature at first falls slightly lower, it nevertheless
       becomes constant when ice begins to form. By then allowing the
       liquid to get just warm, and again observing the temperature of
       the formation of ice, an exact determination may be arrived at.
       It is still better to take a large mass of solution, and induce
       the formation of the first crystals by dropping a small lump
       of ice into the solution already partially over-cooled. This
       only imperceptibly changes the composition of the solution. The
       observation should be made at the point of formation of only a
       very small amount of crystals, as otherwise the composition of
       the solution will become altered from their separation. Every
       precaution must be taken to prevent the access of moisture to the
       interior of the apparatus, which might also alter the composition
       of the solution or properties of the solvent (for instance, when
       using acetic acid).

       With respect to the depression of dilute solutions it is
       known--(1) That the depression increases in almost direct
       proportion to the amount of the substance in solution (always
       per 100 parts of water), for example, for KCl when the solution
       contains 1 part of salt (per 100 parts of water) the depression =
       0·45°, when the solution contains 2 parts of salt = 0·90°, with
       10 parts of salt = 4·4°. (2) The greater the molecular weight
       expressed by the formula (see Chapter VII.), and designated
       by M, the less, under other similar conditions, will be the
       depression _d_, and therefore if the concentration of a solution
       (the amount by weight of substance dissolved per 100 parts of
       water) be designated by _p_, then the fraction M_d_/_p_ or the
       molecular depression for a given class of substances will be a
       constant quantity; for example, in the case of methyl alcohol in
       water 17·3, for acetone about 18·0, for sugar about 18·5. (3) In
       general the molecular depression for substances whose solutions
       do not conduct an electric current is about 18·5, while for
       acids, salts, and such like substances whose solutions do conduct
       electricity, it is _i_ times greater; for instance, for HCl, KI,
       HNO_{3}, KHO, &c., about 36 (_i_ is nearly 2), for borax about
       66, and so on where _i_ varies in the same manner as it does in
       the case of the osmotic pressure of solutions (Note 19). (4)
       Different solvents (water, acetic acid, benzene, &c.) have each
       their corresponding constants of molecular depression (which have
       a certain remote connection with their molecular weight); for
       example, for acetic acid the molecular depression is about 39 and
       not 19 (as it is for water), for benzene 49, for methyl alcohol
       about 17, &c. (5) If the molecular weight M of a substance be
       unknown, then in the case of non-conductors of electricity or for
       a given group, it may be found by determining the depression,
       _d_, for a given concentration, _p_; for example, in the case of
       peroxide of hydrogen, which is a non-conductor of electricity,
       the molecular weight, M, was found to be nearly 34, _i.e._ equal
       to H_{2}O_{2}.

       Similar results have also been found for the fall in the
       vapour tension of solutions (Note 51), and for the rise of
       their boiling points (hence these data may also serve for
       determining the molecular weight of a substance in solution,
       as is shortly described in Chapter VII., Note 27 bis). And as
       these conclusions are also applicable in the case of osmotic
       pressure (Note 19), and a variation in the magnitude of _i_, in
       passing from solutions which do not conduct an electric current
       to those which do conduct electricity is everywhere remarked,
       so it was natural to here seek that causal connection which
       Arrhenius (1888), Ostwald, and others expected to find in the
       supposition that a portion of the substance of the electrolyte
       is already decomposed in the very act of solution, into its
       ions (for example, NaCl into Na and Cl), or into the atoms of
       those individual substances which make their appearance in
       electrolysis, and in this way to explain the fact that _i_ is
       greater for those bodies which conduct an electric current.
       We will not consider here this supposition, known as the
       hypothesis of 'electrolytic dissociation,' not only because
       it wholly belongs to that special branch--physical chemistry,
       and gives scarcely any help towards explaining the chemical
       relations of solutions (particularly their passage into definite
       compounds, their reactions, and their very formation), but
       also because--(1) all the above data (for constant depression,
       osmotic pressure, &c.) only refer to dilute solutions, and are
       not applicable to strong solutions; whilst the chemical interest
       in strong solutions is not less than in dilute solutions, and
       the transition from the former into the latter is consecutive
       and inevitable; (2) because in all homogeneous bodies (although
       it may be insoluble and not an electrolyte) a portion of the
       atoms may he supposed (Clausius) to be passing from one particle
       to another (Chapter X., Note 28), and as it were dissociated,
       but there are no reasons for believing that such a phenomenon
       is proper to the solutions of electrolytes only; (3) because no
       essential mark of difference is observed between the solution of
       electrolytes and non-conductors, although it might be expected
       there would be according to Arrhenius' hypothesis; (4) because it
       is most reasonable to suppose the formation of new, more complex,
       but unstable and easily dissociated compounds in the act of
       solution, than a decomposition, even partial, of the substances
       taken; (5) because if Arrhenius' hypothesis be accepted it
       becomes necessary to admit the existence in solutions of free
       ions, like the atoms Cl or Na, without any apparent expenditure
       of the energy necessary for their disruption, and if in this case
       it can be explained why _i_ then = 2, it is not at all clear why
       solutions of MgSO_{4} give _i_ = 1, although the solution does
       conduct an electric current; (6) because in dilute solutions,
       the approximative proportionality between the depression and
       concentration may be recognised, while admitting the formation
       of hydrates, with as much right as in admitting the solution
       of anhydrous substances, and if the formation of hydrates be
       recognised it is easier to admit that a portion of these hydrates
       is decomposed than to accept the breaking-up into ions; (7)
       because the best conductors of electricity are solutions like the
       sulphates in which it is necessary to recognise the formation
       of associated systems or hydrates; (8) because the cause of
       electro-conductivity can be sooner looked for in this affinity
       and this combination of the substance dissolved with the solvent,
       as is seen from the fact, that (D. P. Konovaloff) neither aniline
       nor acetic acid alone conduct an electric current, a solution
       of aniline in water conducts it badly (and here the affinity is
       very small), while a solution of aniline in acetic acid forms a
       good electrolyte, in which, without doubt, chemical forces are
       acting, bringing aniline, like ammonia, into combination with the
       acetic acid; which is evident from the researches made by Prof.
       Konovaloff upon mixtures (solutions) of aniline and other amines;
       and, lastly, (9) because I, together with many of the chemists
       of the present day, cannot regard the hypothesis of electrolytic
       dissociation in the form given to it up to now by Arrhenius and
       Ostwald, as answering to the sum total of the chemical data
       respecting solutions and dissociation in general. Thus, although
       I consider it superfluous to discuss further the evolution of
       the above theory of solutions, still I think that it would he
       most useful for students of chemistry to consider all the data
       referring to this subject, which can be found in the _Zeitschrift
       für physikalische Chemie_, 1888-1894.

  [50] This fact, which was established by Gay-Lussac, Pierson, and
       v. Babo, is confirmed by the latest observations, and enables us
       to express not only the fall of tension (_p_-_p_´) itself, but
       its ratio to the tension of water (_p_-_p_´)/_p_. It is to be
       remarked that in the absence of any chemical action, the fall of
       pressure is either very small, or does not exist at all (note
       33), and is not proportional to the quantity of the substance
       added. As a rule, the tension is then equal, according to the law
       of Dalton, to the sum of the tensions of the substances taken.
       Hence liquids which are insoluble in each other (for example,
       water and chloride of carbon) present a tension equal to the sum
       of their individual tensions, and therefore such a mixture boils
       at a lower temperature than the more volatile liquid (Magnus,
       Regnault).

  [51] If, in the example of common salt, the fall of tension be divided
       by the tension of water, a figure is obtained which is nearly
       105 times less than the magnitude of the fall of temperature of
       formation of ice. This correlation was theoretically deduced
       by Goldberg, on the basis of the application of the mechanical
       theory of heat, and is repeated by many investigated solutions.

The diminution of the vapour tension of solutions explains the rise in
boiling point due to the solution of solid non-volatile bodies in water.
The temperature of a vapour is the same as that of the solution from
which it is generated, and therefore it follows that the aqueous vapour
given off from a solution will be superheated. A saturated solution
of common salt boils at 108·4°, a solution of 335 parts of nitre in
100 parts of water at 115·9°, and a solution of 325 parts of potassium
chloride in 100 parts of water at 179°, if the temperature of ebullition
be determined by immersing the thermometer bulb in the liquid itself.
This is another proof of the bond which exists between water and the
substance dissolved. And this bond is seen still more clearly in those
cases (for example, in the solution of nitric or formic acid in water)
where the solution boils at a higher temperature than either water or
the volatile substance dissolved in it. For this reason the solutions of
certain gases--for instance, hydriodic or hydrochloric acid--boil above
100°.

The separation of ice from solutions[52] explains both the phenomenon,
well known to sailors, that the ice formed from salt water gives fresh
water, and also the fact that by freezing, just as by evaporation, a
solution is obtained which is richer in salts than before. This is taken
advantage of in cold countries for obtaining a liquor from sea water,
which is then evaporated for the extraction of salt.

  [52] Fritzsche showed that solutions of certain colouring matters yield
       colourless ice, which clearly proves the passage of water only
       into a solid state, without any intermixture of the substance
       dissolved, although the possibility of the admixture in certain
       other cases cannot be denied.

On the removal of part of the water from a solution (by evaporation
or the separation of ice), a saturated solution should be obtained,
and then the solid substance dissolved should separate out. Solutions
saturated at a certain temperature should also separate out a
corresponding portion of the substance dissolved if they be reduced,
by cooling,[53] to a temperature at which the water can no longer hold
the former quantity of the substance in solution. If this separation,
by cooling a saturated solution or by evaporation, take place slowly,
_crystals_ of the substance dissolved are in many cases formed; and
this is the method by which crystals of soluble salts are usually
obtained. Certain solids very easily separate out from their solutions
in perfectly formed crystals, which may attain very large dimensions.
Such are nickel sulphate, alum, sodium carbonate, chrome-alum, copper
sulphate, potassium ferricyanide, and a whole series of other salts.
The most remarkable circumstance in this is that many solids in
separating out from an aqueous solution retain a portion of water,
forming crystallised solid substances which contain water. A portion
of the water previously in the solution remains in the separated
crystals. The water which is thus retained is called the _water of
crystallisation_. Alum, copper sulphate, Glauber's salt, and magnesium
sulphate contain such water, but neither sal-ammoniac, table salt,
nitre, potassium chlorate, silver nitrate, nor sugar, contains any
water of crystallisation. One and the same substance may separate out
from a solution with or without water of crystallisation, according
to the temperature at which the crystals are formed. Thus common salt
in crystallising from its solution in water at the ordinary or at a
higher temperature does not contain water of crystallisation. But if
its separation from the solution takes place at a low temperature,
namely below -5°, then the crystals contain 38 parts of water in 100
parts. Crystals of the same substance which separate out at different
temperatures may contain different amounts of water of crystallisation.
This proves to us that a solid dissolved in water may form various
compounds with it, differing in their properties and composition,
and capable of appearing in a solid separate form like many ordinary
definite compounds. This is indicated by the numerous properties and
phenomena connected with solutions, and gives reason for thinking that
there exist in solutions themselves such compounds of the substance
dissolved, and the solvent or compounds similar to them, only in a
liquid partly decomposed form. Even the _colour of solutions_ may
often confirm this opinion. Copper sulphate forms crystals having a
blue colour and containing water of crystallisation. If the water
of crystallisation be removed by heating the crystals to redness, a
colourless anhydrous substance is obtained (a white powder). From this
it may be seen that the blue colour belongs to the compound of the
copper salt with water. Solutions of copper sulphate are all blue, and
consequently they contain a compound similar to the compound formed by
the salt with its water of crystallisation. Crystals of cobalt chloride
when dissolved in an anhydrous liquid--like alcohol, for instance--give
a blue solution, but when they are dissolved in water a red solution is
obtained. Crystals from the aqueous solution, according to Professor
Potilitzin, contain six times as much water (CoCl_{2},6H_{2}O) for a
given weight of the salt, as those violet crystals (CoCl_{2},H_{2}O)
which are formed by the evaporation of an alcoholic solution.

  [53] As the solubility of certain substances (for example, coniine,
       cerium sulphate, and others) decreases with a rise of temperature
       (between certain limits--see, for example, note 24), so these
       substances do not separate from their saturated solutions on
       cooling but on heating. Thus a solution of manganese sulphate,
       saturated at 70°, becomes cloudy on further heating. The point at
       which a substance separates from its solution with a change of
       temperature gives an easy means of determining the co-efficient
       of solubility, and this was taken advantage of by Prof. Alexéeff
       for determining the solubility of many substances. The phenomenon
       and method of observation are here essentially the same as in
       the determination of the temperature of formation of ice. If a
       solution of a substance which separates out on heating be taken
       (for example, the sulphate of calcium or manganese), then at a
       certain fall of temperature ice will separate out from it, and
       at a certain rise of temperature the salt will separate out.
       From this example, and from general considerations, it is clear
       that the separation of a substance dissolved from a solution
       should present a certain analogy to the separation of ice from a
       solution. In both cases, a heterogeneous system of a solid and a
       liquid is formed from a homogeneous (liquid) system.

That solutions contain particular compounds with water is further shown
by the phenomena of supersaturated solutions, of so-called cryohydrates,
of solutions of certain acids having constant boiling points, and the
properties of compounds containing water of crystallisation whose data
it is indispensable to keep in view in the consideration of solutions.

Supersaturated solutions exhibit the following phenomena:--On the
refrigeration of a saturated solution of certain salts,[54] if the
liquid be brought under certain conditions, the excess of the solid
may sometimes remain in solution and not separate out. A great number
of substances, and more especially sodium sulphate, Na_{2}SO_{4}, or
Glauber's salt, easily form supersaturated solutions. If boiling water
be saturated with this salt, and the solution be poured off from any
remaining undissolved salt, and, the boiling being still continued,
the vessel holding the solution be well closed by cotton wool, or by
fusing up the vessel, or by covering the solution with a layer of oil,
then it will he found that this saturated solution does not separate
out any Glauber's salt whatever on cooling down to the ordinary or even
to a much lower temperature; although without the above precautions a
salt separates out on cooling, in the form of crystals, which contain
Na_{2}SO_{4},10H_{2}O--that is, 180 parts of water for 142 parts of
anhydrous salt. The supersaturated solution may be moved about or
shaken inside the vessel holding it, and no crystallisation will take
place; the salt remains in the solution in as large an amount as at a
higher temperature. If the vessel holding the supersaturated solution
be opened and a crystal of Glauber's salt be thrown in, crystallisation
suddenly takes place.[55] A considerable rise in temperature is noticed
during this rapid separation of crystals, which is due to the fact
that the salt, previously in a liquid state, passes into a solid state.
This bears some resemblance to the fact that water maybe cooled below
0° (even to -10°) if it be left at rest, under certain circumstances,
and evolves heat in suddenly crystallising. Although from this point
of view there is a resemblance, yet in reality the phenomenon of
supersaturated solutions is much more complicated. Thus, on cooling,
a saturated solution of Glauber's salt deposits crystals containing
Na_{2}SO_{4},7H_{2}0,[56] or 126 parts of water per 142 parts of
anhydrous salt, and not 180 parts of water, as in the above-mentioned
salt. The crystals containing 7H_{2}O are distinguished for their
instability; if they stand in contact not only with crystals of
Na_{2}SO_{4},10H_{2}O, but with many other substances, they immediately
become opaque, forming a mixture of anhydrous and deca-hydrated salts.
It is evident that between water and a soluble substance there may be
established different kinds of greater or less stable equilibrium, of
which solutions form a particular case.[57]

  [54] Those salts which separate out with water of crystallisation and
       give several crystallohydrates form supersaturated solutions with
       the greatest facility, and the phenomenon is much more common
       than was previously imagined. The first data were given in the
       last century by Loewitz, in St. Petersburg. Numerous researches
       have proved that supersaturated solutions do not differ from
       ordinary solutions in any of their essential properties. The
       variations in specific gravity, vapour tension, formation of ice,
       &c., take place according to the ordinary laws.

  [55] Inasmuch as air, as has been shown by direct experiment, contains,
       although in very small quantities, minute crystals of salts,
       and among them sodium sulphate, air can bring about the
       crystallisation of a supersaturated solution of sodium sulphate
       in an open vessel, but it has no effect on saturated solutions of
       certain other salts; for example, lead acetate. According to the
       observations of De Boisbaudran, Gernez, and others, isomorphous
       salts (analogous in composition) are capable of inducing
       crystallisation. Thus, a supersaturated solution of nickel
       sulphate crystallises by contact with crystals of sulphates of
       other metals analogous to it, such as those of magnesium, cobalt,
       copper, and manganese. The crystallisation of a supersaturated
       solution, set up by the contact of a minute crystal, starts from
       it in rays with a definite velocity, and it is evident that the
       crystals as they form propagate the crystallisation in definite
       directions. This phenomenon recalls the evolution of organisms
       from germs. An attraction of similar molecules ensues, and they
       dispose themselves in definite similar forms.

  [56] At the present time a view is very generally accepted, which
       regards supersaturated solutions as homogeneous systems, which
       pass into heterogeneous systems (composed of a liquid and a
       solid substance), in all respects exactly resembling the passage
       of water cooled below its freezing point into ice and water,
       or the passage of crystals of rhombic sulphur into monoclinic
       crystals, and of the monoclinic crystals into rhombic. Although
       many phenomena of supersaturation are thus clearly understood,
       yet the spontaneous formation of the unstable hepta-hydrated salt
       (with 7H_{2}O), in the place of the more stable deca-hydrated
       salt (with mol. 10H_{2}O), indicates a property of a saturated
       solution of sodium sulphate which obliges one to admit that
       it has a different structure from an ordinary solution.
       Stcherbacheff asserts, on the basis of his researches, that
       a solution of the deca-hydrated salt gives, on evaporation,
       without the aid of heat, the deca-hydrated salt, whilst after
       heating above 33° it forms a supersaturated solution and the
       hepta-hydrated salt. But in order that this view should be
       accepted, some facts must be discovered distinguishing solutions
       (which are, according to this view, isomeric) containing the
       hepta-hydrated salt from those containing the deca-hydrated
       salt, and all efforts in this direction (the study of the
       properties of the solutions) have given negative results. As
       some crystallohydrates of salts (alums, sugar of lead, calcium
       chloride) melt straightway (without separating out anything),
       whilst others (like Na_{2}SO_{4},10H_{2}O) are broken up, then
       it may be that the latter are only in a state of equilibrium at
       a higher temperature than their melting point. It may here be
       observed that in melting crystals of the deca-hydrated salt,
       there is formed, besides the solid anhydrous salt, a saturated
       solution giving the hepta-hydrated salt, so that this passage
       from the deca-to the hepta-hydrated salt, and the reverse,
       takes place with the formation of the anhydrous (or, it may be,
       monohydrated) salt.

       Moreover, supersaturation (Potilitzin, 1889) only takes
       place with those substances which are capable of giving
       several modifications or several crystallohydrates, _i.e._
       supersaturated solutions separate out, besides the stable normal
       crystallohydrate, hydrates containing less water and also the
       anhydrous salt. This degree of saturation acts upon the substance
       dissolved in a like manner to heat. Sulphate of nickel in a
       solution at 15° to 20° separates out rhombic crystals with
       7H_{2}O, at 30° to 40° cubical crystals, with 6H_{2}O, at 50° to
       70° monoclinic crystals, also containing 6H_{2}O. Crystals of
       the same composition separate out from supersaturated solutions
       at one temperature (17° to 19°), but at different degrees of
       saturation, as was shown by Lecoq de Boisbaudran. The capacity
       to voluntarily separate out slightly hydrated or anhydrous salts
       by the introduction of a crystal into the solution is common to
       all supersaturated solutions. If a salt forms a supersaturated
       solution, then one would expect, according to this view, that
       it should exist in the form of several hydrates or in several
       modifications. Thus Potilitzin concluded that chlorate of
       strontium, which easily gives supersaturated solutions, should
       be capable of forming several hydrates, besides the anhydrous
       salt known; and he succeeded in discovering the existence
       of two hydrates, Sr(ClO_{3})_{2},3H_{2}O and apparently
       Sr(ClO_{3})_{2},8H_{2}O. Besides this, three modifications of the
       common anhydrous salt were obtained, differing from each other
       in their crystalline form. One modification separated out in the
       form of rhombic octahedra, another in oblique plates, and a third
       in long brittle prisms or plates. Further researches showed that
       salts which are not capable of forming supersaturated solutions
       such as the bromates of calcium, strontium, and barium, part
       with their water of hydration with difficulty (they crystallise
       with 1H_{2}O), and decompose very slowly in a vacuum or in dry
       air. In other words the tension of dissociation is very small
       in this class of hydrates. As the hydrates characterised by a
       small dissociation tension are incapable of giving supersaturated
       solutions, so conversely supersaturated solutions give hydrates
       whose tension of dissociation is great (Potilitzin, 1893).

  [57] _Emulsions_, like milk, are composed of a solution of glutinous
       or similar substances, or of oily liquids suspended in a
       liquid in the form of drops, which are clearly visible under a
       microscope, and form an example of a mechanical formation which
       resembles solution. But the difference from solutions is here
       evident. There are, however, solutions which approach very near
       to emulsions in the facility with which the substance dissolved
       separates from them. It has long been known, for example, that
       a particular kind of Prussian blue, KFe_{2}(CN)_{6}, dissolves
       in pure water, but, on the addition of the smallest quantity
       of either of a number of salts, it coagulates and becomes
       quite insoluble. If copper sulphide (CuS), cadmium sulphide
       (CdS), arsenic sulphide (As_{2}S_{5}) (the experiments with
       these substances proceed with great ease, and the solution
       obtained is comparatively stable), and many other metallic
       sulphides, be obtained by a method of double decomposition (by
       precipitating salts of these metals by hydrogen sulphide),
       and be then carefully washed (by allowing the precipitate to
       settle, pouring off the liquid, and again adding sulphuretted
       hydrogen water), then, as was shown by Schulze, Spring, Prost,
       and others, the previously insoluble sulphides pass into
       transparent (for mercury, lead, and silver, reddish brown; for
       copper and iron, greenish brown; for cadmium and indium, yellow;
       and for zinc, colourless) solutions, which may be preserved
       (the weaker they are the longer they keep) and even boiled, but
       which, nevertheless, in time coagulate--that is, separate in an
       insoluble form, and then sometimes become crystalline and quite
       incapable of re-dissolving. Graham and others observed the power
       shown by colloids (_see_ note 18) of forming similar _hydrosols
       or solutions of gelatinous colloids_, and, in describing alumina
       and silica, we shall again have occasion to speak of such
       solutions.

       In the existing state of our knowledge concerning solution, such
       solutions may be looked on as a transition between emulsion and
       ordinary solutions, but no fundamental judgment can be formed
       about them until a study has been made of their relations to
       ordinary solutions (the solutions of even soluble colloids freeze
       immediately on cooling below 0°, and, according to Guthrie, do
       not form cryohydrates), and to supersaturated solutions, with
       which they have certain points in common.

Solutions of salts on refrigeration below 0° deposit ice or crystals
(which then frequently contain water of crystallisation) of the salt
dissolved, and on reaching a certain degree of concentration they
solidify in their entire mass. These solidified masses are termed
_cryohydrates_. My researches on solutions of common salt (1868) showed
that its solution solidifies when it reaches a composition NaCl +
10H_{2}O (180 parts of water per 58·5 parts of salt), which takes place
at about -23°. The solidified solution melts at the same temperature,
and both the portion melted and the remainder preserve the above
composition. Guthrie (1874-1876) obtained the cryohydrates of many
salts, and he showed that certain of them are formed like the above at
comparatively low temperatures, whilst others (for instance, corrosive
sublimate, alums, potassium chlorate, and various colloids) are formed
on a slight cooling, to -2° or even before.[58] In the case of common
salt, the cryohydrate with 10 molecules of water, and in the case of
sodium nitrate, the cryohydrate[59] with 7 molecules of water (_i.e._
126 parts of water per 85 of salt) should be accepted as established
substances, capable of passing from a solid to a liquid state and
conversely; and therefore it may be thought that in cryohydrates we
have solutions which are not only undecomposable by cold, but also have
a definite composition which would present a fresh case of definite
equilibrium between the solvent and the substance dissolved.

  [58] Offer (1880) concludes, from his researches on cryohydrates, that
       they are simple mixtures of ice and salts, having a constant
       melting point, just as there are alloys having a constant point
       of fusion, and solutions of liquids with a constant boiling point
       (_see_ note 60). This does not, however, explain in what form
       a salt is contained, for instance, in the cryohydrate NaCl +
       10H_{2}O. At temperatures above -10° common salt separates out in
       anhydrous crystals, and at temperatures near -10°, in combination
       with water of crystallisation, NaCl + 2H_{2}O, and, therefore,
       it is very improbable that at still lower temperatures it would
       separate without water. If the possibility of the solidified
       cryohydrate containing NaCl + 2H_{2}O and ice be admitted, then
       it is not clear why one of these substances does not melt before
       the other. If alcohol does not extract water from the solid mass,
       leaving the salt behind, this does not prove the presence of ice,
       because alcohol also takes up water from the crystals of many
       hydrated substances (for instance, from NaCl + 2H_{2}O) at about
       their melting-points. Besides which, a simple observation on the
       cryohydrate, NaCl + 10H_{2}O, shows that with the most careful
       cooling it does not on the addition of ice deposit ice, which
       would occur if ice were formed on solidification intermixed with
       the salt.

       I may add with regard to cryohydrates that many of the solutions
       of acids solidify completely on prolonged cooling (for example,
       H_{2}SO_{4},H_{2}O), and then form perfectly definite compounds.
       For the solutions of sulphuric acid (_see_ Chapter XX.)
       Pickering obtained, for instance, a hydrate, H_{2}SO_{4},4H_{2}O
       at -25°. Hydrochloric, nitric, and other acids also give similar
       crystalline hydrates, melting at low temperatures and presenting
       many similarities with the cryohydrates.

  [59] _See_ note 24.

The formation of definite but unstable compounds in the process of
solution becomes evident from the phenomena of a marked decrease of
vapour tension, or from the rise of the temperature of ebullition which
occurs in the solution of certain volatile liquids and gases in water.
As an example, we will take hydriodic acid, HI, a gas which liquefies,
giving a liquid which boils at -20°. A solution of it containing 57 p.c.
of hydriodic acid is distinguished by the fact that if it be heated
the hydriodic acid volatilises together with the water in the same
proportions as they occur in the solution, therefore such a solution
may be distilled unchanged. The solution boils at a higher temperature
than water, at 127°. A portion of the physical properties of the gas
and water have in this case already disappeared--a new substance is
formed, which has its definite boiling point. To put it more correctly,
this is not the temperature of ebullition, but the temperature at which
the compound formed decomposes, forming the vapours of the products of
dissociation, which, on cooling, re-combine. Should a less amount of
hydriodic acid be dissolved in water than the above, then, on heating
such a solution, water only at first distils over, until the solution
attains the above-mentioned composition; it will then distil over
unaltered. If more hydriodic acid be passed into such a solution a fresh
quantity of the gas will dissolve, but it passes off with great ease,
like air from water. It must not, however, be thought that those forces
which determine the formation of ordinary gaseous solutions play no part
whatever in the formation of a solution having a definite boiling point;
that they do react is shown from the fact that such constant gaseous
solutions vary in their composition under different pressures.[60] It
is not, therefore, at every, but only at the ordinary, atmospheric
pressure that a constant boiling solution of hydriodic acid will contain
57 p.c. of the gas. At another pressure the proportion of water and
hydriodic acid will be different. It varies, however, judging from
observations made by Roscoe, very little for considerable variations of
pressure. This variation in composition directly indicates that pressure
exerts an influence on the formation of unstable chemical compounds
which are easily dissociated (with formation of a gas), just as it
influences the solution of gases, only the latter is influenced to a
more considerable degree than the former.[61] Hydrochloric, nitric, and
other acids form _solutions having definite boiling points_, like that
of hydriodic acid. They show further the common property, if containing
but a small proportion of water, that they _fume in air_. Strong
solutions of nitric, hydrochloric, hydriodic, and other gases are even
termed 'fuming acids.' The fuming liquids contain a definite compound
whose temperature of ebullition (decomposition) is higher than 100°,
and contain also an excess of the volatile substance dissolved, which
exhibits a capacity to combine with water and form a hydrate, whose
vapour tension is less than that of aqueous vapour. On evaporating in
air, this dissolved substance meets the atmospheric moisture and forms
a visible vapour (fumes) with it, which consists of the above-mentioned
compound. The attraction or affinity which binds, for instance,
hydriodic acid with water is evinced not only in the evolution of heat
and the diminution of vapour tension (rise of boiling point), but also
in many purely chemical relations. Thus hydriodic acid is produced from
iodine and hydrogen sulphide in the presence of water, but unless water
is present this reaction does not take place.[62]

  [60] For this reason (the want of entire constancy of the composition
       of constant boiling solutions with a change of pressure), the
       existence of definite hydrates formed by volatile substances--for
       instance, by hydrochloric acid and water--is frequently denied.
       It is generally argued as follows: If there did exist a constancy
       of composition, then it would be unaltered by a change of
       pressure. But the distillation of constant boiling hydrates
       is undoubtedly accompanied (judging by the vapour densities
       determined by Bineau), like the distillation of sal ammoniac,
       sulphuric acid, &c., by a complete decomposition of the original
       compound--that is, these substances do not exist in a state
       of vapour, but their products of decomposition (hydrochloric
       acid and water) are gases at the temperature of volatilisation,
       which dissolve in the volatilised and condensed liquids; but
       the solubility of gases in liquids depends on the pressure,
       and, therefore, the composition of constant boiling solutions
       may, and even ought to, vary with a change of pressure, and,
       further, the smaller the pressure and the lower the temperature
       of volatilisation, the more likely is a true compound to be
       obtained. According to the researches of Roscoe and Dittmar
       (1859), the constant boiling solution of hydrochloric acid
       proved to contain 18 p.c. of hydrochloric acid at a pressure of
       3 atmospheres, 20 p.c. at 1 atmosphere, and 23 p.c. at 1/10 of
       an atmosphere. On passing air through the solution until its
       composition became constant (_i.e._ forcing the excess of aqueous
       vapour or of hydrochloric acid to pass away with the air),
       then acid was obtained containing about 20 p.c. at 100°, about
       23 p.c. at 50°, and about 25 p.c. at 0°. From this it is seen
       that by decreasing the pressure and lowering the temperature of
       evaporation one arrives at the same limit, where the composition
       should be taken as HCl + 6H_{2}O, which requires 25·26 p.c. of
       hydrochloric acid. Fuming hydrochloric acid contains more than
       this.

       In the case already considered, as in the case of formic acid
       in the researches of D. P. Konovaloff (note 47), the constant
       boiling solution corresponds with a minimum tension--that
       is, with a boiling point higher than that of either of the
       component elements. But there is another case of constant boiling
       solutions similar to the case of the solution of propyl alcohol,
       C_{3}H_{8}O, when a solution, undecomposed by distillation, boils
       at a lower point than that of the more volatile liquid. However,
       in this case also, if there be solution, the possibility of the
       formation of a definite compound in the form C_{3}H_{8}O + H_{2}O
       cannot be denied, and the tension of the solution is not equal to
       the sum of tensions of the components. There are possible cases
       of constant boiling mixtures even when there is no solution nor
       any loss of tension, and consequently no chemical action, since
       the amount of liquids that are volatilised is determined by
       the product of the vapour densities into their vapour tensions
       (Wanklyn), in consequence of which liquids whose boiling point
       is above 100°--for instance, turpentine and ethereal oils in
       general--when distilled with aqueous vapour, pass over at a
       temperature below 100°. Consequently, it is not in the constancy
       of composition and boiling point (temperature of decomposition)
       that evidence of a distinct chemical action is to be found in
       the above-described solutions of acids, but in the great loss
       of tension, which completely resembles the loss of tension
       observed, for instance, in the perfectly-definite combinations of
       substances with water of crystallisation (see later, note 65).
       Sulphuric acid, H_{2}SO_{4}, as we shall learn later, is also
       decomposed by distillation, like HCl + 6H_{2}O, and exhibits,
       moreover, all the signs of a definite chemical compound. The
       study of the variation of the specific gravities of solutions as
       dependent on their composition (see note 19) shows that phenomena
       of a similar kind, although of different dimensions, take place
       in the formation of both H_{2}SO_{4} from H_{2}O and SO_{3}, and
       of HCl + 6H_{2}O (or of aqueous solutions analogous to it) from
       HCl and H_{2}O.

  [61] The essence of the matter may he thus represented. A gaseous or
       easily volatile substance _A_ forms with a certain quantity of
       water, _n_H_{2}O, a definite complex compound _An_H_{2}O, which
       is stable up to a temperature t° higher than 100°. At this
       temperature it is decomposed into two substances, _A_ + H_{2}O.
       Both boil below _t_° at the ordinary pressure, and therefore at
       _t_° they distil over and re-combine in the receiver. But if a
       part of the substance _An_H_{2}O is decomposed or volatilised,
       a portion of the undecomposed liquid still remains in the
       vessel, which can partially dissolve one of the products of
       decomposition, and that in quantity varying with the pressure and
       temperature, and therefore the solution at a constant boiling
       point will have a slightly different composition at different
       pressures.

  [62] For solutions of hydrochloric acid in water there are still greater
       differences in reactions. For instance, strong solutions
       decompose antimony sulphide (forming hydrogen sulphide, H_{2}S),
       and precipitate common salt from its solutions, whilst weak
       solutions do not act thus.

Many compunds containing water of crystallisation are solid substances
(when melted they are already solutions--_i.e._ liquids); furthermore,
they are capable of being formed from solutions, like ice or aqueous
vapour. They may be called _crystallo-hydrates_. Inasmuch as the direct
presence of ice or aqueous vapour cannot be admitted in solutions
(for these are liquids), although the presence of water may be, so
also there is no basis for acknowledging the presence in solutions
of crystallo-hydrates, although they are obtained from solutions as
such.[63] It is evident that such substances present one of the many
forms of equilibrium between water and a substance dissolved in it.
This form, however, reminds one, in all respects, of solutions--that
is, aqueous compounds which are more or less easily decomposed, with
separation of water and the formation of a less aqueous or an anhydrous
compound. In fact, there are not a few crystals containing water which
lose a part of their water at the ordinary temperature. Of such a kind,
for instance, are the crystals of soda, or sodium carbonate, which,
when separated from an aqueous solution at the ordinary temperature,
are quite transparent; but when left exposed to air, lose a portion
of their water, becoming opaque, and, in the process, lose their
crystalline appearance, although preserving their original form.
This process of the separation of water at the ordinary temperature
is termed the _efflorescence_ of crystals. Efflorescence takes place
more rapidly under the receiver of an air pump, and especially at a
gentle heat. This breaking up of a crystal is dissociation at the
ordinary temperature. Solutions are decomposed in exactly the same
manner.[64] The tension of the aqueous vapour which is given off
from crystallo-hydrates is naturally, as with solutions, less than
the vapour tension of water itself[65] at the same temperature, and
therefore many anhydrous salts which are capable of combining with
water absorb aqueous vapour from moist air; that is, they act like a
cold body on which water is deposited from steam. It is on this that
the desiccation of gases is based, and it must further be remarked in
this respect that certain substances--for instance, potassium carbonate
(K_{2}CO_{3}) and calcium chloride (CaCl_{2})--not only absorb the water
necessary for the formation of a solid crystalline compound, but also
give solutions, or _deliquesce_, as it is termed, in moist air. Many
crystals do not effloresce in the least at the ordinary temperature;
for example, copper sulphate, which may be preserved for an indefinite
length of time without efflorescing, but when placed under the receiver
of an air pump, if efflorescence be once started, it goes on at the
ordinary temperature. The temperature at which the complete separation
of water from crystals takes place varies considerably, not only for
different substances, but also for different portions of the contained
water. Very often the temperature at which dissociation begins is very
much higher than the boiling point of water. So, for example, copper
sulphate, which contains 36 p.c. of water, gives up 28·8 p.c. at 100°,
and the remaining quantity, namely 7·2 p.c., only at 240°. Alum, out of
the 45·5 p.c. of water which it contains, gives up 18·9 p.c. at 100°,
17·7 p.c. at 120°, 7·7 p.c. at 180°, and 1 p.c. at 280°; it only loses
the last quantity (1 p.c.) at its temperature of decomposition. These
examples clearly show that the annexation of water of crystallisation
is accompanied by a rather profound, although, in comparison with
instances which we shall consider later, still inconsiderable, change
of its properties. In certain cases the water of crystallisation is
only given off when the solid form of the substance is destroyed: when
the crystals melt on heating. The crystals are then said _to melt in
their water of crystallisation_. Further, after the separation of the
water, a solid substance remains behind, so that by further heating it
acquires a solid form. This is seen most clearly in crystals of sugar of
lead or lead acetate, which melt in their water of crystallisation at a
temperature of 56·25°, and in so doing begin to lose water. On reaching
a temperature of 100° the sugar of lead solidifies, having lost all its
water; and then at a temperature of 280°, the anhydrous and solidified
salt again melts.[65 bis]

  [63] Supersaturated solutions give an excellent proof in this respect.
       Thus a solution of copper sulphate generally crystallises in
       penta-hydrated crystals, CuSO_{4} + 5H_{2}O, and its saturated
       solution gives such crystals if it be brought into contact with
       the minutest possible crystal of the same kind. But, according
       to the observations of Lecoq de Boisbaudran, if a crystal of
       ferrous sulphate (an isomorphous salt, _see_ note 55), FeSO_{4}
       + 7H_{2}O, be placed in a saturated solution of copper sulphate,
       then crystals of hepta-hydrated salt, CuSO_{4} + 7H_{2}O,
       are obtained. It is evident that neither the penta-nor the
       hepta-hydrated salt is contained as such in the solution. The
       solution presents its own particular liquid form of equilibrium.

  [64] Efflorescence, like every evaporation, proceeds from the surface.
       In the interior of crystals which have effloresced there is
       usually found a non-effloresced mass, so that the majority of
       effloresced crystals of washing soda show, in their fracture, a
       transparent nucleus coated by an effloresced, opaque, powdery
       mass. It is a remarkable circumstance in this respect that
       efflorescence proceeds in a completely regular and uniform
       manner, so that the angles and planes of similar crystallographic
       character effloresce simultaneously, and in this respect the
       crystalline form determines those parts of crystals where
       efflorescence starts, and the order in which it continues. In
       solutions evaporation also proceeds from the surface, and the
       first crystals which appear on its reaching the required degree
       of saturation are also formed at the surface. After falling to
       the bottom the crystals naturally continue to grow (_see_ Chapter
       X.).

  [65] According to Lesc[oe]ur (1883), at 100° a concentrated solution of
       barium hydroxide, BaH_{2}O_{2}, on first depositing crystals
       (with + H_{2}O) has a tension of about 630 mm. (instead of
       760 mm., the tension of water), which decreases (because the
       solution evaporates) to 45 mm., when all the water is expelled
       from the crystals, BaH_{2}O_{2} + H_{2}O, which are formed, but
       they also lose water (dissociate, effloresce at 100°), leaving
       the hydroxide, BaH_{2}O_{2}, which is perfectly undecomposable
       at 100°--that is, does not part with water. At 73° (the tension
       of water is then 265 mm.) a solution, containing 33H_{2}O,
       on crystallising has a tension of 230 mm.; the crystals,
       BaH_{2}O_{2} + 8H_{2}O, which separate out, have a tension of
       160 mm.; on losing water they give BaH_{2}O_{2} + H_{2}O. This
       substance does not decompose at 73°, and therefore its tension =
       0. In those crystallohydrates which effloresce at the ordinary
       temperature, the tension of dissociation nearly approximates
       to that of the aqueous vapour, as Lesc[oe]ur (1891) showed. To
       this category of compounds belong B_{2}O_{3}(3 + _x_)H_{2}O,
       C_{2}O_{4}H_{2}(2 + _x_)H_{2}O, BaO(9 + _x_)H_{2}O, and SrO(9
       + _x_)H_{2}O. And a still greater tension is possessed by
       Na_{2}SO_{4}10H_{2}O, Na_{2}CO_{3}10H_{2}O, and MgSO_{4}(7 +
       _x_)H_{2}O. Müller-Erzbach (1884) determines the tension (with
       reference to liquid water) by placing tubes of the same length
       with water and the substances experimented with in a desiccator,
       the rate of loss of water giving the relative tension. Thus,
       at the ordinary temperature, crystals of sodium phosphate,
       Na_{2}HPO_{4} + 12H_{2}O, present a tension of 0·7 compared
       with water, until they lose 5H_{2}O, then 0·4 until they lose
       5H_{2}O more, and on losing the last equivalent of water the
       tension falls to 0·04 compared with water. It is clear that
       the different molecules of water are held by an unequal force.
       Out of the five molecules of water in copper sulphate the two
       first are comparatively easily separated even at the ordinary
       temperature (but only after several days in a desiccator,
       according to Latchinoff); the next two are more difficultly
       separated, and the last equivalent is retained even at 100°. This
       is another indication of the capacity of CuSO_{4} to form three
       hydrates, CuSO_{4}5H_{2}O, CuSO_{4}3H_{2}O, and CuSO_{4}H_{2}O.
       The researches of Andreae on the tension of dissociation of
       hydrated sulphate of copper showed (1891) the existence of three
       provinces, characterised at a given temperature by a constant
       tension: (1) between 3-5, (2) between 1-3, and lastly (3) between
       0-1 molecule of water, which again confirms the existence of
       three hydrates of the above composition for this salt.

  [65 bis] Sodium acetate (C_{2}H_{3}O_{2}Na,3H_{2}O) melts at 58°, but
       re-solidifies only on contact with a crystal, otherwise it
       may remain liquid even at 0°, and may be used for obtaining a
       constant temperature. According to Jeannel, the latent heat
       of fusion is about 28 calories, and according to Pickering
       the heat of solution 35 calories. When melted this salt boils
       at 123°--that is, the tension of the vapour given off at that
       temperature equals the atmospheric pressure.

It is most important to recognise in respect to the water of
crystallisation that its ratio to the quantity of the substance with
which it is combined is always a constant quantity. However often we
may prepare copper sulphate, we shall always find 36·14 p.c. of water
in its crystals, and these crystals always lose four-fifths of their
water at 100°, and one-fifth of the whole amount of the water contained
remains in the crystals at 100°, and is only expelled from them at a
temperature of about 240°. What has been said about crystals of copper
sulphate refers also to crystals of every other substance, which contain
water of crystallisation. It is impossible in any of these cases to
increase either the relative proportion of the salt or of the water,
without changing the homogeneity of the substance. If once a portion of
the water be lost--for instance, if once efflorescence takes place--a
mixture is obtained, and not a homogeneous substance, namely a mixture
of a substance deprived of water with a substance which has not yet lost
water--_i.e._ decomposition has already commenced. This constant ratio
is an example of the fact that in chemical compounds the quantity of
the component parts is quite definite; that is, it is an example of the
so-called _definite chemical compounds_. They may be distinguished from
solutions, and from all other so-called indefinite chemical compounds,
in that at least one, and sometimes both, of the component parts may be
added in a large quantity to an indefinite chemical compound, without
destroying its homogeneity, as in solutions, whilst it is impossible
to add any one of the component parts to a definite chemical compound
without destroying the homogeneity of the entire mass. Definite
chemical compounds only decompose at a certain rise in temperature;
on a lowering in temperature they do not, at least with very few
exceptions, yield their components like solutions which form ice or
compounds with water of crystallisation. This leads to the assumption
that solutions contain water as water,[66] although it may sometimes
be in a very small quantity. Therefore solutions which are capable
of solidifying completely (for instance, crystallo-hydrates capable
of melting) such as the compound of 84-1/2 parts of sulphuric acid,
H_{2}SO_{4}, with 15-1/2 parts of water, H_{2}O, or H_{2}SO_{4},H_{2}O
(or H_{4}SO_{5}), appear as true definite chemical compounds. If,
then, we imagine such a definite compound in a liquid state, and admit
that it partially decomposes in this state, separating water--not as
ice or vapour (for then the system would be heterogeneous, including
substances in different physical states), but in a liquid form, when
the system will be homogeneous--we shall form an idea of a solution as
an unstable, dissociating fluid state of equilibrium between water and
the substance dissolved. Moreover, it should be remarked that, judging
by experiment, many substances give with water not one but _diverse_
compounds,[67] which is seen in the capacity of one substance to form
with water many various _crystallo-hydrates_, or compounds with water
of crystallisation, showing diverse and independent properties. From
these considerations, _solutions[68] may be regarded as fluid, unstable,
definite chemical compounds in a state of dissociation_.[69]

  [66] Such a phenomenon frequently presents itself in purely chemical
       action. For instance, let a liquid substance _A_ give, with
       another liquid substance _B_, under the conditions of an
       experiment, a mere minute quantity of a solid or gaseous
       substance _C_. This small quantity will separate out (pass
       away from the sphere of action, as Berthollet expressed it),
       and the remaining masses of _A_ and _B_ will again give _C_;
       consequently, under these conditions action will go on to the
       end. Such, it seems to me, is the action in solutions when they
       yield ice or vapour indicating the presence of water.

  [67] Certain substances are capable of forming together only one
       compound, others several, and these of the most varied degrees
       of stability. The compounds of water are instances of this
       kind. In solutions the existence of several different definite
       compounds must be acknowledged, but many of these have not
       yet been obtained in a free state, and it may be that they
       cannot be obtained in any other but a liquid form--that is,
       dissolved; just as there are many undoubted definite compounds
       which only exist in one physical state. Among the hydrates such
       instances occur. The compound CO_{2} + 8H_{2}O (_see_ note
       31), according to Wroblewski, only occurs in a solid form.
       Hydrates like H_{2}S + 12H_{2}O (De Forcrand and Villard), HBr
       + H_{2}O (Roozeboom), can only be accepted on the basis of a
       decrease of tension, but present themselves as very transient
       substances, incapable of existing in a stable free state.
       Even sulphuric acid, H_{2}SO_{4}, itself, which undoubtedly
       is a definite compound, fumes in a liquid form, giving off
       the anhydride, SO_{3}--that is, it exhibits a very unstable
       equilibrium. The crystallo-hydrates of chlorine, Cl_{2} +
       8H_{2}O, of hydrogen sulphide, H_{2}S + 12H_{2}O (it is formed
       at 0°, and is completely decomposed at +1°, as then 1 vol. of
       water only dissolves 4 vols. of hydrogen sulphide, while at 0·1°
       it dissolves about 100 vols.), and of many other gases, are
       instances of hydrates which are very unstable.

  [68] Of such a kind are also other indefinite chemical compounds; for
       example, metallic alloys. These are solid substances or
       solidified solutions of metals. They also contain definite
       compounds, and may contain an excess of one of the metals.
       According to the experiments of Laurie (1888), the alloys of zinc
       with copper in respect to the electro-motive force in galvanic
       batteries behave just like zinc if the proportion of copper in
       the alloy does not exceed a certain percentage--that is, until
       a definite compound is attained--for in that case particles of
       free zinc are present; but if a copper surface be taken, and it
       be covered by only one-thousandth part of its area of zinc, then
       only the zinc will act in a galvanic battery.

  [69] According to the above supposition, the condition of solutions
       in the sense of the kinetic hypothesis of matter (that is, on
       the supposition of an internal motion of molecules and atoms)
       may be represented in the following form:--In a homogeneous
       liquid--for instance, water--the molecules occur in a certain
       state of, although mobile, still stable, equilibrium. When a
       substance _A_ dissolves in water, its molecules form with several
       molecules of water, systems _An_H_{2}O, which are so unstable
       that when surrounded by molecules of water they decompose and
       re-form, so that _A_ passes from one mass of molecules of water
       to another, and the molecules of water which were at this
       moment in harmonious motion with _A_ in the form of the system
       _An_H_{2}O, in the next instant may have already succeeded in
       getting free. The addition of water or of molecules of _A_
       may either only alter the number of free molecules, which in
       their turn enter into systems _An_H_{2}O, or they may introduce
       conditions for the possibility of building up new systems
       _Am_H_{2}O, where _m_ is either greater or less than _n_. If
       in the solution the relation of the molecules be the same as
       in the system _Am_H_{2}O, then the addition of fresh molecules
       of water or of _A_ would be followed by the formation of new
       molecules _An_H_{2}O. The relative quantity, stability, and
       composition of these systems or definite compounds will vary in
       one or another solution. I adopted this view of solutions (1887,
       Pickering subsequently put forward a similar view) after a most
       intimate study of the variation of their specific gravities, to
       which my book, cited in note 19, is devoted. Definite compounds,
       _An__{1}H_{2}O and _Am__{1}H_{2}O, existing in a free--for
       instance, solid--form, may in certain cases be held in solutions
       in a dissociated state (although but partially); they are similar
       in their structure to those definite substances which are formed
       in solutions, but it is not necessary to assume that such
       systems as Na_{2}SO_{4} + 10H_{2}O, or Na_{2}SO_{4} + 7H_{2}O,
       or Na_{2}SO_{4}, are contained in solutions. The comparatively
       more stable systems _An__{1}H_{2}O which exist in a free state
       and change their physical state must present, although within
       certain limits of temperature, an entirely harmonious kind of
       motion of _A_ with _n__{1}H_{2}O; the property also and state of
       systems _An_H_{2}O and _Am_H_{2}O, occurring in solutions, is
       that they are in a liquid form, although partially dissociated.
       Substances _A__{1}, which give solutions, are distinguished by
       the fact that they can form such unstable systems _An_H_{2}O,
       but besides them they can give other much more stable systems
       _An__{1}H_{2}O. Thus ethylene, C_{2}H_{4}, in dissolving in
       water, probably forms a system C_{2}H_{4}_n_H_{2}O, which easily
       splits up into C_{2}H_{4} and H_{2}O, but it also gives the
       system of alcohol, C_{2}H_{4},H_{2}O or C_{2}H_{6}O, which is
       comparatively stable. Thus oxygen can dissolve in water, and it
       can combine with it, forming peroxide of hydrogen. Turpentine,
       C_{10}H_{16}, does not dissolve in water, but it combines with it
       as a comparatively stable hydrate. In other words, the chemical
       structure of hydrates, or of the definite compounds which are
       contained in solutions, is distinguished not only by its original
       peculiarities but also by a diversity of stability. A similar
       structure to hydrates must be acknowledged in crystallo-hydrates.
       On melting they give actual (real) solutions. As substances which
       give crystallo-hydrates, like salts, are capable of forming a
       number of diverse hydrates, and as the greater the number of
       molecules of water (_n_) they (_An_H_{2}O) contain, the lower
       is the temperature of their formation, and as the more easily
       they decompose the more water they hold, therefore, in the first
       place, the isolation of hydrates holding much water existing in
       aqueous solutions may be soonest looked for at low temperatures
       (although, perhaps, in certain cases they cannot exist in the
       solid state); and, secondly, the stability also of such higher
       hydrates will be at a minimum under the ordinary circumstances
       of the occurrence of liquid water. Hence a further more detailed
       investigation of cryohydrates may help to the elucidation of
       the nature of solutions. But it may be foreseen that certain
       cryohydrates will, like metallic alloys, present solidified
       mixtures of ice with the salts themselves and their more stable
       hydrates, and others will be definite compounds.

In regarding solutions from this point of view they come under the head
of those definite compounds with which chemistry is mainly concerned.[70]

  [70] The above representation of solutions, &c., considering them as
       a particular state of definite compounds, excludes the
       independent existence of indefinite compounds; by this means that
       unity of chemical conception is obtained which cannot be arrived
       at by admitting the physico-mechanical conception of indefinite
       compounds. The gradual transition from typical solutions (as
       of gases in water, and of weak saline solutions) to sulphuric
       acid, and from it and its definite, but yet unstable and liquid,
       compounds, to clearly defined compounds, such as salts and their
       crystallo-hydrates, is so imperceptible, that in denying that
       solutions pertain to the number of definite but dissociating
       compounds, we risk denying the definiteness of the atomic
       composition of such substances as sulphuric acid or of molten
       crystallo-hydrates. I repeat, however, that for the present the
       theory of solutions cannot be considered as firmly established.
       The above opinion about them is nothing more than a hypothesis
       which endeavours to satisfy those comparatively limited data
       which we have for the present about solutions, and of those
       cases of their transition into definite compounds. By submitting
       solutions to the Daltonic conception of atomism, I hope that we
       may not only attain to a general harmonious chemical doctrine,
       but also that new motives for investigation and research will
       appear in the problem of solutions, which must either confirm the
       proposed theory or replace it by another fuller and truer one;
       and I for my part cannot consider this to be the case with any of
       the other present doctrines of solutions (note 49).

We saw above that copper sulphate loses four-fifths of its water at
100° and the remainder at 240°. This means that there are two definite
compounds of water with the anhydrous salt. Washing soda or carbonate of
sodium, Na_{2}CO_{3} separates out as crystals, Na_{2}CO_{3},10H_{2}O,
containing 62·9 p.c. of water by weight, from its solutions at the
ordinary temperature. When a solution of the same salt deposits crystals
at a low temperature, about -20°, then these crystals contain 71·8
parts of water per 28·2 parts of anhydrous salt. Further, the crystals
are obtained together with ice, and are left behind when it melts. If
ordinary soda, with 62·9 p.c. of water, be cautiously melted in its
own water of crystallisation, there remains a salt, in a solid state,
containing only 14·5 p.c. of water, and a liquid is obtained which
contains the solution of a salt which separates out crystals at 34°,
which contain 46 p.c. of water and do not effloresce in air. Lastly,
if a supersaturated solution of soda be prepared, then at temperatures
below 8° it deposits crystals containing 54·3 p.c. of water. Thus as
many as five compounds of anhydrous soda with water are known; and
they are dissimilar in their properties and crystalline form, and even
in their solubility. It is to be observed that the greatest amount of
water in the crystals corresponds with a temperature of -20°, and the
smallest to the highest temperature. There is apparently no relation
between the above quantities of water and the salts, but this is only
because in each case the amount of water and anhydrous salt was given
in percentages; but if it be calculated for one and the same quantity
of anhydrous salt, or of water, a great regularity will be observed in
the amounts of the component parts in all these compounds. It appears
that for 106 parts of anhydrous salt in the crystals separated out at
-20° there are 270 parts of water; in the crystals obtained at 15° there
are 180 parts of water; in the crystals obtained from a supersaturated
solution 126 parts, in the crystals which separate out at 34°, 90
parts, and the crystals with the smallest amount of water, 18 parts.
On comparing these quantities of water it may easily be seen that they
are in simple proportion to each other, for they are all divisible by
18, and are in the ratio 15 : 10 : 7 : 5 : 1. Naturally, direct
experiment, however carefully it be conducted, is hampered with errors,
but taking these unavoidable experimental errors into consideration,
it will be seen that for a given quantity of an anhydrous substance
there occur, in several of its compounds with water, quantities of
water which are in very simple multiple proportion. This is observed
in, and is common to, all definite chemical compounds. This rule is
called _the law of multiple proportions_. It was discovered by Dalton,
and will be evolved in further detail subsequently in this work. For
the present we will only state that the law of definite composition
enables the composition of substances to be expressed by formulæ,
and the law of multiple proportions permits the application of whole
numbers as coefficients of the symbols of the elements in these formulæ.
Thus the formula Na_{2}CO_{3},10H_{2}O shows directly that in this
crystallo-hydrate there are 180 parts of water to 106 parts by weight of
the anhydrous salt, because the formula of soda, Na_{2}CO_{3}, directly
answers to a weight of 106, and the formula of water to 18 parts, by
weight, which are here taken 10 times.

In the above examples of the combinations of water, we saw the gradually
increasing intensity of the bond between water and a substance with
which it forms a homogeneous compound. There is a series of such
compounds with water, in which the water is held with very great force,
and is only given up at a very high temperature, and sometimes cannot
be separated by any degree of heat without the entire decomposition
of the substance. In these compounds there is generally no outward
sign whatever of their containing water. A perfectly new substance is
formed from an anhydrous substance and water, in which sometimes the
properties of neither one nor the other substance are observable. In
the majority of cases, a considerable amount of heat is evolved in the
formation of such compounds with water. Sometimes the heat evolved is
so intense that a red heat is produced and light is emitted. It is
hardly to be wondered at, after this, that stable compounds are formed
by such a combination. Their decomposition requires great heat; a large
amount of work is necessary to separate them into their component
parts. All such compounds are definite, and, generally, completely and
clearly definite. The number of such definite compounds with water or
_hydrates_, in the narrow sense of the word, is generally inconsiderable
for each anhydrous substance; in the greater number of cases, there is
formed only one such combination of a substance with water, one hydrate,
having so great a stability. The water contained in these compounds is
often called _water of constitution_--_i.e._ water which enters into the
structure or composition of the given substance. By this it is desired
to express, that in other cases the molecules of water are, as it were,
separate from the molecules of that substance with which it is combined.
It is supposed that in the formation of hydrates this water, even in
the smallest particles, forms one complete whole with the anhydrous
substance. Many examples of the formation of such hydrates might be
cited. The most familiar example in practice is the hydrate of lime,
or so-called 'slaked' lime. Lime is prepared by burning limestone, by
which the carbonic anhydride is expelled from it, and there remains
a white stony mass, which is dense, compact, and rather tenacious.
Lime is usually sold in this form, and bears the name of 'quick' or
'unslaked' lime. If water be poured over such lime, a great rise in
temperature is remarked either directly, or after a certain time. The
whole mass becomes hot, part of the water is evaporated, the stony mass
in absorbing water crumbles into powder, and if the water be taken in
sufficient quantity and the lime be pure and well burnt, not a particle
of the original stony mass is left--it all crumbles into powder. If
the water be in excess, then naturally a portion of it remains and
forms a solution. This process is called 'slaking' lime. Slaked lime is
used in practice in intermixture with sand as mortar. Slaked lime is a
definite hydrate of lime. If it is dried at 100° it retains 24·3 p.c.
of water. This water can only be expelled at a temperature above 400°,
and then quicklime is re-obtained. The heat evolved in the combination
of lime with water is so intense that it can set fire to wood, sulphur,
gunpowder, &c. Even on mixing lime with ice the temperature rises to
100°. If lime be moistened with a small quantity of water in the dark,
a luminous effect is observed. But, nevertheless, water may still be
separated from this hydrate.[71] If phosphorus be burnt in dry air, a
white substance called 'phosphoric anhydride' is obtained. It combines
with water with such energy, that the experiment must be conducted with
great caution. A red heat is produced in the formation of the compound,
and it is impossible to separate the water from the resultant hydrate
at any temperature. The hydrate formed by phosphoric anhydride is a
substance which is totally undecomposable into its original component
parts by the action of heat. Almost as energetic a combination occurs
when sulphuric anhydride, SO_{3}, combines with water, forming its
hydrate, sulphuric acid, H_{2}SO_{4}. In both cases definite compounds
are produced, but the latter substance, as a liquid, and capable of
decomposition by heat, forms an evident link with solutions. If 80 parts
of sulphuric anhydride retain 18 parts of water, this water cannot be
separated from the anhydride, even at a temperature of 300°. It is only
by the addition of phosphoric anhydride, or by a series of chemical
transformations, that this water can be separated from its compound
with sulphuric anhydride. Oil of vitriol, or sulphuric acid, is such a
compound. If a larger proportion of water be taken, it will combine
with the H_{2}SO_{4}; for instance, if 36 parts of water per 80 parts of
sulphuric anhydride be taken, a compound is formed which crystallises in
the cold, and melts at +8°, whilst oil of vitriol does not solidify even
at -30°. If still more water be taken, the oil of vitriol will dissolve
in the remaining quantity of water. An evolution of heat takes place,
not only on the addition of the water of constitution, but in a less
degree on further additions of water.[72] And therefore there is no
distinct boundary, but only a gradual transition, between those chemical
phenomena which are expressed in the formation of solutions and those
which take place in the formation of the most stable hydrates.[73]

  [71] In combining with water one part by weight of lime evolves 245
       units of heat. A high temperature is obtained, because the
       specific heat of the resulting product is small. Sodium oxide,
       Na_{2}O, in reacting on water, H_{2}O, and forming caustic soda
       (sodium hydroxide), NaHO, evolves 552 units of heat for each part
       by weight of sodium oxide.

  [72] The diagram given in note 28 shows the evolution of heat on the
       mixture of sulphuric acid, or monohydrate (H_{2}SO_{4}, _i.e._
       SO_{3} + H_{2}O), with different quantities of water per 100
       vols. of the resultant solution. Every 98 grams of sulphuric
       acid (H_{2}SO_{4}) evolve, on the addition of 18 grams of water,
       6,379 units of heat; with twice or three times the quantity of
       water 9,418 and 11,137 units of heat, and with an infinitely
       large quantity of water 17,860 units of heat, according to the
       determinations of Thomsen. He also showed that when H_{2}SO_{4}
       is formed from SO_{3} (= 80) and H_{2}O (= 18), 21,308 units
       of heat are evolved per 98 parts by weight of the resultant
       sulphuric acid.

  [73] Thus, for different hydrates the stability with which they hold
       water is very dissimilar. Certain hydrates hold water very
       loosely, and in combining with it evolve little heat. From
       other hydrates the water cannot be separated by any degree of
       heat, even if they are formed from anhydrides (_i.e._ anhydrous
       substances) and water with little evolution of heat; for
       instance, acetic anhydride in combining with water evolves an
       inconsiderable amount of heat, but the water cannot then be
       expelled from it. If the hydrate (acetic acid) formed by this
       combination be strongly heated it either volatilises without
       change, or decomposes into new substances, but it does not again
       yield the original substances--_i.e._, the anhydride and water,
       at least in a liquid form. Here is an instance which gives the
       reason for calling the water entering into the composition of the
       hydrate, water of constitution. Such, for example, is the water
       entering into the so-called caustic soda or sodium hydroxide
       (_see_ note 71). But there are hydrates which easily part with
       their water; yet this water cannot be considered as water of
       crystallisation, not only because sometimes such hydrates have no
       crystalline form, but also because, in perfectly analogous cases,
       very stable hydrates are formed, which are capable of particular
       kinds of chemical reactions, as we shall subsequently learn.
       Such, for example, is the unstable hydrated oxide of copper,
       which is not formed from water and oxide of copper, but which is
       obtained just like far more stable hydrates, for example, the
       hydrated oxide of barium BaH_{2}O_{2} equal to BaO + H_{2}O, by
       the double decomposition of the solution of salts with alkalies.
       In a word, there is no distinct boundary either between the water
       of hydrates and of crystallisation, or between solution and
       hydration.

       It must be observed that in separating from an aqueous solution,
       many substances, without having a crystalline form, hold water in
       the same unstable state as in crystals; only this water cannot
       be termed 'water of crystallisation' if the substance which
       separates out has no crystalline form. The hydrates of alumina
       and silica are examples of such unstable hydrates. If these
       substances are separated from an aqueous solution by a chemical
       process, then they always contain water. The formation of a new
       chemical compound containing water is here particularly evident,
       for alumina and silica in an anhydrous state have chemical
       properties differing from those they show when combined with
       water, and do not combine directly with it. The entire series of
       colloids on separating from water form similar compounds with
       it, which have the aspect of solid gelatinous substances. Water
       is held in a considerable quantity in solidified glue or boiled
       albumin. It cannot be expelled from them by pressure; hence,
       in this case there has ensued some kind of combination of the
       substance with water. This water, however, is easily separated on
       drying; but not the whole of it, a portion being retained, and
       this portion is considered to belong to the hydrate, although
       in this case it is very difficult, if not impossible, to obtain
       definite compounds. The absence of any distinct boundary lines
       between solutions, crystallo-hydrates, and ordinary hydrates
       above referred to, is very clearly seen in such examples.

We have thus considered many aspects and degrees of combination of
various substances with water, or instances of the compounds of water,
when it and other substances form new homogeneous substances, which
in this case will evidently be complex--_i.e._ made up of different
substances--and although they are homogeneous, yet it must be admitted
that in them there exist those component parts which entered into their
composition, inasmuch as these parts may be re-obtained from them. It
must not be imagined that water really exists in hydrate of lime, any
more than that ice or steam exists in water. When we say that water
occurs in the composition of a certain hydrate, we only wish to point
out that there are chemical transformations in which it is possible to
obtain that hydrate by means of water, and other transformations in
which this water may be separated out from the hydrate. This is all
simply expressed by the words, that water enters into the composition of
this hydrate. If a hydrate be formed by feeble bonds, and be decomposed
even at the ordinary temperature, and be a liquid, then the water
appears as one of the products of dissociation, and this gives an idea
of what solutions are, and forms the fundamental distinction between
them and other hydrates in which the water is combined with greater
stability.




                               CHAPTER II

                   THE COMPOSITION OF WATER, HYDROGEN


The question now arises, Is not _water_ itself a _compound substance_?
Cannot it be formed by the mutual combination of some component parts?
Cannot it be broken up into its component parts? There cannot be the
least doubt that if it does split up, and if it is a compound, then it
is a _definite_ one characterised by the stability of the union between
those component parts from which it is formed. From the fact alone that
water passes into all physical states as a homogeneous whole, without
in the least varying chemically in its properties and without splitting
up into its component parts (neither solutions nor many hydrates can
be distilled--they are split up), we must conclude, from this fact
alone, that if water is a compound then it is a stable and definite
chemical compound capable of entering into many other combinations.
Like many other great discoveries in the province of chemistry, it is
to the end of the last century that we are indebted for the important
discovery that water is not a simple substance, that it is composed of
two substances like a number of other compound substances. This was
proved by two of the methods by which the compound nature of bodies
may be directly determined; by analysis and by synthesis--that is,
by a method of the decomposition of water into, and of the formation
of water from, its component parts. In 1781 Cavendish first obtained
water by burning hydrogen in oxygen, both of which gases were already
known to him. He concluded from this that water was composed of two
substances. But he did not make more accurate experiments, which would
have shown the relative quantities of the component parts in water,
and which would have determined its complex nature with certainty.
Although his experiments were the first, and although the conclusion
he drew from them was true, yet such novel ideas as the complex nature
of water are not easily recognised so long as there is no series of
researches which entirely and indubitably proves the truth of such a
conclusion. The fundamental experiments which proved the complexity
of water by the method of synthesis, and of its formation from other
substances, were made in 1789 by Monge, Lavoisier, Fourcroy, and
Vauquelin. They obtained four ounces of water by burning hydrogen,
and found that water consists of 15 parts of hydrogen and 85 parts of
oxygen. It was also proved that the weight of water formed was equal
to the sum of the weights of the component parts entering into its
composition; consequently, water contains all the matter entering into
oxygen and hydrogen. The complexity of water was proved in this manner
by a method of synthesis. But we will turn to its analysis--_i.e._ to
its decomposition into its component parts. The analysis may be more or
less complete. Either both component parts may be obtained in a separate
state, or else only one is separated and the other is converted into a
new compound in which its amount may be determined by weighing. This
will be a reaction of substitution, such as is often taken advantage of
for analysis. The first analysis of water was thus conducted in 1784
by Lavoisier and Meusnier. The apparatus they arranged consisted of a
glass retort containing water previously purified, and of which the
weight had been determined. The neck of the retort was inserted into
a porcelain tube, placed inside an oven, and heated to a red heat by
charcoal. Iron filings, which decompose water at a red heat, were placed
inside this tube. The end of the tube was connected with a worm, for
condensing any water which might pass through the tube undecomposed.
This condensed water was collected in a separate flask. The gas formed
by the decomposition was collected over water in a bell jar. The aqueous
vapour in passing over the red-hot iron was decomposed, and a gas was
formed from it whose weight could be determined from its volume, its
density being known. Besides the water which passed through the tube
unaltered, a certain quantity of water disappeared in the experiment,
and this quantity, in the experiments of Lavoisier and Meusnier, was
equal to the weight of gas which was collected in the bell jar plus the
increase in weight of the iron filings. Hence the water was decomposed
into a gas, which was collected in the bell jar, and a substance, which
combined with the iron; consequently, it is composed of these two
component parts. This was the first analysis of water ever made; but
here only one (and not both) of the gaseous component parts of water was
collected separately. Both the component parts of water can, however,
be simultaneously obtained in a free state. For this purpose the
decomposition is brought about by a galvanic current or by heat, as we
shall learn directly.[1]

  [1] The first experiments of the synthesis and decomposition of water
      did not afford, however, an entirely convincing proof that water
      was composed of hydrogen and oxygen only. Davy, who investigated
      the decomposition of water by the galvanic current, thought for a
      long time that, besides the gases, an acid and alkali were also
      obtained. He was only convinced of the fact that water contains
      nothing but hydrogen and oxygen by a long series of researches,
      which showed him that the appearance of an acid and alkali in the
      decomposition of water proceeds from the presence of impurities
      (especially from the presence of ammonium nitrate) in water. A
      final comprehension of the composition of water is obtained from
      the accurate determination of the quantities of the component
      parts which enter into its composition. It will be seen from
      this how many data are necessary for proving the composition of
      water--that is, of the transformations of which it is capable.
      What has been said of water refers to all other compounds; the
      investigation of each one, the entire proof of its composition,
      can only be obtained by the accumulation of a large mass of data
      referring to it.

Water is a bad conductor of electricity--that is, pure water does not
transmit a feeble current; but if any salt or acid be dissolved in it,
then its conductivity increases, and _on the passage of a current_
through acidified water _it is decomposed_ into its component parts.
Some sulphuric acid is generally added to the water. By immersing
platinum plates (electrodes) in this water (platinum is chosen because
it is not acted on by acids, whilst many other metals are chemically
acted on by acids), and connecting them with a galvanic battery, it
will be observed that bubbles of gas appear on these plates. The gas
which separates is called _detonating gas_,[2] because, on ignition,
it very easily explodes.[3] What takes place is as follows:--First,
the water, by the action of the current, is decomposed into two gases.
The mixture of these gases forms detonating gas. When detonating gas
is brought into contact with an incandescent substance--for instance,
a lighted taper--the gases re-combine, forming water, the combination
being accompanied by a great evolution of heat, and therefore the vapour
of the water formed expands considerably, which it does very rapidly,
and as a consequence, an explosion takes place--that is, sound and
increase of pressure, and atmospheric disturbance, as in the explosion
of gunpowder.

  [2] This gas is collected in a voltameter.

  [3] In order to observe this explosion without the slightest danger,
      it is best to proceed in the following manner. Some soapy water is
      prepared, so that it easily forms soap bubbles, and it is poured
      into an iron trough. In this water, the end of a gas-conducting
      tube is immersed. This tube is connected with any suitable
      apparatus, in which detonating gas is evolved. Soap bubbles,
      full of this gas, are then formed. If the apparatus in which the
      gas is produced be then removed (otherwise the explosion might
      travel into the interior of the apparatus), and a lighted taper be
      brought to the soap bubbles, a very sharp explosion takes place.
      The bubbles should be small to avoid any danger; ten, each about
      the size of a pea, suffice to give a sharp report, like a pistol
      shot.

In order to discover what gases are obtained by the decomposition of
water, the gases which separate at each electrode must be collected
separately. For this purpose a V-shaped tube is taken; one of its ends
is open and the other fused up. A platinum wire, terminating inside
the tube in a plate, is fused into the closed end; the closed end is
entirely filled with water[4] acidified with sulphuric acid, and another
platinum wire, terminating in a plate, is immersed in the open end. If
a current from a galvanic battery be now passed through the wires an
evolution of gases will be observed, and the gas which is obtained in
the open branch passes into the air, while that in the closed branch
accumulates above the water. As this gas accumulates it displaces the
water, which continues to descend in the closed and ascend into the open
branch of the tubes. When the water, in this way, reaches the top of the
open end, the passage of the current is stopped, and the gas which was
evolved from one of the electrodes only is obtained in the apparatus.
By this means it is easy to prove that a particular gas appears at each
electrode. If the closed end be connected with the negative pole--_i.e._
with that joined to the zinc--then the gas collected in the apparatus
is capable of burning. This may be demonstrated by the following
experiment:--The bent tube is taken off the stand, and its open end
stopped up with the thumb and inclined in such a manner that the gas
passes from the closed to the open end. It will then be found, on
applying a lighted lamp or taper, that the gas burns. This combustible
gas is _hydrogen_. If the same experiment be carried on with a current
passing in the opposite direction--that is, if the closed end be joined
up with the positive pole (_i.e._ with the carbon, copper, or platinum),
then the gas which is evolved from it does not itself burn, but it
supports combustion very vigorously, so that a smouldering taper in it
immediately bursts into flame. This gas, which is collected at the anode
or positive pole, is _oxygen_, which is obtained, as we saw before (in
the Introduction), from mercury oxide and is contained in air.

  [4] In order to fill the tube with water, it is turned up, so that
      the closed end points downwards and the open end upwards, and
      water acidified with sulphuric acid is poured into it.

Thus in the decomposition of water oxygen appears at the positive pole
and hydrogen at the negative pole,[4 bis] so that detonating gas will be
a mixture of both. Hydrogen burns in air from the fact that in doing so
it re-forms water, with the oxygen of the air. Detonating gas explodes
from the fact that the hydrogen burns in the oxygen mixed with it. It
is very easy to measure the relative quantities of one and the other
gas which are evolved in the decomposition of water. For this purpose
a funnel is taken, whose orifice is closed by a cork through which two
platinum wires pass. These wires are connected with a battery. Acidified
water is poured into the funnel, and a glass cylinder full of water
is placed over the end of each wire (fig. 18). On passing a current,
hydrogen and oxygen collect in these cylinders, and it will easily be
seen that two volumes of hydrogen are evolved for every one volume of
oxygen. This signifies that, in decomposing, water gives two volumes of
hydrogen and one volume of oxygen.

  [4 bis] Owing to the gradual but steady progress made during the
      last twenty-five years in the production of an electric current
      from the dynamo and its transmission over considerable distances,
      the electrolytic decomposition of many compound bodies has
      acquired great importance, and the use of the electric current is
      making its way into many chemical manufactures. Hence, Prof. D.
      A. Lachinoff's proposal to obtain hydrogen and oxygen (both of
      which have many applications) by means of electrolysis (either of
      a 10 to 15 per cent. solution of caustic soda or a 15 per cent.
      solution of sulphuric acid) may find a practical application, at
      all events in the future. In general, owing to their simplicity,
      electrolytic methods have a great future, but as yet, so long
      as the production of an electric current remains so costly,
      their application is limited. And for this reason, although
      certain of these methods are mentioned in this work, they are
      not specially considered, the more so since a profitable and
      proper use of the electric current for chemical purposes requires
      special electro-technical knowledge which beginners cannot he
      assumed to have, and therefore, an exposition of the principles
      of electrotechnology as applied, to the production of chemical
      transformations, although referred to in places, does not come
      within the scope of the present work.

[Illustration: FIG. 18.--Decomposition of water by the galvanic current,
for determining the relation between the volumes of hydrogen and oxygen.]

Water is also decomposed into its component parts by _the action of
heat_. At the melting point of silver (960°), and in its presence,
water is decomposed and the oxygen absorbed by the molten silver,
which dissolves it so long as it is liquid. But directly the silver
solidifies the oxygen is expelled from it. However, this experiment
is not entirely convincing; it might be thought that in this case the
decomposition of the water did not proceed from the action of heat, but
from the action of the silver on water--that silver decomposes water,
taking up the oxygen. If steam be passed through a red-hot tube, whose
internal temperature attains 1,000°, then a portion[5] of the water
decomposes into its component parts, forming detonating gas. But on
passing into the cooler portions of the apparatus this detonating gas
again reunites and forms water. The hydrogen and oxygen obtained combine
together at a lower temperature.[6] Apparently the problem--to show the
decomposability of water at high temperatures--is unattainable. It was
considered as such before Henri Sainte-Claire Deville (in the fifties)
introduced the conception of dissociation into chemistry, as of a change
of chemical state resembling evaporation, if decomposition be likened to
boiling, and before he had demonstrated the decomposability of water by
the action of heat in an experiment which will presently be described.
In order to demonstrate clearly the _dissociation_ of water, or its
decomposability by heat, at a temperature approaching that at which it
is formed, it was necessary to separate the hydrogen from the oxygen at
a high temperature, without allowing the mixture to cool. Deville took
advantage of the difference between the densities of hydrogen and oxygen.

  [5] As water is formed by the combination of oxygen and hydrogen, with
      a considerable evolution of heat, and as it can also be
      decomposed, this reaction is a reversible one (_see_
      Introduction), and consequently at a high temperature the
      decomposition of water cannot be complete--it is limited by the
      opposite reaction. Strictly speaking, it is not known how much
      water is decomposed at a given temperature, although many efforts
      (Bunsen, and others) have been made in various directions to solve
      this question. Not knowing the coefficient of expansion, and the
      specific heat of gases at such high temperatures, renders all
      calculations (from observations of the pressure on explosion)
      doubtful.

  [6] Grove, in 1847, observed that a platinum wire fused in the
      oxyhydrogen flame--that is, having acquired the temperature
      of the formation of water--and having formed a molten drop at
      its end which fell into water, evolved detonating gas--that
      is, decomposed water. It therefore follows that water already
      decomposes at the temperature of its formation. At that time, this
      formed a scientific paradox; this we shall unravel only with the
      development of the conceptions of dissociation, introduced into
      science by Henri Sainte-Claire Deville, in 1857. These conceptions
      form an important epoch in science, and their development is one
      of the problems of modern chemistry. The essence of the matter is
      that, at high temperatures, water exists but also decomposes, just
      as a volatile liquid, at a certain temperature, exists both as a
      liquid and as a vapour. Similarly as a volatile liquid saturates
      a space, attaining its maximum tension, so also the products of
      dissociation have their maximum tension, and once that is attained
      decomposition ceases, just as evaporation ceases. Under like
      conditions, if the vapour be allowed to escape (and therefore its
      partial pressure be diminished), evaporation recommences, so also
      if the products of decomposition be removed, decomposition again
      continues. These simple conceptions of dissociation introduce
      infinitely varied consequences into the mechanism of chemical
      reactions, and therefore we shall have occasion to return to
      them very often. We may add that Grove also concluded that water
      was decomposed at a white heat, from the fact that he obtained
      detonating gas by passing steam through a tube with a wire heated
      strongly by an electric current, and also by passing steam over
      molten oxide of lead, he obtained, on the one hand, litharge (=
      oxide of lead and oxygen), and on the other, metallic lead formed
      by the action of hydrogen.

[Illustration: FIG. 19.--Decomposition of water by the action of heat,
and the separation of the hydrogen formed by its permeating through a
porous tube.]

A wide porcelain tube P (fig. 19) is placed in a furnace, which can be
raised to a high temperature (it should be heated with small pieces of
good coke). In this tube there is inserted a second tube T, of smaller
diameter, made of unglazed earthenware and therefore porous. The ends
of the tube are luted to the wide tube, and two tubes, C and C', are
inserted into the ends, as shown in the drawing. With this arrangement
it is possible for a gas to pass into the annular space between the
walls of the two tubes, from whence it can be collected. Steam from a
retort or flask is passed through the tube D, into the inner porous
tube T. This steam on entering the red-hot space is decomposed into
hydrogen and oxygen. The densities of these gases are very different,
hydrogen being sixteen times lighter than oxygen. Light gases, as we
saw above, penetrate through porous surfaces very much more rapidly
than denser gases, and therefore the hydrogen passes through the
pores of the tube into the annular space very much more rapidly than
the oxygen. The hydrogen which separates out into the annular space
can only be collected when this space does not contain any oxygen.
If any air remains in this space, then the hydrogen which separates
out will combine with its oxygen and form water. For this reason a
gas incapable of supporting combustion--for instance, nitrogen or
carbonic anhydride--is previously passed into the annular space. Thus
the carbonic anhydride is passed through the tube C, and the hydrogen,
separated from the steam, is collected through the tube C', and will be
partly mixed with carbonic anhydride. A certain portion of the carbonic
anhydride will penetrate through the pores of the unglazed tube into
the interior of the tube T. The oxygen will remain in this tube, and
the volume of the remaining oxygen will be half that of the volume of
hydrogen which separates out from the annular space.[6 bis]

  [6 bis] Part of the oxygen will also penetrate through the pores
      of the tube; but, as was said before, a much smaller quantity
      than the hydrogen, and as the density of oxygen is sixteen times
      greater than that of hydrogen, the volume of oxygen which passes
      through the porous walls will be four times less than the volume
      of hydrogen (the quantities of gases passing through porous
      walls are inversely proportional to the square roots of their
      densities). The oxygen which separates out into the annular
      space will combine, at a certain fall of temperature, with the
      hydrogen; but as each volume of oxygen only requires two volumes
      of hydrogen, whilst at least four volumes of hydrogen will pass
      through the porous walls for every volume of oxygen that passes,
      therefore, part of the hydrogen will remain free, and can be
      collected from the annular space. A corresponding quantity of
      oxygen remaining from the decomposition of the water can be
      collected from the internal tube.

The decomposition of water is effected much more easily by a method of
substitution, taking advantage of the affinity of substances for the
oxygen or the hydrogen of water. If a substance be added to water,
which takes up the oxygen and replaces the hydrogen--then we shall
obtain the latter gas from the water. Thus with sodium, water gives
hydrogen, and with chlorine, which takes up the hydrogen, oxygen is
obtained.

Hydrogen is evolved from water by many metals, which are capable
of forming oxides in air--that is, which are capable of burning or
combining with oxygen. The capacity of metals for combining with
oxygen, and therefore for decomposing water, or for the evolution of
hydrogen, is very dissimilar.[7] Among metals, potassium and sodium
exhibit considerable energy in this respect. The first occurs in
potash, the second in soda. They are both lighter than water, soft, and
easily change in air. By bringing one or the other of them in contact
with water at the ordinary temperature,[8] a quantity of hydrogen,
corresponding with the amount of the metal taken, may be directly
obtained. One gram of hydrogen, occupying a volume of 11·16 litres at 0°
and 760 mm., is evolved from every 39 grams of potassium, or 23 grams of
sodium. The phenomenon may be observed in the following way: a solution
of sodium in mercury--or 'sodium amalgam,' as it is generally called--is
poured into a vessel containing water, and owing to its weight sinks to
the bottom; the sodium held in the mercury then acts on the water like
pure sodium, liberating hydrogen. The mercury does not act here, and the
same amount of it as was taken for dissolving the sodium is obtained in
the residue. The hydrogen is evolved gradually in the form of bubbles,
which pass through the liquid.

  [7] In order to demonstrate the difference of the affinity of oxygen
      for different elements, it is enough to compare the amounts of
      heats which are evolved in their combination with 16 parts by
      weight of oxygen; in the case of sodium (when Na_{2}O is formed,
      or 46 parts of Na combine with 16 parts of oxygen, according to
      Beketoff) 100,000 calories (or units of heat), are evolved, for
      hydrogen (when water, H_{2}O, is formed) 69,000 calories, for
      iron (when the oxide FeO is formed) 69,000, and if the oxide
      Fe_{2}O_{3} is formed, 64,000 calories, for zinc (ZnO is formed)
      86,000 calories, for lead (when PbO is formed) 51,000 calories,
      for copper (when CuO is formed) 38,000 calories, and for mercury
      (HgO is formed) 31,000 calories.

      These figures cannot correspond directly with the magnitude of the
      affinities, for the physical and mechanical side of the matter is
      very different in the different cases. Hydrogen is a gas, and, in
      combining with oxygen, gives a liquid; consequently it changes
      its physical state, and, in doing so, evolves heat. But zinc and
      copper are solids, and, in combining with oxygen, give solid
      oxides. The oxygen, previously a gas, now passes into a solid or
      liquid state, and, therefore, also must have given up its store of
      heat in forming oxides. As we shall afterwards see, the degree of
      contraction (and consequently of mechanical work) was different
      in the different cases, and therefore the figures expressing the
      heat of combination cannot directly depend on the affinities,
      on the loss of internal energy previously in the elements.
      Nevertheless, the figures above cited correspond, in a certain
      degree, with the order in which the elements stand in respect
      to their affinity for oxygen, as may be seen from the fact that
      the mercury oxide, which evolves the least heat (among the above
      examples), is the least stable is easily decomposed, giving up its
      oxygen; whilst sodium, the formation of whose oxide is accompanied
      by the greatest evolution of heat, is able to decompose all the
      other oxides, taking up their oxygen. In order to generalise the
      connection between affinity and the evolution and the absorption
      of heat, which is evident in its general features, and was firmly
      established by the researches of Favre and Silbermann (about
      1840), and then of Thomsen (in Denmark) and Berthelot (in France),
      many investigators, especially the one last mentioned, established
      the _law of maximum work_. This states that only those chemical
      reactions take place of their own accord in which the greatest
      amount of chemical (latent, potential) energy is transformed into
      heat. But, in the first place, we are not able, judging from what
      has been said above, to distinguish that heat which corresponds
      with purely chemical action from the sum total of the heat
      observed in a reaction (in the calorimeter); in the second place,
      there are evidently endothermal reactions which proceed under
      the same circumstances as exothermal (carbon burns in the vapour
      of sulphur with absorption of heat, whilst in oxygen it evolves
      heat); and, in the third place, there are reversible reactions,
      which when taking place in one direction evolve heat, and when
      taking place in the opposite direction absorb it; and, therefore,
      the principle of maximum work in its elementary form is not
      supported by science. But the subject continues to be developed,
      and will probably lead to a general law, such as thermal chemistry
      does not at present possess.

  [8] If a piece of metallic sodium be thrown into water, it floats on it
      (owing to its lightness), keeps in a state of continual motion
      (owing to the evolution of hydrogen on all sides), and immediately
      decomposes the water, evolving hydrogen, which can be lighted.
      This experiment may, however, lead to an explosion should the
      sodium stick to the walls of the vessel, and begin to act on the
      limited mass of water immediately adjacent to it (probably in
      this case NaHO forms with Na, Na_{2}O, which acts on the water,
      evolving much heat and rapidly forming steam), and the experiment
      should therefore be carried on with caution. The decomposition
      of water by sodium may he better demonstrated, and with greater
      safety, in the following manner. Into a glass cylinder filled
      with mercury, and immersed in a mercury bath, water is first
      introduced, which will, owing to its lightness, rise to the top,
      and then a piece of sodium wrapped in paper is introduced with
      forceps into the cylinder. The metal rises through the mercury
      to the surface of the water, on which it remains, and evolves
      hydrogen, which collects in the cylinder, and may be tested after
      the experiment has been completed. The safest method of making
      this experiment is, however, as follows. The sodium (cleaned from
      the naphtha in which it is kept) is either wrapped in fine copper
      gauze and held by forceps, or else held in forceps at the end of
      which a small copper cage is attached, and is then held under
      water. The evolution of hydrogen goes on quietly, and it may he
      collected in a bell jar and then lighted.

Beyond the hydrogen evolved and a solid substance, which remains in
solution (it may be obtained by evaporating the resultant solution) no
other products are here obtained. Consequently, from the two substances
(water and sodium) taken, the same number of new substances (hydrogen
and the substance dissolved in water) have been obtained, from which we
may conclude that the reaction which here takes place is a reaction of
double decomposition or of substitution. The resultant solid is nothing
else but the so-called caustic soda (sodium hydroxide), which is made
up of sodium, oxygen, and half of the hydrogen contained in the water.
Therefore, the substitution took place between the hydrogen and the
sodium, namely half of the hydrogen in the water was replaced by the
sodium, and was evolved in a free state. Hence the reaction which takes
place here may be expressed by the equation H_{2}O + Na = NaHO + H; the
meaning of this is clear from what has already been said.[9]

  [9] This reaction is vigorously exothermal, _i.e._ it is accompanied
      by the evolution of heat. If a sufficient quantity of water
      be taken the whole of the sodium hydroxide, NaHO, formed is
      dissolved, and about 42,500 units of heat are evolved per 23 grams
      of sodium taken. As 40 grams of sodium hydroxide are produced,
      and they in dissolving, judging from direct experiment, evolve
      about 10,000 calories; therefore, without an excess of water, and
      without the formation of a solution, the reaction would evolve
      about 32,500 calories. We shall afterwards learn that hydrogen
      contains in its smallest isolable particles H_{2} and not H,
      and therefore it follows that the reaction should be written
      thus--2Na + 2H_{2}O = H_{2} + 2NaOH, and it then corresponds with
      an evolution of heat of +65,000 calories. And as N. N. Beketoff
      showed that Na_{2}O, or anhydrous oxide of sodium, forms the
      hydrate, or sodium hydroxide (caustic soda), 2NaHO, with water,
      evolving about 35,500 calories, therefore the reaction 2Na +
      H_{2}O = H_{2} + Na_{2}O corresponds to 29,500 calories. This
      quantity of heat is less than that which is evolved in combining
      with water, in the formation of caustic soda, and therefore it is
      not to be wondered at that the hydrate, NaHO, is always formed
      and not the anhydrous substance Na_{2}O. That such a conclusion,
      which agrees with facts, is inevitable is also seen from the fact
      that, according to Beketoff, the anhydrous sodium oxide, Na_{2}O,
      acts directly on hydrogen, with separation of sodium, Na_{2}O + H
      = NaHO + Na. This reaction is accompanied by an evolution of heat
      equal to about 3,000 calories, because Na_{2}O + H_{2}O gives, as
      we saw, 35,500 calories and Na + H_{2}O evolves 32,500 calories.
      However, an opposite reaction also takes place--NaHO + Na =
      Na_{2}O + H (both with the aid of heat)--consequently, in this
      case heat is absorbed. In this we see an example of calorimetric
      calculations and the limited application of the law of maximum
      work for the general phenomena of reversible reactions, to which
      the case just considered belongs. But it must be remarked that
      all reversible reactions evolve or absorb but little heat, and
      the reason of the law of maximum work, not being universal must
      first of all be looked for in the fact that we have no means of
      separating the heat which corresponds with the purely chemical
      process from the sum total of the heat observed, and as the
      structure of a number of substances is altered by heat and also
      by contact, we can scarcely hope that the time approaches when
      such a distinction will be possible. A heated substance, in point
      of fact, has no longer the original energy of its atoms--that is,
      the act of heating not only alters the store of motion of the
      molecules but also of the atoms forming the molecules, in other
      words, it makes the beginning of or preparation for chemical
      change. From this it must be concluded that thermochemistry, or
      the study of the heat accompanying chemical transformations,
      cannot he identified with chemical mechanics. Thermo-chemical data
      form a part of it, but they alone cannot give it.

Sodium and potassium act on water at the ordinary temperature. Other
heavier metals only act on it with a rise of temperature, and then not
so rapidly or vigorously. Thus magnesium and calcium only liberate
hydrogen from water at its boiling point, and zinc and iron only a red
heat, whilst a whole series of heavy metals, such as copper, lead,
mercury, silver, gold, and platinum, do not in the least decompose water
at any temperature, and do not replace its hydrogen.

From this it is clear that hydrogen may be obtained by the decomposition
of steam by the action of iron (or zinc) with a rise of temperature.
The experiment is conducted in the following manner: pieces of iron
(filings, nails, &c.), are placed in a porcelain tube, which is then
subjected to a strong heat and steam passed through it. The steam,
coming into contact with the iron, gives up its oxygen to it, and thus
the hydrogen is set free and passes out at the other end of the tube
together with undecomposed steam. This method, which is historically
very significant,[10] is practically inconvenient, as it requires a
rather high temperature. Further, this reaction, as a reversible one
(a red-hot mass of iron decomposes a current of steam, forming oxide
and hydrogen; and a mass of oxide of iron, heated to redness in a
stream of hydrogen, forms iron and steam), does not proceed in virtue
of the comparatively small difference between the affinity of oxygen
for iron (or zinc) and for hydrogen, but only because the hydrogen
escapes, as it is formed, in virtue of its elasticity.[11] If the
oxygen compounds--that is, the oxides--which are obtained from the
iron or zinc, be able to pass into solution, then the affinity acting
in solution is added, and the reaction may become non-reversible, and
proceed with comparatively much greater facility.[12] As the oxides
of iron and zinc, by themselves insoluble in water, are capable
of combining with (have an affinity for) acid oxides (as we shall
afterwards fully consider), and form saline and soluble substances, with
acids, or hydrates having acid properties, hence by the action of such
hydrates, or of their aqueous solutions,[13] iron and zinc are able to
liberate hydrogen with great ease at the ordinary temperature--that
is, they act on solutions of acids just as sodium acts on water.[14]
Sulphuric acid, H_{2}SO_{4}, is usually chosen for this purpose; the
hydrogen is displaced from it by many metals with much greater facility
than directly from water, and such a displacement is accompanied by
the evolution of a large amount of heat.[15] When the hydrogen in
sulphuric acid is replaced by a metal, a substance is obtained which is
called a salt of sulphuric acid or a sulphate. Thus, by the action of
zinc on sulphuric acid, hydrogen and zinc sulphate ZnSO_{4},[15 bis]
are obtained. The latter is a solid substance, soluble in water. In
order that the action of the metal on the acid should go on regularly,
and to the end, it is necessary that the acid should be diluted with
water, which dissolves the salt as it is formed; otherwise the salt
covers the metal, and hinders the acid from attacking it. Usually the
acid is diluted with from three to five times its volume of water, and
the metal is covered with this solution. In order that the metal should
act rapidly on the acid, it should present a large surface, so that a
maximum amount of the reacting substances may come into contact in a
given time. For this purpose the zinc is used as strips of sheet zinc,
or in the granulated form (that is, zinc which has been poured from a
certain height, in a molten state, into water). The iron should be in
the form of wire, nails, filings, or cuttings.

[Illustration: FIG. 20.--Apparatus for the preparation of hydrogen from
zinc and sulphuric acid.]

  [10] The composition of water, as we saw above, was determined by
       passing steam over red-hot iron; the same method has been used
       for making hydrogen for filling balloons. An oxide having the
       composition Fe_{3}O_{4} is formed in the reaction, so that it is
       expressed by the equation 3Fe + 4H_{2}O = Fe_{3}O_{4} + 8H.

  [11] The reaction between iron and water (note 10) is reversible. By
       heating the oxide in a current of hydrogen, water and iron are
       obtained. From this it follows, from the principle of chemical
       equilibria, that if iron and hydrogen be taken, and also oxygen,
       but in such a quantity that it is insufficient for combination
       with both substances, then it will divide itself between the two;
       part of it will combine with the iron and the other part with
       the hydrogen, but a portion of both will remain in an uncombined
       state.

       Therefore, if iron and water be placed in a closed space,
       decomposition of the water will proceed on heating to the
       temperature at which the reaction 3Fe + 4H_{2}O = Fe_{3}O_{4} +
       8H commences; but it ceases, does not go on to the end, because
       the conditions for a reverse reaction are attained, and a state
       of equilibrium will ensue after the decomposition of a certain
       quantity of water. Here again (_see_ note 9) the reversibility is
       connected with the small heat effect, and again both reactions
       (direct and reverse) proceed at a red heat. But if, in the
       above-described reaction, the hydrogen escapes as it is evolved,
       then its partial pressure does not increase with its formation,
       and therefore all the iron can he oxidised by the water. In
       this we see the elements of that influence of mass to which we
       shall have occasion to return later. With copper and lead there
       will be no decomposition, either at the ordinary or at a high
       temperature, because the affinity of these metals for oxygen is
       much less than that of hydrogen.

  [12] In general, if reversible as well as non-reversible reactions
       can take place between substances acting on each other, then,
       judging by our present knowledge, the non-reversible reactions
       take place in the majority of cases, which obliges one to
       acknowledge the action, in this case, of comparatively strong
       affinities. The reaction, Zn + H_{2}SO_{4} = H_{2} + ZnSO_{4},
       which takes place in solutions at the ordinary temperature, is
       scarcely reversible under these conditions, but at a certain high
       temperature it becomes reversible, because at this temperature
       zinc sulphate and sulphuric acid split up, and the action must
       take place between the water and zinc. From the preceding
       proposition results proceed which are in some cases verified
       by experiment. If the action of zinc or iron on a solution of
       sulphuric acid presents a non-reversible reaction, then we may
       by this means obtain hydrogen in a very compressed state, and
       compressed hydrogen will not act on solutions of sulphates of
       the above-named metals. This is verified in reality as far as
       was possible in the experiments to keep up the compression or
       pressure of the hydrogen. Those metals which do not evolve
       hydrogen with acids, on the contrary, should, at least at an
       increase of pressure, be displaced by hydrogen. And in fact
       Brunner showed that gaseous hydrogen displaces platinum and
       palladium from the aqueous solutions of their chlorine compounds,
       but not gold, and Beketoff succeeded in showing that silver and
       mercury, under a considerable pressure, are separated from the
       solutions of certain of their compounds by means of hydrogen.
       Reaction already commences under a pressure of six atmospheres,
       if a weak solution of silver sulphate be taken; with a stronger
       solution a much greater pressure is required, however, for the
       separation of the silver.

  [13] For the same reason, many metals in acting on solutions of the
       alkalis displace hydrogen. Aluminium acts particularly clearly
       in this respect, because its oxide gives a soluble compound with
       alkalis. For the same reason tin, in acting on hydrochloric acid,
       evolves hydrogen, and silicon does the same with hydrofluoric
       acid. It is evident that in such cases the sum of all the
       affinities plays a part; for instance, taking the action of
       zinc on sulphuric acid, we have the affinity of zinc for oxygen
       (forming zinc oxide, ZnO), the affinity of its oxide for
       sulphuric anhydride, SO_{3} (forming zinc sulphate, ZnSO_{4}),
       and the affinity of the resultant salt, ZnSO_{4}, for water.
       It is only the first-named affinity that acts in the reaction
       between water and the metal, if no account is taken of those
       forces (of a physico-mechanical character) which act between the
       molecules (for instance, the cohesion between the molecules of
       the oxide) and those forces (of a chemical character) which act
       between the atoms forming the molecule, for instance, between
       the atoms of hydrogen giving the molecule H_{2} containing two
       atoms. I consider it necessary to remark, that the hypothesis of
       the affinity or endeavour of heterogeneous atoms to enter into a
       common system and in harmonious motion (_i.e._ to form a compound
       molecule) must inevitably be in accordance with the hypothesis of
       forces including homogeneous atoms to form complex molecules (for
       instance, H_{2}), and to build up the latter into solid or liquid
       substances, in which the existence of an attraction between the
       homogeneous particles must certainly be admitted. Therefore,
       those forces which bring about solution must also be taken into
       consideration. These are all forces of one and the same series,
       and in this may be seen the great difficulties surrounding the
       study of molecular mechanics and its province--chemical mechanics.

  [14] It is acknowledged that zinc itself acts on water, even at the
       ordinary temperature, but that the action is confined to small
       masses and only proceeds at the surface. In reality, zinc, in the
       form of a very fine powder, or so-called 'zinc dust,' is capable
       of decomposing water with the formation of oxide (hydrated) and
       hydrogen. The oxide formed acts on sulphuric acid, water then
       dissolves the salt produced, and the action continues because
       one of the products of the action of water on zinc, zinc oxide,
       is removed from the surface. One might naturally imagine that
       the reaction does not proceed directly between the metal and
       water, but between the metal and the acid, but such a simple
       representation, which we shall cite afterwards, hides the
       mechanism of the reaction, and does not permit of its actual
       complexity being seen.

  [15] According to Thomsen the reaction between zinc and a very weak
       solution of sulphuric acid evolves about 38,000 calories (zinc
       sulphate being formed) per 65 parts by weight of zinc; and 56
       parts by weight of iron--which combine, like 65 parts by weight
       of zinc, with 16 parts by weight of oxygen--evolve about 25,000
       calories (forming ferrous sulphate, FeSO_{4}). Paracelsus
       observed the action of metals on acids in the seventeenth
       century; but it was not until the eighteenth century that Lémery
       determined that the gas which is evolved in this action is a
       particular one which differs from air and is capable of burning.
       Even Boyle confused it with air. Cavendish determined the chief
       properties of the gas discovered by Paracelsus. At first it was
       called 'inflammable air'; later, when it was recognised that in
       burning it gives water, it was called hydrogen, from the Greek
       words for water and generator.

  [15 bis] If, when the sulphuric acid is poured over the zinc, the
      evolution of the hydrogen proceed too slowly, it may be greatly
      accelerated by adding a small quantity of a solution of CuSO_{4}
      or PtCl_{4} to the acid. The reason of this is explained in Chap.
      XVI., note 10 bis.

The usual method of obtaining hydrogen is as follows:--A certain
quantity of granulated zinc is put into a double-necked, or Woulfe's,
bottle. Into one neck a funnel is placed, reaching to the bottom of
the bottle, so that the liquid poured in may prevent the hydrogen from
escaping through it. The gas escapes through a special gas conducting
tube, which is firmly fixed, by a cork, into the other neck, and ends in
a water bath (fig. 20), under the orifice of a glass cylinder full of
water.[16] If sulphuric acid be now poured into the Woulfe's bottle it
will soon be seen that bubbles of a gas are evolved, which is hydrogen.
The first part of the gas evolved should not be collected, as it is
mixed with the air originally in the apparatus. This precaution should
be taken in the preparation of all gases. Time must be allowed for the
gas evolved to displace all the air from the apparatus, otherwise in
testing the combustibility of the hydrogen an explosion may occur from
the formation of detonating gas (the mixture of the oxygen of the air
with the hydrogen).[17]

  [16] As laboratory experiments with gases require a certain preliminary
       knowledge, we will describe certain _practical methods for the
       collection and preparation of gases_. When in laboratory practice
       an intermittent supply of hydrogen (or other gas which is evolved
       without the aid of heat) is required the apparatus represented
       in fig. 21 is the most convenient. It consists of two bottles,
       having orifices at the bottom, in which corks with tubes are
       placed, and these tubes are connected by an india-rubber tube
       (sometimes furnished with a spring clamp). Zinc is placed in one
       bottle, and dilute sulphuric acid in the other. The neck of the
       former is closed by a cork, which is fitted with a gas-conducting
       tube with a stopcock. If the two bottles are connected with each
       other and the stopcock be opened, the acid will flow to the zinc
       and evolve hydrogen. If the stopcock be closed, the hydrogen
       will force out the acid from the bottle containing the zinc, and
       the action will cease. Or the vessel containing the acid may be
       placed at a lower level than that containing the zinc, when all
       the liquid will flow into it, and in order to start the action
       the acid vessel may be placed on a higher level than the other,
       and the acid will flow to the zinc. It can also be employed for
       collecting gases (as an aspirator or gasometer).

       [Illustration: FIG. 21.--A very convenient apparatus for the
       preparation of gases obtained without heat. It may also replace
       an aspirator or gasometer.]

       In laboratory practice, however, other forms of apparatus are
       generally employed for exhausting, collecting, and holding gases.
       We will here cite the most usual forms. An _aspirator_ usually
       consists of a vessel furnished with a stopcock at the bottom.
       A stout cork, through which a glass tube passes, is fixed into
       the neck of this vessel. If the vessel be filled up with water
       to the cork and the bottom stopcock is opened, then the water
       will run out and draw gas in. For this purpose the glass tube is
       connected with the apparatus from which it is desired to pump out
       or exhaust the gas.

       [Illustration: FIG. 22.--Continuous aspirator. The tube _d_
       should be more than 32 feet long.]

       The aspirator represented in fig. 22 may be recommended for its
       continuous action. It consists of a tube _d_ which widens out at
       the top, the lower part being long and narrow. In the expanded
       upper portion _c_, two tubes are sealed; one, _e_, for drawing in
       the gas, whilst the other, _b_, is connected to the water supply
       _w_. The amount of water supplied through the tube _b_ must be
       less than the amount which can be carried off by the tube _d_.
       Owing to this the water in the tube _d_ will flow through it in
       cylinders alternating with cylinders of gas, which will be thus
       carried away. The gas which is drawn through may be collected
       from the end of the tube _d_, but this form of pump is usually
       employed where the air or gas aspirated is not to be collected.
       If the tube _d_ is of considerable length, say 40 ft. or more,
       a very fair vacuum will be produced, the amount of which is
       shown by the gauge _g_; it is often used for filtering under
       reduced pressure, as shown in the figure. If water be replaced by
       mercury, and the length of the tube _d_ be greater than 760 mm.,
       the aspirator may be employed as an air-pump, and all the air may
       be exhausted from a limited space; for instance, by connecting
       _g_ with a hollow sphere.

       [Illustration: FIG. 23.--Gasholder.]

       _Gasholders_ are often used for collecting and holding gases.
       They are made of glass, copper, or tin plate. The usual form
       is shown in fig. 23. The lower vessel _B_ is made hermetically
       tight--_i.e._, impervious to gases--and is filled with water.
       A funnel is attached to this vessel (on several supports).
       The vessel _B_ communicates with the bottom of the funnel by
       a stopcock _b_ and a tube _a_, reaching to the bottom of the
       vessel _B_. If water be poured into the funnel and the stopcocks
       _a_ and _b_ opened, the water will run through _a_, and the air
       escape from the vessel _B_ by _b_. A glass tube _f_ runs up the
       side of the vessel _B_, with which it communicates at the top
       and bottom, and shows the amount of water and gas the gasholder
       contains. In order to fill the gasholder with a gas, it is first
       filled with water, the cocks _a_, _b_ and _e_ are closed, the nut
       _d_ unscrewed, and the end of the tube conducting the gas from
       the apparatus in which it is generated is passed into _d_. As
       the gas fills the gasholder, the water runs out at _d_. If the
       pressure of a gas be not greater than the atmospheric pressure
       and it be required to collect it in the gasholder, then the
       stopcock _e_ is put into communication with the space containing
       the gas. Then, having opened the orifice _d_, the gasholder
       acts like an aspirator; the gas will pass through _e_, and the
       water run out at _d_. If the cocks be closed, the gas collected
       in the gasholder may be easily preserved and transported. If
       it be desired to transfer this gas into another vessel, then a
       gas-conducting tube is attached to _e_, the cock _a_ opened, _b_
       and _d_ closed, and the gas will then pass out at _e_, owing to
       its pressure in the apparatus being greater than the atmospheric
       pressure, due to the pressure of the water poured into the
       funnel. If it be required to fill a cylinder or flask with the
       gas, it is filled with water and inverted in the funnel, and the
       stopcocks _b_ and _a_ opened. Then water will run through _a_,
       and the gas will escape from the gasholder into the cylinder
       through _b_.

  [17] When it is required to prepare hydrogen in large quantities for
       filling balloons, copper vessels or wooden casks lined with
       lead are employed; they are filled with scrap iron, over which
       dilute sulphuric acid is poured. The hydrogen generated from a
       number of casks is carried through lead pipes into special casks
       containing water (in order to cool the gas) and lime (in order
       to remove acid fumes). To avoid loss of gas all the joints are
       made hermetically tight with cement or tar. In order to fill his
       gigantic balloon (of 25,000 cubic metres capacity), Giffard, in
       1878, constructed a complicated apparatus for giving a continuous
       supply of hydrogen, in which a mixture of sulphuric acid and
       water was continually run into vessels containing iron, and
       from which the solution of iron sulphate formed was continually
       drawn off. When coal gas, extracted from coal, is employed for
       filling balloons, it should be as light, or as rich in hydrogen,
       as possible. For this reason, only the last portions of the
       gas coming from the retorts are collected, and, besides this,
       it is then sometimes passed through red-hot vessels, in order
       to decompose the hydrocarbons as much as possible; charcoal
       is deposited in the red-hot vessels, and hydrogen remains as
       gas. Coal gas may be yet further enriched in hydrogen, and
       consequently rendered lighter, by passing it over an ignited
       mixture of charcoal and lime.

       L. Mond (London) proposes to manufacture hydrogen on a large
       scale from water gas (_see infra_, and Chapters VIII. and IX.),
       which contains a mixture of oxide of carbon (CO) and hydrogen,
       and is produced by the action of steam upon incandescent coke
       (C + H_{2}O = CO + H_{2}). He destroys the oxide of carbon
       by converting it into carbon and carbonic anhydride (2CO =
       C + CO_{2}), which is easily done by means of incandescent,
       finely-divided metallic nickel; the carbon then remains with
       the nickel, from which it may be removed by burning it in air,
       and the nickel can then be used over again (_see_ Chapter IX.,
       Note 24 bis). The CO_{2} formed is removed from the hydrogen by
       passing it through milk of lime. This process should apparently
       give hydrogen on a large scale more economically than any of the
       methods hitherto proposed.

Hydrogen, besides being contained in water, is also contained in many
other substances,[18] and may be obtained from them. As examples of
this, it may be mentioned (1) that a mixture of formate of sodium,
CHNaO_{2}, and caustic soda, NaHO, when heated to redness, forms sodium
carbonate, Na_{2}CO_{3}, and hydrogen, H_{2};[19] (2) that a number of
organic substances are decomposed at a red heat, forming hydrogen, among
other gases, and thus it is that hydrogen is contained in ordinary coal
gas.

  [18] Of the metals, only a very few combine with hydrogen (for example,
       sodium), and give substances which are easily decomposed. Of
       the non-metals, the halogens (fluorine, chlorine, bromine,
       and iodine) most easily form hydrogen compounds; of these the
       hydrogen compound of chlorine, and still more that of fluorine,
       is stable, whilst those of bromine and iodine are easily
       decomposed, especially the latter. The other non-metals--for
       instance, sulphur, carbon, and phosphorus--give hydrogen
       compounds of different composition and properties, but they are
       all less stable than water. The number of the carbon compounds of
       hydrogen is enormous, but there are very few among them which are
       not decomposed, with separation of the carbon and hydrogen, at a
       red heat.

  [19] The reaction expressed by the equation CNaHO_{2} + NaHO =
       CNa_{2}O_{3} + H_{2} may be effected in a glass vessel, like
       the decomposition of copper carbonate or mercury oxide (_see_
       Introduction); it is non-reversible, and takes place without the
       presence of water, and therefore Pictet (_see_ later) made use of
       it to obtain hydrogen under great pressure.

Charcoal itself liberates hydrogen from steam at a high temperature;[20]
but the reaction which here takes place is distinguished by a certain
complexity, and will therefore be considered later.

  [20] The reaction between charcoal and superheated steam is a double
       one--that is, there may be formed either carbonic oxide, CO
       (according to the equation H_{2}O + C = H_{2} + CO), or carbonic
       anhydride CO_{2} (according to the equation 2H_{2}O + C = 2H_{2}
       + CO_{2}), and the resulting mixture is called _water-gas_; we
       shall speak of it in Chapter IX.

_The properties of hydrogen._--Hydrogen presents us with an example
of a gas which at first sight does not differ from air. It is not
surprising, therefore, that Paracelsus, having discovered that an
aëriform substance is obtained by the action of metals on sulphuric
acid, did not determine exactly its difference from air. In fact,
hydrogen, like air, is colourless, and has no smell;[21] but a more
intimate acquaintance with its properties proves it to be entirely
different from air. The first sign which distinguishes hydrogen from
air is its combustibility. This property is so easily observed that
it is the one to which recourse is usually had in order to recognise
hydrogen, if it is evolved in a reaction, although there are many other
combustible gases. But before speaking of the combustibility and other
chemical properties of hydrogen, we will first describe the physical
properties of this gas, as we did in the case of water. It is easy to
show that it is one of the lightest gases.[22] If passed into the bottom
of a flask full of air, hydrogen will not remain in it, but, owing
to its lightness, rapidly escapes and mixes with the atmosphere. If,
however, a cylinder whose orifice is turned downwards be filled with
hydrogen, it will not escape, or, more correctly, it will only slowly
mix with the atmosphere. This may be demonstrated by the fact that a
lighted taper sets fire to the hydrogen at the orifice of the cylinder,
and is itself extinguished inside the cylinder. Hence, hydrogen, being
itself combustible, does not support combustion. The great lightness
of hydrogen is taken advantage of for balloons. Ordinary coal gas,
which is often also used for the same purpose, is only about twice
as light as air, whilst hydrogen is 14-1/2 times lighter than air.
A very simple experiment with soap bubbles very well illustrates the
application of hydrogen for filling balloons. Charles, of Paris, showed
the lightness of hydrogen in this way, and constructed a balloon filled
with hydrogen almost simultaneously with Montgolfier. One litre of pure
and dry hydrogen[23] at 0° and 760 mm. pressure weighs 0·08986 gram;
that is, hydrogen is almost 14-1/2 (more exactly, 14·39) times lighter
than air. It is the lightest of all gases. The small density of hydrogen
determines many remarkable properties which it shows; thus, hydrogen
passes exceedingly rapidly through fine orifices, its molecules (Chapter
I.) being endued with the greatest velocity.[24] At pressures somewhat
higher than the atmospheric pressure, all other gases exhibit a greater
compressibility and co-efficient of expansion than they should according
to the laws of Mariotte and Gay-Lussac; whilst hydrogen, on the
contrary, is compressed to a less degree than it should be from the law
of Mariotte,[25] and with a rise of pressure it expands slightly less
than at the atmospheric pressure.[26] However, hydrogen, like air and
many other gases which are permanent at the ordinary temperature, does
not pass into a liquid state under a very considerable pressure,[27] but
is compressed into a lesser volume than would follow from Mariotte's
law.[28] From this it may be concluded that the absolute boiling point
of hydrogen, and of gases resembling it,[29] lies very much below the
ordinary temperature; that is, that the liquefaction of this _gas_
is only possible at low temperatures, and under great pressures.[30]
This conclusion was verified (1877) by the experiments of Pictet
and Cailletet.[31] They compressed gases at a very low temperature,
and then allowed them to expand, either by directly decreasing the
pressure or by allowing them to escape into the air, by which means
the temperature fell still lower, and then, just as steam when rapidly
rarefied[32] deposits liquid water in the form of a fog, hydrogen in
expanding forms a fog, thus indicating its passage into a liquid state.
But as yet it has been impossible to preserve this liquid, even for a
short time, to determine its properties, notwithstanding the employment
of a temperature of -200° and a pressure of 200 atmospheres,[33] although
by these means the gases of the atmosphere may be kept in a liquid
state for a long time. This is due to the fact that the absolute boiling
point of hydrogen lies lower than that of all other known gases, which
also depends on the extreme lightness of hydrogen.[34]

  [21] Hydrogen obtained by the action of zinc or iron on sulphuric acid
       generally smells of hydrogen sulphide (like rotten eggs), which
       it contains in admixture. As a rule such hydrogen is not so
       pure as that obtained by the action of an electric current or
       of sodium on water. The impurity of the hydrogen depends on the
       impurities contained in the zinc, or iron, and sulphuric acid,
       and on secondary reactions which take place simultaneously with
       the main reaction. Impure hydrogen may be easily freed from
       the impurities it contains: some of them--namely, those having
       acid properties--are absorbed by caustic soda, and therefore
       may be removed by passing the hydrogen through a solution of
       this substance; another series of impurities is absorbed by a
       solution of mercuric chloride; and, lastly, a third series is
       absorbed by a solution of potassium permanganate. If absolutely
       _pure hydrogen_ be required, it is sometimes obtained by the
       decomposition of water (previously boiled to expel all air,
       and mixed with pure sulphuric acid) by the galvanic current.
       Only the gas evolved at the negative electrode is collected.
       Or else, an apparatus like that which gives detonating gas is
       used, the positive electrode, however, being immersed under
       mercury containing zinc in solution. The oxygen which is evolved
       at this electrode then immediately, at the moment of its
       evolution, combines with the zinc, and this compound dissolves
       in the sulphuric acid and forms zinc sulphate, which remains in
       solution, and therefore the hydrogen generated will be quite free
       from oxygen.

  [22] An inverted beaker is attached to one arm of the beam of a
       tolerably sensitive balance, and its weight counterpoised by
       weights in the pan attached to the other arm, If the beaker
       be then filled with hydrogen it rises, owing to the air being
       replaced by hydrogen. Thus, at the ordinary temperature of a
       room, a litre of air weighs about 1·2 gram, and on replacing the
       air by hydrogen a decrease in weight of about 1 gram per litre is
       obtained. Moist hydrogen is heavier than dry--for aqueous vapour
       is nine times heavier than hydrogen. In filling balloons it is
       usually calculated that (it being impossible to have perfectly
       dry hydrogen or to obtain it quite free from air) the lifting
       force due to the difference between the weights of equal volumes
       of hydrogen and air is equal to 1 kilogram (= 1,000 grams) per
       cubic metre (= 1,000 litres).

  [23] The density of hydrogen in relation to the air has been repeatedly
       determined by accurate experiments. The first determination, made
       by Lavoisier, was not very exact; taking the density of air as
       unity, he obtained 0·0769 for that of hydrogen--that is, hydrogen
       as thirteen times lighter than air. More accurate determinations
       are due to Thomsen, who obtained the figure 0·0693; Berzelius
       and Dulong, who obtained 0·0688; and Dumas and Boussingault, who
       obtained 0·06945. Regnault, and more recently Le Duc (1892),
       took two spheres of considerable capacity, which contained equal
       volumes of air (thus avoiding the necessity of any correction
       for weighing them in air). Both spheres were attached to the
       scale pans of a balance. One was sealed up, and the other first
       weighed empty and then full of hydrogen. Thus, knowing the weight
       of the hydrogen filling the sphere, and the capacity of the
       sphere, it was easy to find the weight of a litre of hydrogen;
       and, knowing the weight of a litre of air at the same temperature
       and pressure, it was easy to calculate the density of hydrogen.
       Regnault, by these experiments, found the average density of
       hydrogen to be 0·06926 in relation to air; Le Duc, 0·06948 (with
       a possible error of ±0·00001), and this latter figure must now be
       looked upon as near to the truth.

       In this work I shall always refer the densities of all gases to
       hydrogen, and not to air; I will therefore give, for the sake of
       clearness, the weight of a litre of dry pure hydrogen in grams
       at a temperature _t_° and under a pressure _H_ (measured in
       millimetres of mercury at 0°, in lat. 45°). The weight of a litre
       of hydrogen

            = 0·08986 × (_H_/760) × 1/(1 + 0·00367_t_) gram.

       For aëronauts it is very useful to know, besides this, the
       weight of the air at different heights, and I therefore insert
       the adjoining table, constructed on the basis of Glaisher's
       data, for the temperature and moisture of the atmospheric strata
       in clear weather. All the figures are given in the metrical
       system--1,000 millimetres = 39·37 inches, 1,000 kilograms =
       2204·3375 lbs., 1,000 cubic metres = 35,316·6 cubic feet. The
       starting temperature at the earth's surface is taken as = 15° C.,
       its moisture 60 p.c., pressure 760 millimetres. The pressures
       are taken as indicated by an _aneroid barometer_, assumed to be
       corrected at the sea level and at lat. 45° C. If the height above
       the level of the sea equal _z_ kilometres, then the weight of 1
       cubic metre of air may be approximately taken as 1·222- 0·12_z_ +
       0·00377_z_^2 kilogram.

       +--------+-----------+--------+--------+--------------------+
       |Pressure|Temperature|Moisture| Height | Weight of the air  |
       |        |           |        |(metres)|(1,000 cubic metres)|
       |--------+-----------+--------+--------+--------------------+
       | 760 mm.|  15°   C. | 60 p.c.|     0  |   1222  kilos.     |
       | 700  " |  11·0° "  | 64  "  |   690  |   1141    "        |
       | 650  " |   7·6° "  | 64  "  |  1300  |   1073    "        |
       | 600  " |   4·3° "  | 63  "  |  1960  |   1003    "        |
       | 550  " |  -1·0° "  | 62  "  |  2660  |    931    "        |
       | 500  " |  -2·4° "  | 58  "  |  3420  |    857    "        |
       | 450  " |  -5·8° "  | 52  "  |  4250  |    781    "        |
       | 400  " |  -9·1° "  | 44  "  |  5170  |    703    "        |
       | 350  " | -12·5° "  | 36  "  |  6190  |    624    "        |
       | 300  " | -15·9° "  | 27  "  |  7360  |    542    "        |
       | 250  " | -19·2° "  | 18  "  |  8720  |    457    "        |
       +--------+-----------+--------+--------+--------------------+

       Although the figures in this table are calculated with every
       possible care from average data, yet they can only be taken
       approximately, for in every separate case the conditions, both at
       the earth's surface and in the atmosphere, will differ from those
       here taken. In calculating the height to which a balloon can
       ascend, it is evident that the density of gas in relation to air
       must be known. This density for ordinary coal gas is from 0·6 to
       0·35, and for hydrogen with its ordinary contents of moisture and
       air from 0·1 to 0·15.

       Hence, for instance, it may be calculated that a balloon of 1,000
       cubic metres capacity filled with pure hydrogen, and weighing
       (the envelope, tackle, people, and ballast) 727 kilograms, will
       only ascend to a height of about 4,250 metres.

  [24] If a cracked flask be filled with hydrogen and its neck immersed
       under water or mercury, then the liquid will rise up into the
       flask, owing to the hydrogen passing through the cracks about 3·8
       times quicker than the air is able to pass through these cracks
       into the flask. The same phenomenon may be better observed if,
       instead of a flask, a tube be employed, whose end is closed by a
       porous substance, such as graphite, unglazed earthenware, or a
       gypsum plate.

  [25] According to Boyle and Mariotte's law, for a given gas at a
       constant temperature the volume decreases by as many times as
       the pressure increases; that is, this law requires that the
       product of the volume _v_ and the pressure _p_ for a given gas
       should be a constant quantity: _pv_ = _C_, a constant quantity
       which does not vary with a change of pressure. This equation
       does very nearly and exactly express the observed relation
       between the volume and pressure, but only within comparatively
       small variations of pressure, density, and volume. If these
       variations be in any degree considerable, the quantity _pv_
       proves to be dependent on the pressure, and it either increases
       or diminishes with an increase of pressure. In the former case
       the compressibility is less than it should he according to
       Mariotte's law, in the latter case it is greater. We will call
       the first case a positive discrepancy (because then _d(pv)/d(p)_
       is greater than zero), and the second case a negative discrepancy
       (because then _d(pv)/d(p)_ is less than zero). Determinations
       made by myself (in the seventies), M. L. Kirpicheff, and V. A.
       Hemilian showed that all known gases at low pressures--_i.e._
       when considerably rarefied--present positive discrepancies. On
       the other hand, it appears from the researches of Cailletet,
       Natterer, and Amagat that all gases under great pressures (when
       the volume obtained is 500-1,000 times less than under the
       atmospheric pressure) also present positive discrepancies. Thus
       under a pressure of 2,700 atmospheres air is compressed, not
       2,700 times, but only 800, and hydrogen 1,000 times. Hence the
       positive kind of discrepancy is, so to say, normal to gases. And
       this is easily intelligible. If a gas followed Mariotte's law, or
       if it were compressed to a greater extent than is shown by this
       law, then under great pressures it would attain a density greater
       than that of solid and liquid substances, which is in itself
       improbable and even impossible by reason of the fact that solid
       and liquid substances are themselves but little compressible.
       For instance, a cubic centimetre of oxygen at 0° and under the
       atmospheric pressure weighs about 0·0014 gram, and at a pressure
       of 3,000 atmospheres (this pressure is attained in guns) it
       would, if it followed Mariotte's law, weigh 4·2 grams--that is,
       would be about four times heavier than water--and at a pressure
       of 10,000 atmospheres it would be heavier than mercury. Besides
       this, positive discrepancies are probable because the molecules
       of a gas themselves must occupy a certain volume. Considering
       that Mariotte's law, strictly speaking, applies only to the
       intermolecular space, we can understand the necessity of positive
       discrepancies. If we designate the volume of the molecules of a
       gas by _b_ (like van der Waals, _see_ Chap. I., Note 34), then it
       must be expected that _p(v-b) = C_. Hence _pv = C + bp_, which
       expresses a positive discrepancy. Supposing that for hydrogen
       _pv_ = 1,000, at a pressure of one metre of mercury, according to
       the results of Regnault's, Amagat's, and Natterer's experiments,
       we obtain _b_ as approximately 0·7 to 0·9.

       Thus the increase of _pv_ with the increase of pressure must
       be considered as the normal law of the compressibility of
       gases. Hydrogen presents such a positive compressibility at
       all pressures, for it presents positive discrepancies from
       Mariotte's law, according to Regnault, at all pressures above the
       atmospheric pressure. Hence hydrogen is, so to say, a perfect
       gas. No other gas behaves so simply with a change of pressure.
       All other gases at pressures from 1 to 30 atmospheres present
       negative discrepancies--that is, they are then compressed to a
       greater degree than should follow from Mariotte's law, as was
       shown by the determinations of Regnault, which were verified
       when repeated by myself and Boguzsky. Thus, for example, on
       changing the pressure from 4 to 20 metres of mercury--that is,
       on increasing the pressure five times--the volume only decreased
       4·93 times when hydrogen was taken, and 5·06 when air was taken.

       The positive discrepancies from the law at low pressures are
       of particular interest, and, according to the above-mentioned
       determinations made by myself, Kirpicheff, and Hemilian, and
       verified (by two methods) by K. D. Kraevitch and Prof. Ramsay
       (London, 1894), they are proper to all gases (even to those which
       are easily compressed into a liquid state, such as carbonic and
       sulphurous anhydrides). These discrepancies approach the case
       of a very high rarefaction of gases, where a gas is near to a
       condition of maximum dispersion of its molecules, and perhaps
       presents a passage towards the substance termed 'luminiferous
       ether' which fills up interplanetary and interstellar space. If
       we suppose that gases are rarefiable to a definite limit only,
       having attained which they (like solids) do not alter in volume
       with a decrease of pressure, then on the one hand the passage of
       the atmosphere at its upper limits into a homogeneous ethereal
       medium becomes comprehensible, and on the other hand it would be
       expected that gases would, in a state of high rarefaction (_i.e._
       when small masses of gases occupy large volumes, or when furthest
       removed from a liquid state), present positive discrepancies from
       Boyle and Mariotte's law. Our present acquaintance with this
       province of highly rarefied gases is very limited (because direct
       measurements are exceedingly difficult to make, and are hampered
       by possible errors of experiment, which may be considerable), and
       its further development promises to elucidate much in respect to
       natural phenomena. To the three states of matter (solid, liquid,
       and gaseous) it is evident a fourth must yet be added, the
       ethereal or ultra-gaseous (as Crookes proposed), understanding by
       this, matter in its highest possible state of rarefaction.

  [26] The law of Gay-Lussac states that all gases in all conditions
       present one coefficient of expansion 0·00367; that is, when
       heated from 0° to 100° they expand like air; namely, a thousand
       volumes of a gas measured at 0° will occupy 1367 volumes at
       100°. Regnault, about 1850, showed that Gay-Lussac's law is not
       entirely correct, and that different gases, and also one and
       the same gas at different pressures, have not quite the same
       coefficients of expansion. Thus the expansion of air between 0°
       and 100° is 0·367 under the ordinary pressure of one atmosphere,
       and at three atmospheres it is 0·371, the expansion of hydrogen
       is 0·366, and of carbonic anhydride 0·37. Regnault, however,
       did not directly determine the change of volume between 0° and
       100°, but measured the variation of tension with the change of
       temperature; but since gases do not entirely follow Mariotte's
       law, the change of volume cannot be directly judged by the
       variation of tension. The investigations carried on by myself and
       Kayander, about 1870, showed the variation of volume on heating
       from 0° to 100° under a constant pressure. These investigations
       confirmed Regnault's conclusion that Gay-Lussac's law is not
       entirely correct, and further showed (1) that the expansion
       per volume from 0° to 100° under a pressure of one atmosphere,
       for air = 0·368, for hydrogen = 0·367, for carbonic anhydride
       = 0·373, for hydrogen bromide = 0·386, &c.; (2) that for gases
       which are more compressible than should follow from Mariotte's
       law the expansion by heat increases with the pressure--for
       example, for air at a pressure of three and a half atmospheres,
       it equals 0·371, for carbonic anhydride at one atmosphere
       it equals 0·373, at three atmospheres 0·389, and at eight
       atmospheres 0·413; (3) that for gases which are less compressible
       than should follow from Mariotte's law, the expansion by heat
       decreases with an increase of pressure--for example, for hydrogen
       at one atmosphere 0·367, at eight atmospheres 0·369, for air
       at a quarter of an atmosphere 0·370, at one atmosphere 0·368;
       and hydrogen like _air_ (and all gases) is less compressed _at
       low pressures_ than should follow from Mariotte's law (_see_
       Note 25). Hence, hydrogen, starting from zero to the highest
       pressures, exhibits a gradually, although only slightly, varying
       coefficient of expansion, whilst for air and other gases at the
       atmospheric and higher pressures, the coefficient of expansion
       increases with the increase of pressure, so long as their
       compressibility is greater than should follow from Mariotte's
       law. But when at considerable pressures, this kind of discrepancy
       passes into the normal (_see_ Note 25), then the coefficient of
       expansion of all gases decreases with an increase of pressure,
       as is seen from the researches of Amagat. The difference between
       the two coefficients of expansion, for a constant pressure
       and for a constant volume, is explained by these relations.
       Thus, for example, for air at a pressure of one atmosphere the
       true coefficient of expansion (the volume varying at constant
       pressure) = 0·00368 (according to Mendeléeff and Kayander) and
       the variation of tension (at a constant volume, according to
       Regnault) = 0·00367.

  [27] Permanent gases are those which cannot be liquefied by an increase
       of pressure alone. With a rise of temperature, all gases and
       vapours become permanent gases. As we shall afterwards learn,
       carbonic anhydride becomes a permanent gas at temperatures above
       31°, and at lower temperatures it has a maximum tension, and may
       be liquefied by pressure alone.

       _The liquefaction_ of gases, accomplished by Faraday (_see_
       Ammonia, Chapter VI.) and others, in the first half of this
       century, showed that a number of substances are capable, like
       water, of taking all three physical states, and that there is
       no essential difference between vapours and gases, the only
       distinction being that the boiling points (or the temperature
       at which the tension = 760 mm.) of liquids lie above the
       ordinary temperature, and those of liquefied gases below, and
       consequently a gas is a superheated vapour, or vapour heated
       above the boiling point, or removed from saturation, rarefied,
       having a lower tension than that maximum which is proper to a
       given temperature and substance. We will here cite the _maximum
       tensions_ of certain liquids and gases _at various temperatures_,
       because they may be taken advantage of for obtaining constant
       temperatures by changing the pressure at which boiling or the
       formation of saturated vapours takes place. (I may remark that
       the dependence between the tension of the saturated vapours of
       various substances and the temperature is very complex, and
       usually requires three or four independent constants, which
       vary with the nature of the substance, and are found from the
       dependence of the tension _p_ on the temperature _t_ given
       by experiment; but in 1892 K. D. Kraevitch showed that this
       dependence is determined by the properties of a substance, such
       as its density, specific heat, and latent heat of evaporation.)
       The temperatures (according to the air thermometer) are placed
       on the left, and the tension in millimetres of mercury (at 0°)
       on the right-hand side of the equations. Carbon bisulphide,
       CS_{2}, 0° = 127·9; 10° = 198·5; 20° = 298·1; 30° = 431·6; 40°
       = 617·5; 50° = 857·1. Chlorobenzene, C_{6}H_{5}Cl, 70° = 97·9;
       80° = 141·8; 90° = 208·4; 100° = 292·8; 110° = 402·6; 120° =
       542·8; 130° = 719·0. Aniline, C_{6}H_{7}N, 150° = 283·7; 160°
       = 387·0; 170° = 515·6; 180° = 677·2; 185° = 771·5. Methyl
       salicylate, C_{8}H_{8}O_{3}, 180° = 294·4; 190° = 330·9; 200° =
       432·4; 210° = 557·5; 220° = 710·2; 224° = 779·9. Mercury, Hg,
       300° = 246·8; 310° = 304·9; 320° = 373·7; 330° = 454·4; 340°
       = 548·6; 350° = 658·0; 359° = 770·9. Sulphur, S, 395° = 300;
       423° = 500; 443° = 700; 452° = 800; 459° = 900. These figures
       (Ramsay and Young) show the possibility of obtaining constant
       temperatures in the vapours of boiling liquids by altering
       the pressure. We may add the following boiling points under
       a pressure of 760 mm. (according to the air thermometer by
       Collendar and Griffiths, 1891): aniline, 184° = 13; naphthalene,
       217° = 94; benzophenone, 305° = 82; mercury, 356° = 76;
       triphenyl-methane, 356° = 44; sulphur, 444° = 53. And melting
       points: tin, 231° = 68; bismuth, 269° = 22; lead, 327° = 69;
       and zinc, 417° = 57. These data may be used for obtaining a
       constant temperature and for verifying thermometers. The same
       object may be attained by the melting points of certain salts,
       determined according to the air thermometer by V. Meyer and
       Riddle (1893): NaCl, 851°; NaBr, 727°; NaI, 650°; KCl, 760°;
       KBr, 715°; KI, 623°; K_{2}CO_{3}, 1045°; Na_{2}CO_{3}, 1098°;
       Na_{2}B_{4}O_{7}, 873°; Na_{2}SO_{4}, 843°; K_{2}SO_{4}, 1073°.
       The tension of liquefied gases is expressed in atmospheres.
       Sulphurous anhydride, SO_{2},-30° = 0·4;-20° = 0·6;-10° = 1; 0°
       = 1·5; +10° = 2·3; 20° = 3·2; 30° = 5·3. Ammonia, NH_{3},-40°
       = 0·7;-30° = 1·1;-20° = 1·8; -10° = 2·8; 0° = 4·2; +10° = 6·0;
       20° = 8·4. Carbonic anhydride, CO_{2},-115° = 0·033; -80° =
       1;-70° = 2·1;-60° = 3·9;-50° = 6·8;-40° = 10;-20° = 23; 0° = 35;
       +10° = 46; 20° = 58. Nitrous oxide, N_{2}O,-125° = 0·033;-92° =
       1;-80° = 1·9;-50° = 7·6; -20° = 23·1; 0° = 36·1; +20° = 55·3.
       Ethylene, C_{2}H_{4},-140° = 0·033;-130° = 0·1; -103° = 1;-40°
       = 13;-1° = 42. Air,-191° = 1;-158° = 14;-140° = 39. Nitrogen,
       N_{2},-203° = 0·085;-193° = 1;-160° = 14;-146° = 32. The methods
       of liquefying gases (by pressure and cold) will be described
       under ammonia, nitrous oxide, sulphurous anhydride, and in later
       footnotes. We will now turn our attention to the fact that the
       evaporation of volatile liquids, under various, and especially
       under low, pressures, gives an easy means for obtaining _low
       temperatures_. Thus liquefied carbonic anhydride, under the
       ordinary pressure, reduces the temperature to -80°, and when it
       evaporates in a rarefied atmosphere (under an air-pump) to 25 mm.
       (= 0·033 atmosphere) the temperature, judging by the above-cited
       figures, falls to -115° (Dewar). Even the evaporation of liquids
       of common occurrence, under low pressures easily attainable with
       an air-pump, may produce low temperatures, which may be again
       taken advantage of for obtaining still lower temperatures. Water
       boiling in a vacuum becomes cold, and under a pressure of less
       than 4·5 mm. it freezes, because its tension at 0° is 4·5 mm.
       A sufficiently low temperature may be obtained by forcing fine
       streams of air through common ether, or liquid carbon bisulphide,
       CS_{2}, or methyl chloride, CH_{3}Cl, and other similar volatile
       liquids. In the adjoining table are given, for certain gases, (1)
       the number of atmospheres necessary for their liquefaction at
       15°, and (2) the boiling points of the resultant liquids under a
       pressure of 760 mm.

                   C_{2}H_{4}  N_{2}O  CO_{2},  H_{2}S
                  (1)    42          31     52      10
                  (2)  -103°        -92°   -80°    -74°

           AsH_{3},  NH_{3}  HCl  CH_{3}Cl  C_{2}N_{2}  SO_{2}
          (1)   8         7      25       4        4         3
          (2) -58°      -38°    -35°    -24°     -21°      -10°

  [28] Natterer's determinations (1851-1854), together with Amagat's
       results (1880-1888), show that the compressibility of hydrogen,
       under high pressures, may be expressed by the following figures:--

                  _p_   =     1     100    1000    2500
                  _v_   =     1  0·0107  0·0019  0·0013
                  _pv_  =     1    1·07     1·9    3·25
                  _s_   =  0·11    10·3      58      85

       where _p_ = the pressure in metres of mercury, _v_ = the volume,
       if the volume taken under a pressure of 1 metre = 1, and _s_
       the weight of a litre of hydrogen at 20° in grams. If hydrogen
       followed Mariotte's law, then under a pressure of 2,500 metres,
       one litre would contain not 85, but 265 grams. It is evident
       from the above figures that the weight of a litre of the gas
       approaches a limit as the pressure increases, which is doubtless
       the density of the gas when liquefied, and therefore the weight
       of a litre of liquid hydrogen will probably be near 100 grams
       (density about 0·1, being less than that of all other liquids).

  [29] Cagniard de Latour, on heating ether in a closed tube to about
       190°, observed that at this temperature the liquid is transformed
       into vapour occupying the original volume--that is, having the
       same density as the liquid. The further investigations made by
       Drion and myself showed that every liquid has such an _absolute
       boiling point_, above which it cannot exist as a liquid and
       is transformed into a dense gas. In order to grasp the true
       signification of this absolute boiling temperature, it must be
       remembered that the liquid state is characterised by a cohesion
       of its particles which does not exist in vapours and gases. The
       cohesion of liquids is expressed in their capillary phenomena
       (the breaks in a column of liquid, drop formation, and rise
       in capillary tubes, &c.), and the product of the density of a
       liquid into the height to which it rises in a capillary tube (of
       a definite diameter) may serve as the measure of the magnitude
       of cohesion. Thus, in a tube of 1 mm. diameter, water at 15°
       rises (the height being corrected for the meniscus) 14·8 mm.,
       and ether at _t°_ to a height 5·35-0·028_t°_ mm. The cohesion
       of a liquid is lessened by heating, and therefore the capillary
       heights are also diminished. It has been shown by experiment that
       this decrement is proportional to the temperature, and hence by
       the aid of capillary observations we are able to form an idea
       that at a certain rise of temperature the cohesion may become =
       0. For ether, according to the above formula, this would occur
       at 191°. If the cohesion disappear from a liquid it becomes a
       gas, for cohesion is the only point of difference between these
       two states. A liquid in evaporating and overcoming the force of
       cohesion absorbs heat. Therefore, the absolute boiling point was
       defined by me (1861) as that temperature at which (_a_) a liquid
       cannot exist as a liquid, but forms a gas which cannot pass into
       a liquid state under any pressure whatever; (_b_) cohesion = 0;
       and (_c_) the latent heat of evaporation = 0.

       This definition was but little known until Andrews (1869)
       explained the matter from another aspect. Starting from gases,
       he discovered that carbonic anhydride cannot be liquefied by any
       degree of compression at temperatures above 31°, whilst at lower
       temperatures it can be liquefied. He called this temperature the
       _critical temperature_. It is evident that it is the same as
       the absolute boiling point. We shall afterwards designate it by
       _tc_. At low temperatures a gas which is subjected to a pressure
       greater than its maximum tension (Note 27) is transformed into
       a liquid, which, in evaporating, gives a saturated vapour
       possessing this maximum tension; whilst at temperatures above
       tc the pressure to which the gas is subjected may increase
       indefinitely. However, under these conditions the volume of the
       gas does not change indefinitely but approaches a definite limit
       (_see_ Note 28)--that is, it resembles in this respect a liquid
       or a solid which is altered but little in volume by pressure.
       The volume which a liquid or gas occupies at _tc_ is termed the
       _critical volume_, and corresponds with the _critical pressure_,
       which we will designate by _pc_ and express in atmospheres. It
       is evident from what has been said that the discrepancies from
       Mariotte and Boyle's law, the absolute boiling point, the density
       in liquid and compressed gaseous states, and the properties of
       liquids, must all he intimately connected together. We will
       consider these relations in one of the following notes. At
       present we will supplement the above observations by the values
       of _tc_ and _pc_ for certain liquids and gases which have been
       investigated in this respect--

       +---------------------+------++----------------------+------+
       |                _tc_ | _pc_ ||                _tc_  | _pc_ |
       +---------------------+------++----------------------+------+
       | N_{2}         -146° |  33  || H_{2}S        +108°  |  92  |
       | CO            -140° |  39  || C_{2}N_{2}    +124°  |  62  |
       | O_{2}         -119° |  50  || NH_{3}        +131°  | 114  |
       | CH_{4}        -100° |  50  || CH_{3}Cl      +141°  |  73  |
       | NO             -93° |  71  || SO_{2}        +155°  |  79  |
       | C_{2}H_{4}     +10° |  51  || C_{5}H_{10}   +192°  |  34  |
       | CO_{2}         +32° |  77  || C_{4}H_{10}O  +193°  |  40  |
       | N_{2}O         +53° |  75  || CHCl_{3}      +268°  |  55  |
       | C_{2}H_{2}     +37° |  68  || CS_{2}        +278°  |  78  |
       | HCl            +52° |  86  || C_{6}H_{6}    +292°  |  60  |
       | H_{2}O        +365° | 200  || C_{6}H_{5}F   +287°  |  45  |
       | CH_{3}OH      +240° |  79  || C_{6}H_{5}Cl  +360°  |  45  |
       | C_{2}H_{5}OH  +243° |  63  || C_{6}H_{5}Br  +397°  |  45  |
       | CH_{3}COOH    +322° |  57  || C_{6}H_{5}I   +448°  |  45  |
       +---------------------+------++----------------------+------+

       Young and Guy (1891) showed that _tc_ and _pc_ clearly depend
       upon the composition and molecular weight.

  [30] I came to this conclusion in 1870 (_Ann. Phys. Chem._ 141, 623).

  [31] Pictet, in his researches, effected the direct liquefaction of
       many gases which up to that time had not been liquefied. He
       employed the apparatus used for the manufacture of ice on a large
       scale, employing the vaporisation of liquid sulphurous anhydride,
       which may be liquefied by pressure alone. This anhydride is a gas
       which is transformed into a liquid at the ordinary temperature
       under a pressure of several atmospheres (_see_ Note 27), and
       boils at -10° at the ordinary atmospheric pressure. This liquid,
       like all others, boils at a lower temperature under a diminished
       pressure, and by continually pumping out the gas which comes
       off by means of a powerful air-pump its boiling point falls as
       low as -75°. Consequently, if on the one hand we force liquid
       sulphurous anhydride into a vessel, and on the other hand pump
       out the gas from the same vessel by powerful air-pumps, then the
       liquefied gas will boil in the vessel, and cause the temperature
       in it to fall to -75°. If a second vessel is placed inside this
       vessel, then another gas may be easily liquefied in it at the low
       temperature produced by the boiling liquid sulphurous anhydride.
       Pictet in this manner easily liquefied carbonic anhydride, CO_{2}
       (at -60° under a pressure of from four to six atmospheres).
       This gas is more refractory to liquefaction than sulphurous
       anhydride, but for this reason it gives on evaporating a still
       lower temperature than can be attained by the evaporation of
       sulphurous anhydride. A temperature of -80° may be obtained by
       the evaporation of liquid carbonic anhydride at a pressure of
       760 mm., and in an atmosphere rarefied by a powerful pump the
       temperature falls to -140°. By employing such low temperatures, it
       was possible, with the aid of pressure, to liquefy the majority
       of the other gases. It is evident that special pumps which are
       capable of rarefying gases are necessary to reduce the pressure
       in the chambers in which the sulphurous and carbonic anhydride
       boil; and that, in order to re-condense the resultant gases into
       liquids, special force pumps are required for pumping the liquid
       anhydrides into the refrigerating chamber. Thus, in Pictet's
       apparatus (fig. 24), the carbonic anhydride was liquefied by the
       aid of the pumps E F, which compressed the gas (at a pressure of
       4-6 atmospheres) and forced it into the tube K, vigorously cooled
       by being surrounded by boiling liquid sulphurous anhydride, which
       was condensed in the tube C by the pump B, and rarefied by the
       pump A. The liquefied carbonic anhydride flowed down the tube K
       into the tube H, in which it was subjected to a low pressure by
       the pump E, and thus gave a very low temperature of about -140°.
       The pump E carried off the vapour of the carbonic anhydride, and
       conducted it to the pump F, by which it was again liquefied.
       The carbonic anhydride thus made an entire circuit--that
       is, it passed from a rarefied vapour of small tension and
       low temperature into a compressed and cooled gas, which was
       transformed into a liquid, which again vaporised and produced a
       low temperature.

       [Illustration: FIG. 24.--General arrangement of the apparatus
       employed by Pictet for liquefying gases.]

       Inside the wide inclined tube H, where the carbonic acid
       evaporated, was placed a second and narrow tube M containing
       hydrogen, which was generated in the vessel L from a mixture of
       sodium formate and caustic soda (CHO_{2}Na + NaHO = Na_{2}CO_{3}
       + H_{2}). This mixture gives hydrogen on heating the vessel L.
       This vessel and the tube M were made of thick copper, and could
       withstand great pressures. They were, moreover, hermetically
       connected together and closed up. Thus the hydrogen which was
       evolved had no outlet, accumulated in a limited space, and its
       pressure increased in proportion to the amount of it evolved.
       This pressure was recorded on a metallic manometer R attached to
       the end of the tube M. As the hydrogen in this tube was submitted
       to a very low temperature and a powerful pressure, all the
       necessary conditions were present for its liquefaction. When the
       pressure in the tube H became steady--_i.e._ when the temperature
       had fallen to -140° and the manometer R indicated a pressure
       of 650 atmospheres in the tube M--then this pressure did not
       rise with a further evolution of hydrogen in the vessel L. This
       served as an indication that the tension of the vapour of the
       hydrogen had attained a maximum corresponding with -140°, and that
       consequently all the excess of the gas was condensed to a liquid.
       Pictet convinced himself of this by opening the cock N, when the
       liquid hydrogen rushed out from the orifice. But, on leaving a
       space where the pressure was equal to 650 atmospheres, and coming
       into contact with air under the ordinary pressure, the liquid or
       powerfully compressed hydrogen expanded, began to boil, absorbed
       still more heat, and became still colder. In doing so a portion
       of the liquid hydrogen, according to Pictet, passed into a solid
       state, and did not fall in drops into a vessel placed under the
       outlet N, but as pieces of solid matter, which struck against
       the sides of the vessel like shot and immediately vaporised.
       Thus, although it was impossible to see and keep the liquefied
       hydrogen, still it was clear that it passed not only into a
       liquid, but also into a solid state. Pictet in his experiments
       obtained other gases which had not previously been liquefied,
       especially oxygen and nitrogen, in a liquid and solid state.
       Pictet supposed that liquid and solid hydrogen has the properties
       of a metal, like iron.

  [32] At the same time (1879) as Pictet was working on the liquefaction
       of gases in Switzerland, Cailletet, in Paris, was occupied on
       the same subject, and his results, although not so convincing as
       Pictet's, still showed that the majority of gases, previously
       unliquefied, were capable of passing into a liquid state.
       Cailletet subjected gases to a pressure of several hundred
       atmospheres in narrow thick-walled glass tubes (fig. 25); he
       then cooled the compressed gas as far as possible by surrounding
       it with a freezing mixture; a cock was then rapidly opened
       for the outlet of mercury from the tube containing the gas,
       which consequently rapidly and vigorously expanded. This rapid
       expansion of the gas would produce great cold, just as the
       rapid compression of a gas evolves heat and causes a rise in
       temperature. This cold was produced at the expense of the gas
       itself, for in rapidly expanding its particles were not able to
       absorb heat from the walls of the tube, and in cooling a portion
       of the expanding gas was transformed into liquid. This was seen
       from the formation of cloud-like drops like a fog which rendered
       the gas opaque. Thus Cailletet proved the possibility of the
       liquefaction of gases, but he did not isolate the liquids. The
       method of Cailletet allows the passage of gases into liquids
       being observed with greater facility and simplicity than Pictet's
       method, which requires a very complicated and expensive apparatus.

       [Illustration: FIG. 25.--Cailletet's apparatus for liquefying
       gases.]

       The methods of Pictet and Cailletet were afterwards improved by
       Olszewski, Wroblewski, Dewar, and others. In order to obtain a
       still lower temperature they employed, instead of carbonic acid
       gas, liquid ethylene or nitrogen and oxygen, whose evaporation
       at low pressures produces a much lower temperature (to -200°).
       They also improved on the methods of determining such low
       temperatures, but the methods were not essentially altered; they
       obtained nitrogen and oxygen in a liquid, and nitrogen even in a
       solid, state, but no one has yet succeeded in seeing hydrogen in
       a liquid form.

       The most illustrative and instructive results (because they
       gave the possibility of maintaining a very low temperature
       and the liquefied gas, even air, for a length of time) were
       obtained in recent years by Prof. Dewar in the Royal Institution
       of London, which is glorified by the names of Davy, Faraday,
       and Tyndall. Dewar, with the aid of powerful pumps, obtained
       many kilograms of oxygen and air (the boiling point under the
       atmospheric pressure =-190°) in a liquid state and kept them in
       this state for a length of time by means of open glass vessels
       with double walls, having a vacuum between them, which prevented
       the rapid transference of heat, and so gave the possibility of
       maintaining very low temperatures inside the vessel for a long
       period of time. The liquefied oxygen or air can be poured from
       one vessel into another and used for any investigations. Thus
       in June 1894, Prof. Dewar showed that at the low temperature
       produced by liquid oxygen many substances become phosphorescent
       (become self-luminous; for instance, oxygen on passing into a
       vacuum) and fluoresce (emit light after being illuminated; for
       instance, paraffin, glue, &c.) much more powerfully than at the
       ordinary temperature; also that solids then greatly alter in
       their mechanical properties, &c. I had the opportunity (1894)
       at Prof. Dewar's of seeing many such experiments in which open
       vessels containing pounds of liquid oxygen were employed, and
       in following the progress made in researches conducted at low
       temperatures, it is my firm impression that the study of many
       phenomena at low temperatures should widen the horizon of natural
       science as much as the investigation of phenomena made at the
       highest temperatures attained in the voltaic arc.

  [33] The investigations of S. Wroblewski in Cracow give reason to
       believe that Pictet could not have obtained liquid hydrogen in
       the interior of his apparatus, and that if he did obtain it,
       it could only have been at the moment of its outrush due to
       the fall in temperature following its sudden expansion. Pictet
       calculated that he obtained a temperature of -140°, but in reality
       it hardly fell below -120°, judging from the latest data for
       the vaporisation of carbonic anhydride under low pressure. The
       difference lies in the method of determining low temperatures.
       Judging from other properties of hydrogen (_see_ Note 34), one
       would think that its absolute boiling point lies far below -120°,
       and even -140° (according to the calculation of Sarrau, on the
       basis of its compressibility, at -174°). But even at -200° (if the
       methods of determining such low temperatures be correct) hydrogen
       does not give a liquid even under a pressure of several hundred
       atmospheres. However, on expansion a fog is formed and a liquid
       state attained, but the liquid does not separate.

  [34] After the idea of the absolute temperature of ebullition (_tc_,
       Note 29) had been worked out (about 1870), and its connection
       with the deviations from Mariotte's law had become evident, and
       especially after the liquefaction of permanent gases, general
       attention was turned to the development of the fundamental
       conceptions of the gaseous and liquid states of matter. Some
       investigators directed their energies to the further study of
       vapours (for instance, Ramsay and Young), gases (Amagat), and
       liquids (Zaencheffsky, Nadeschdin, and others), especially to
       liquids near _tc_ and _pc_; others (Konovaloff and De Heen)
       endeavoured to discover the relation between liquids under
       ordinary conditions (removed from _tc_ and _pc_) and gases,
       whilst a third class of investigators (van der Waals, Clausius,
       and others), starting from the generally-accepted principles
       of the mechanical theory of heat and the kinetic theory of
       gases, and assuming in gases the existence of those forces which
       certainly act in liquids, deduced the connection between the
       properties of one and the other. It would be out of place in
       an elementary handbook like the present to enunciate the whole
       mass of conclusions arrived at by this method, but it is well to
       give an idea of the results of van der Waals' considerations,
       for they explain the gradual uninterrupted passage from a liquid
       into a gaseous state in the simplest manner, and, although the
       deduction cannot be considered as complete and decisive (_see_
       Note 25), nevertheless it penetrates so deeply into the essence
       of the matter that its signification is not only reflected in a
       great number of physical investigations, but also in the province
       of chemistry, where instances of the passage of substances from
       a gaseous to a liquid state are so common, and where the very
       processes of dissociation, decomposition, and combination must be
       identified with a change of physical state of the participating
       substances, which has been elaborated by Gibbs, Lavenig, and
       others.

       For a _given quantity_ (weight, mass) _of a definite substance_,
       its state is expressed by three variables--volume _v_, pressure
       (elasticity, tension) _p_, and temperature _t_. Although the
       compressibility--[_i.e._, _d(v)_/_d(p)_]--of liquids is small,
       still it is clearly expressed, and varies not only with the
       nature of liquids but also with their pressure and temperature
       (at _tc_ the compressibility of liquids is very considerable).
       Although gases, according to Mariotte's law, with small
       variations of pressure, are uniformly compressed, nevertheless
       the dependence of their volume _v_ on _t_ and _p_ is very
       complex. This also applies to the coefficient of expansion [=
       _d(v)_/_d(t)_, or _d(p)_/_d(t)_], which also varies with _t_ and
       _p_, both for gases (_see_ Note 26), and for liquids (at _tc_ it
       is very considerable, and often exceeds that of gases, 0·00367).
       Hence, the _equation of condition_ must include three variables,
       _v_, _p_, and _t_. For a so-called perfect (ideal) gas, or for
       inconsiderable variations of density, the elementary expression
       _pv_ = _R_[Greek: a](1 + [Greek: a]_t_), or _pv_ = _R_(273 + _t_)
       should be accepted, where _R_ is a constant varying with the mass
       and nature of a gas, as expressing this dependence, because it
       includes in itself the laws of Gay-Lussac and Mariotte, for at a
       constant pressure the volume varies proportionally to
       1 + [Greek: a]_t_, and when _t_ is constant the product of _tv_ is
       constant. In its simplest form the equation may be expressed thus:

                              _pv_ = _RT_;

       where _T_ denotes what is termed the absolute temperature, or the
       ordinary temperature + 273--that is, _T_ = _t_ + 273.

       Starting from the supposition of the existence of an attraction
       or internal pressure (expressed by _a_) proportional to the
       square of the density (or inversely proportional to the square of
       the volume), and of the existence of a real volume or diminished
       length of path (expressed by _b_) for each gaseous molecule, van
       der Waals gives for gases the following more complex equation of
       condition:--

              (_p_ + _a_/_v_^2)(_v_-_b_) = 1 + 0·00367_t_;

       if at 0° under a pressure _p_ = 1 (for example, under the
       atmospheric pressure), the volume (for instance, a litre) of
       a gas or vapour he taken as 1, and therefore _v_ and _b_ be
       expressed by the same units as _p_ and _a_. The deviations
       from both the laws of Mariotte and Gay-Lussac are expressed
       by the above equation. Thus, for hydrogen _a_ must be taken
       as infinitely small, and _b_ = 0·0009, judging by the data
       for 1,000 and 2,500 metres pressure (Note 28). For other
       permanent gases, for which (Note 28) I showed (about 1870) from
       Regnault's and Natterer's data, a decrement of _pv_, followed
       by an increment, which was confirmed (about 1880) by fresh
       determinations made by Amagat, this phenomena may be expressed
       in definite magnitudes of _a_ and _b_ (although van der Waals'
       formula is not applicable in the case of very small pressures)
       with sufficient accuracy for contemporary requirements. It
       is evident that van der Waals' formula can also express the
       difference of the coefficients of expansion of gases with a
       change of pressure, and according to the methods of determination
       (Note 26). Besides this, van der Waals' formula shows that at
       temperatures above 273(8_a_/27_b_-1) only one actual volume
       (gaseous) is possible, whilst at lower temperatures, by varying
       the pressure, three different volumes--liquid, gaseous, and
       partly liquid, partly saturated-vaporous--are possible. It is
       evident that the above temperature is the absolute boiling
       point--that is (_tc_) = 273(8_a_/27_b_-1). It is found under the
       condition that all three possible volumes (the three roots of
       van der Waals' cubic equation) are then similar and equal (_vc_
       = 3_b_). The pressure in this case (_pc_) = _a_/(27_b_^2). These
       ratios between the constants _a_ and _b_ and the conditions of
       _critical state_--_i.e._ (_tc_) and (_pc_)--give the possibility
       of determining the one magnitude from the other. Thus for ether
       (Note 29), (_tc_) = 193°, (_tp_) = 40, hence _a_ = 0·0307, _b_
       = 0·00533, and (_vc_) = 0·016. That mass of ether which at
       a pressure of one atmosphere at 0° occupies one volume--for
       instance, a litre--occupies, according to the above-mentioned
       condition, this critical volume. And as the density of the vapour
       of ether compared with hydrogen = 37, and a litre of hydrogen
       at 0° and under the atmospheric pressure weighs 0·0896 gram,
       then a litre of ether vapour weighs 3·32 grams; therefore, in a
       critical state (at 193° and 40 atmospheres) 3·32 grams occupy
       0·016 litre, or 16 c.c.; therefore 1 gram occupies a volume of
       about 5 c.c., and the weight of 1 c.c. of ether will then be
       0·21. According to the investigations of Ramsay and Young (1887),
       the critical volume of ether was approximately such at about the
       absolute boiling point, but the compressibility of the liquid is
       so great that the slightest change of pressure or temperature
       has a considerable effect on the volume. But the investigations
       of the above savants gave another indirect demonstration of the
       truth of van der Waals' equation. They also found for ether that
       the isochords, or the lines of equal volumes (if both _t_ and
       _p_ vary), are generally straight lines. Thus the volume of 10
       c.c. for 1 gram of ether corresponds with pressures (expressed
       in metres of mercury) equal to 0·135_t_-3·3 (for example, at
       180° the pressure = 21 metres, and at 280° it = 34·5 metres).
       The rectilinear form of the isochord (when _v_ = _a_ constant
       quantity) is a direct result of van der Waals' formula.

       When, in 1883, I demonstrated that the specific gravity of
       liquids decreases in proportion to the rise of temperature
       [S_{_t_} = S_{_0_}-K_t_ or S_{_t_} = S_{_0_}(1-K_t_)], or that
       the volumes increase in inverse proportion to the binomial
       1-K_t_, that is, V_{_t_} = V_{_0_}(1-K_t_)^{-1}, where K is the
       modulus of expansion, which varies with the nature of the liquid,
       then, in general, not only does a connection arise between
       gases and liquids with respect to a change of volume, but also
       it would appear possible, by applying van der Waals' formula,
       to judge, from the phenomena of the expansion of liquids, as
       to their transition into vapour, and to connect together all
       the principal properties of liquids, which up to this time had
       not been considered to be in direct dependence. Thus Thorpe and
       Rücker found that 2(_tc_) + 273 = 1/K, where K is the modulus
       of expansion in the above-mentioned formula. For example, the
       expansion of ether is expressed with sufficient accuracy from
       0° to 100° by the equation S_{_t_} = 0·736/(1-0·00154_t_), or
       V_{_t_} = 1/(1-0·00154_t_), where 0·00154 is the modulus of
       expansion, and therefore (_tc_) = 188°, or by direct observation
       193°. For silicon tetrachloride, SiCl_{4}, the modulus equals
       0·00136, from whence (_tc_) = 231°, and by experiment 230°. On
       the other hand, D. P. Konovaloff, admitting that the external
       pressure _p_ in liquids is insignificant when compared with the
       internal (_a_ in van der Waals' formula), and that the work in
       the expansion of liquids is proportional to their temperature
       (as in gases), directly deduced, from van der Waals' formula,
       the above-mentioned formula for the expansion of liquids, V_{t}
       = 1/(1-K_t_), and also the magnitude of the latent heat of
       evaporation, cohesion, and compressibility under pressure. In
       this way van der Waals' formula embraces the gaseous, critical,
       and _liquid states_ of substances, and shows the connection
       between them. On this account, although van der Waals' formula
       cannot be considered as perfectly general and accurate, yet it is
       not only very much more exact than _pv_ = _RT_, but it is also
       more comprehensive, because it applies both to gases and liquids.
       Further research will naturally give a closer proximity to
       truth, and will show the connection between composition and the
       constants (_a_ and _b_); but a great scientific progress is seen
       in this form of the equation of state.

       Clausius (in 1880), taking into consideration the variability of
       _a_, in van der Waals' formula, with the temperature, gave the
       following equation of condition:--

            (_p_ + _a_/(_T_(_v_ + _c_)^2)) (_v_-_b_) = _RT_.

       Sarrau applied this formula to Amagat's data for hydrogen, and
       found _a_ = 0·0551, _c_ =-0·00043, _b_ = 0·00089, and therefore
       calculated its absolute boiling point as -174°, and (_pc_) = 99
       atmospheres. But as similar calculations for oxygen (-105°),
       nitrogen (-124°), and marsh gas (-76°) gave _tc_ higher than
       it really is, the absolute boiling point of hydrogen must lie
       below -174°.

Although a substance which passes with great difficulty into a liquid
state by the action of physico-mechanical forces, hydrogen loses its
gaseous state (that is, its elasticity, or the physical energy of its
molecules, or their rapid progressive motion) with comparative ease
under the influence of chemical attraction,[35] which is not only shown
from the fact that hydrogen and oxygen (two permanent gases) form liquid
water, but also from many phenomena of the absorption of hydrogen.

  [35] This and a number of similar cases clearly show how great are the
       internal chemical forces compared with physical and mechanical
       forces.

Hydrogen is vigorously condensed by certain solids; for example, by
charcoal and by spongy platinum. If a piece of freshly ignited charcoal
be introduced into a cylinder full of hydrogen standing in a mercury
bath, then the charcoal absorbs as much as twice its volume of hydrogen.
Spongy platinum condenses still more hydrogen. But _palladium_, a grey
metal which occurs with platinum, absorbs more hydrogen than any other
metal. Graham showed that when heated to a red heat and cooled in an
atmosphere of hydrogen, palladium retains as much as 600 volumes of
hydrogen. When once absorbed it retains the hydrogen at the ordinary
temperature, and only parts with it when heated to a red heat.[36] This
capacity of certain dense metals for the absorption of hydrogen explains
the property of hydrogen of passing through metallic tubes.[37] It is
termed _occlusion_, and presents a similar phenomenon to solution; it is
based on the capacity of metals of forming unstable easily dissociating
compounds[38] with hydrogen, similar to those which salts form with
water.

  [36] The property of palladium of absorbing hydrogen, and of increasing
       in volume in so doing, may be easily demonstrated by taking a
       sheet of palladium varnished on one side, and using it as a
       cathode. The hydrogen which is evolved by the action of the
       current is retained by the unvarnished surface, as a consequence
       of which the sheet curls up. By attaching a pointer (for
       instance, a quill) to the end of the sheet this bending effect is
       rendered strikingly evident, and on reversing the current (when
       oxygen will be evolved and combine with the absorbed hydrogen,
       forming water) it may be shown that on losing the hydrogen the
       palladium regains its original form.

  [37] Deville discovered that iron and platinum become pervious to
       hydrogen at a red heat. He speaks of this in the following
       terms:--'The permeability of such homogeneous substances as
       platinum and iron is quite different from the passage of gases
       through such non-compact substances as clay and graphite. The
       permeability of metals depends on their expansion, brought
       about by heat, and proves that metals and alloys have a certain
       porosity.' However, Graham proved that it is only hydrogen which
       is capable of passing through the above-named metals in this
       manner. Oxygen, nitrogen, ammonia, and many other gases, only
       pass through in extremely minute quantities. Graham showed that
       at a red heat about 500 c.c. of hydrogen pass per minute through
       a surface of one square metre of platinum 1·1 mm. thick, but that
       with other gases the amount transmitted is hardly perceptible.
       Indiarubber has the same capacity for allowing the transference
       of hydrogen through its substance (_see_ Chapter III.), but at
       the ordinary temperature one square metre, 0·014 mm. thick,
       transmits only 127 c.c. of hydrogen per minute. In the experiment
       on the decomposition of water by heat in porous tubes, the clay
       tube may be exchanged for a platinum one with advantage. Graham
       showed that by placing a platinum tube containing hydrogen under
       these conditions, and surrounding it by a tube containing air,
       the transference of the hydrogen may be observed by the decrease
       of pressure in the platinum tube. In one hour almost all the
       hydrogen (97 p.c.) had passed from the tube, without being
       replaced by air. It is evident that the occlusion and passage
       of hydrogen through metals capable of occluding it are not only
       intimately connected together, but are dependent on the capacity
       of metals to form compounds of various degrees of stability with
       hydrogen--like salts with water.

  [38] It appeared on further investigation that palladium gives a
       definite compound, Pd_{2}H (_see_ further) with hydrogen; but
       what was most instructive was the investigation of sodium
       hydride, Na_{2}H, which clearly showed that the origin and
       properties of such compounds are in entire accordance with the
       conceptions of dissociation.

       Since hydrogen is a gas which is difficult to condense, it is
       little soluble in water and other liquids. At 0° a hundred
       volumes of water dissolve 1·9 volume of hydrogen, and alcohol 6·9
       volumes measured at 0° and 760 mm. Molten iron absorbs hydrogen,
       but in solidifying, it expels it. The solution of hydrogen by
       metals is to a certain degree based on its affinity for metals,
       and must be likened to the solution of metals in mercury and to
       the formation of alloys. In its chemical properties hydrogen,
       as we shall see later, has much of a metallic character. Pictet
       (_see_ Note 31) even affirms that liquid hydrogen has metallic
       properties. The metallic properties of hydrogen are also evinced
       in the fact that it is a good conductor of heat, which is not the
       case with other gases (Magnus).

At the ordinary temperature hydrogen very feebly and rarely enters into
chemical reaction. The capacity of gaseous hydrogen for reaction becomes
evident only under a change of circumstances--by compression, heating,
or the action of light, or at the moment of its evolution. However,
under these circumstances it _combines_ directly with only a very few of
the elements. Hydrogen combines directly with oxygen, sulphur, carbon,
potassium, and certain other elements, but it does not combine directly
with either the majority of the metals or with nitrogen, phosphorus, &c.
Compounds of hydrogen with certain elements on which it does not act
directly are, however, known; they are not obtained by a direct method,
but by reactions of decomposition, or of double decomposition, of other
hydrogen compounds. The property of hydrogen of combining with oxygen
at a red heat determines its combustibility. We have already seen that
hydrogen easily takes fire, and that it then burns with a pale--that is,
non-luminous--flame.[39] Hydrogen does not combine with the oxygen of
the atmosphere at the ordinary temperature; but this combination takes
place at a red heat,[40] and is accompanied by the evolution of much
heat. The product of this combination is water--that is, a compound
of oxygen and hydrogen. This is the _synthesis of water_, and we have
already noticed its analysis or decomposition into its component parts.
The synthesis of water may be very easily observed if a cold glass bell
jar be placed over a burning hydrogen flame, and, better still, if the
hydrogen flame be lighted in the tube of a condenser. The water will
condense in drops as it is formed on the walls of the condenser and
trickle down.[41]

  [39] If it be desired to obtain a perfectly colourless hydrogen
       flame, it must issue from a platinum nozzle, as the glass end of
       a gas-conducting tube imparts a yellow tint to the flame, owing
       to the presence of sodium in the glass.

  [40] Let us imagine that a stream of hydrogen passes along a tube, and
       let us mentally divide this stream into several parts,
       consecutively passing out from the orifice of the tube. The first
       part is lighted--that is, brought to a state of incandescence,
       in which state it combines with the oxygen of the atmosphere. A
       considerable amount of heat is evolved in the combination. The
       heat evolved then, so to say, ignites the second part of hydrogen
       coming from the tube, and, therefore, when once ignited, the
       hydrogen continues to burn, if there be a continual supply of it,
       and if the atmosphere in which it burns be unlimited and contains
       oxygen.

  [41] The combustibility of hydrogen may be shown by the direct
       decomposition of water by sodium. If a pellet of sodium be thrown
       into a vessel containing water, it floats on the water and
       evolves hydrogen, which may be lighted. The presence of sodium
       imparts a yellow tint to the flame. If potassium be taken, the
       hydrogen bursts into flame spontaneously, because sufficient heat
       is evolved in the reaction to ignite the hydrogen. The flame is
        violet by the potassium. If sodium be thrown not on
       to water, but on to an acid, it will evolve more heat, and the
       hydrogen will then also burst into flame. These experiments must
       be carried on with caution, as, sometimes towards the end, a
       mass of sodium oxide (Note 8) is produced, and flies about; it
       is therefore best to cover the vessel in which the experiment is
       carried on.

Light does not aid the combination of hydrogen and oxygen, so that a
mixture of these two gases does not change when exposed to the action
of light; but an electric spark acts just like a flame, and this is
taken advantage of for inflaming a mixture of oxygen and hydrogen, or
detonating gas, inside a vessel, as will be explained in the following
chapters. As hydrogen (and oxygen also) is condensed by spongy platinum,
by which a rise of temperature ensues, and as platinum acts by contact
(Introduction), therefore hydrogen also combines with oxygen, under
the influence of platinum, as Döbereiner showed. If spongy platinum be
thrown into a mixture of hydrogen and oxygen, an explosion takes place.
If a mixture of the gases be passed over spongy platinum, combination
also ensues, and the platinum becomes red-hot.[42]

  [42] This property of spongy platinum is made use of in the so-called
       hydrogen cigar-lighter. It consists of a glass cylinder or
       beaker, inside which there is a small lead stand (which is not
       acted on by sulphuric acid), on which a piece of zinc is laid.
       This zinc is covered by a bell, which is open at the bottom
       and furnished with a cock at the top. Sulphuric acid is poured
       into the space between the bell and the sides of the outer
       glass cylinder, and will thus compress the gas in the bell. If
       the cock of the cylinder be opened the gas will escape by it,
       and will be replaced by the acid, which, coming into contact
       with the zinc, evolves hydrogen, and it will escape through the
       cock. If the cock be closed, then the hydrogen evolved will
       increase the pressure of the gas in the bell, and thus again
       force the acid into the space between the bell and the walls of
       the outer cylinder. Thus the action of the acid on the zinc may
       be stopped or started at will by opening or shutting the cock,
       and consequently a stream of hydrogen may be always turned on.
       Now, if a piece of spongy platinum be placed in this stream, the
       hydrogen will take light, because the spongy platinum becomes
       hot in condensing the hydrogen and inflames it. The considerable
       rise in temperature of the platinum depends, among other things,
       on the fact that the hydrogen condensed in its pores comes into
       contact with previously absorbed and condensed atmospheric
       oxygen, with which hydrogen combines with great facility in this
       form. In this manner the hydrogen cigar-lighter gives a stream of
       burning hydrogen when the cock is open. In order that it should
       work regularly it is necessary that the spongy platinum should be
       quite clean, and it is best enveloped in a thin sheet of platinum
       foil, which protects it from dust. In any case, after some time
       it will be necessary to clean the platinum, which may be easily
       done by boiling it in nitric acid, which does not dissolve the
       platinum, but clears it of all dirt. This imperfection has given
       rise to several other forms, in which an electric spark is made
       to pass before the orifice from which the hydrogen escapes. This
       is arranged in such a manner that the zinc of a galvanic element
       is immersed when the cock is turned, or a small coil giving a
       spark is put into circuit on turning the hydrogen on.

Although gaseous hydrogen does not act directly[43] on many substances,
yet in a _nascent state_ reaction often takes place. Thus, for
instance, water on which sodium amalgam is acting contains hydrogen in
a nascent state. The hydrogen is here evolved from a liquid, and at
the first moment of its formation must be in a condensed state.[44]
In this condition it is capable of reacting on substances on which it
does not act in a gaseous state.[44 bis] Reactions of substitution or
displacement of metals by hydrogen at the moment of its formation are
particularly numerous.[45]

  [43] Under conditions similar to those in which hydrogen combines
       with oxygen it is also capable of combining with chlorine. A
       mixture of hydrogen and chlorine explodes on the passage of an
       electric spark through it, or on contact with an incandescent
       substance, and also in the presence of spongy platinum; but,
       besides this, the action of light alone is enough to bring about
       the combination of hydrogen and chlorine. If a mixture of equal
       volumes of hydrogen and chlorine be exposed to the action of
       sunlight, complete combination rapidly ensues, accompanied by a
       report. Hydrogen does not combine directly with carbon, either at
       the ordinary temperature or by the action of heat and pressure.
       But if an electric current be passed through carbon electrodes
       at a short distance from each other (as in the electric light
       or voltaic arc), so as to form an electric arc in which the
       particles of carbon are carried from one pole to the other, then,
       in the intense heat to which the carbon is subjected in this
       case, it is capable of combining with hydrogen. A gas of peculiar
       smell called acetylene, C_{2}H_{2}, is thus formed from carbon
       and hydrogen.

  [44] There is another explanation of the facility with which hydrogen
       reacts in a nascent state. We shall afterwards learn that the
       molecule of hydrogen contains two atoms, H_{2}, but there are
       elements the molecules of which only contain one atom--for
       instance, mercury. Therefore, every reaction of gaseous hydrogen
       must be accompanied by the disruption of that bond which exists
       between the atoms forming a molecule. At the moment of evolution,
       however, it is supposed that free atoms exist, and in this
       condition, according to the hypothesis, act energetically. This
       hypothesis is not based upon facts, and the idea that hydrogen is
       condensed at the moment of its evolution is more natural, and is
       in accordance with the fact (Note 12) that compressed hydrogen
       displaces palladium and silver (Brunner, Beketoff)--that is, acts
       as at the moment of its liberation.

  [44 bis] There is a very intimate and evident relation between the
       phenomena which take place in the action of spongy platinum and
       the phenomena of the action in a nascent state. The combination
       of hydrogen with aldehyde may be taken as an example. Aldehyde is
       a volatile liquid with an aromatic smell, boiling at 21°, soluble
       in water, and absorbing oxygen from the atmosphere, and in this
       absorption forming acetic acid--the substance which is found in
       ordinary vinegar. If sodium amalgam be thrown into an aqueous
       solution of aldehyde, the greater part of the hydrogen evolved
       combines with the aldehyde, forming alcohol--a substance also
       soluble in water, which forms the principle of all spirituous
       liquors, boils at 78°, and contains the same amount of oxygen
       and carbon as aldehyde, but more hydrogen. The composition of
       aldehyde is C_{2}H_{4}O, that of alcohol C_{2}H_{6}O.

  [45] When, for instance, an acid and zinc are added to a salt of
       silver, the silver is reduced; but this may be explained as a
       reaction of the zinc, and not of the hydrogen at the moment
       of its formation. There are, however, examples to which this
       explanation is entirely inapplicable; thus, for instance,
       hydrogen, at the moment of its liberation easily takes up oxygen
       from its compounds with nitrogen if they be in solution, and
       converts the nitrogen into its hydrogen-compound. Here the
       nitrogen and hydrogen, so to speak, meet at the moment of their
       liberation, and in this state combine together.

       It is evident from this that the elastic gaseous state of
       hydrogen fixes the limit of its energy: prevents it from entering
       into those combinations of which it is capable. In the nascent
       state we have hydrogen which is not in a gaseous state, and its
       action is then much more energetic. At the moment of evolution
       that heat, which would be latent in the gaseous hydrogen, is
       transmitted to its molecules, and consequently they are in a
       state of strain, and can hence act on many substances.

Metals, as we shall afterwards see, are in many cases able to replace
each other; they also, and in some cases still more easily, replace and
are replaced by hydrogen. We have already seen examples of this in the
formation of hydrogen from water, sulphuric acid, &c. In all these
cases the metals sodium, iron, or zinc displace the hydrogen which
occurs in these compounds. Hydrogen may be displaced from many of its
compounds by metals in exactly the same manner as it is displaced from
water; so, for example, hydrochloric acid, which is formed directly
by the combination of hydrogen with chlorine, gives hydrogen by the
action of a great many metals, just as sulphuric acid does. Potassium
and sodium also displace hydrogen from its compounds with nitrogen; it
is only from its compounds with carbon that hydrogen is not displaced
by metals. Hydrogen, in its turn, is able to replace metals; this is
accomplished most easily on heating, and with those metals which do not
themselves displace hydrogen. If hydrogen be passed over the compounds
of many metals with oxygen at a red heat, it takes up the oxygen from
the metals and displaces them just as it is itself displaced by metals.
If hydrogen be passed over the compound of oxygen with copper at a
red heat, then metallic copper and water are obtained--CuO + H_{2} =
H_{2}O + Cu. This kind of double decomposition is called _reduction_
with respect to the metal, which is thus reduced to a metallic state
from its combination with oxygen. But it must be recollected that all
metals do not displace hydrogen from its compound with oxygen, and,
conversely, hydrogen is not able to displace all metals from their
compounds with oxygen; thus it does not displace potassium, calcium, or
aluminium from its compounds with oxygen. If the metals be arranged in
the following series: K, Na, Ca, Al ... Fe, Zn, Hg ... Cu, Pb, Ag, Au,
then the first are able to take up oxygen from water--that is, displace
hydrogen--whilst the last do not act thus, but are, on the contrary,
reduced by hydrogen--that is, have, as is said, a less affinity for
oxygen than hydrogen, whilst potassium, sodium, and calcium have more.
This is also expressed by the amount of heat evolved in the act of
combination with oxygen (_see_ Note 7), and is shown by the fact that
potassium and sodium and other similar metals evolve heat in decomposing
water; but copper, silver, and the like do not do this, because in
combining with oxygen they evolve less heat than hydrogen does, and
therefore it happens that when hydrogen reduces these metals heat is
evolved. Thus, for example, if 16 grams of oxygen combine with copper,
38,000 units of heat are evolved; and when 16 grams of oxygen combine
with hydrogen, forming water, 69,000 units of heat are evolved; whilst
23 grams of sodium, in combining with 16 grams of oxygen, evolve 100,000
units of heat. This example clearly shows that chemical reactions which
proceed directly and unaided evolve heat. Sodium decomposes water and
hydrogen reduces copper, because they are _exothermal_ reactions, or
those which evolve heat; copper does not decompose water, because such
a reaction would be accompanied by an absorption (or secretion) of heat,
or belongs to the class of _endothermal_ reactions in which heat is
absorbed; and such reactions do not generally proceed directly, although
they may take place with the aid of energy (electrical, thermal, &c.)
borrowed from some foreign source.[46]

  [46] Several numerical data and reflections bearing on this matter are
       enumerated in Notes 7, 9, and 11. It must be observed that the
       action of iron or zinc on water is reversible. But the reaction
       CuO + H_{2} = Cu + H_{2}O is not reversible; the difference
       between the degrees of affinity is very great in this case,
       and, therefore, so far as is at present known, no hydrogen is
       liberated even in the presence of a large excess of water. It
       is to be further remarked, that under the conditions of the
       dissociation of water, copper is not oxidised by water, because
       the oxide of copper is reduced by free hydrogen. If a definite
       amount of a metal and acid be taken and their reaction be carried
       on in a closed space, then the evolution of hydrogen will cease,
       when its tension equals that at which compressed hydrogen
       displaces the metal. The result depends upon the nature of the
       metal and the strength of the solution of acid. Tammann and
       Nernst (1892) found that the metals stand in the following order
       in respect to this limiting tension of hydrogen:--Na, Mg, Zn, Al,
       Cd, Fe, Ni.

[Illustration: FIG. 26.--Apparatus employed by Dumas for determining the
composition of water. Described in text.]

The reduction of metals by hydrogen is taken advantage of for
_determining the exact composition of water by weight_. Copper oxide is
usually chosen for this purpose. It is heated to redness in hydrogen,
and the quantity of water thus formed is determined, when the quantity
of oxygen which occurs in it is found from the loss of weight of the
copper oxide. The copper oxide must be weighed immediately before and
after the experiment. The difference shows the weight of the oxygen
which entered into the composition of the water formed. In this manner
only solids have to be weighed, which is a very great gain in the
accuracy of the results obtained.[47] Dulong and Berzelius (1819) were
the first to determine the composition of water by this method, and
they found that water contains 88·91 of oxygen and 11·09 of hydrogen in
100 parts by weight, or 8·008 parts of oxygen per one part of hydrogen.
Dumas (1842) improved on this method,[48] and found that water contains
12·575 parts of hydrogen per 100 parts of oxygen--that is, 7·990 parts
of oxygen per 1 part of hydrogen--and therefore it is usually accepted
that _water contains eight parts by weight of oxygen to_ one _part by
weight of hydrogen_. By whatever method water be obtained, it will
always present the same composition. Whether it be taken from nature
and purified, or whether it be obtained from hydrogen by oxidation,
or whether it be separated from any of its compounds, or obtained by
some double decomposition--it will in every case contain one part by
weight of hydrogen and eight parts of oxygen. This is because water
is a definite chemical compound. Detonating gas, from which it may be
formed, is a simple mixture of oxygen and hydrogen, although a mixture
of the same composition as water. All the properties of both constituent
gases are preserved in detonating gas. Either one or the other gas may
be added to it without destroying its homogeneity. The fundamental
properties of oxygen and hydrogen are not found in water, and neither
of the gases can be directly combined with it. But they may be evolved
from it. In the formation of water there is an evolution of heat; for
the decomposition of water heat is required. All this is expressed by
the words, _Water is a definite chemical compound of hydrogen with
oxygen_. Taking the symbol of hydrogen, H, as expressing a unit quantity
by weight of this substance, and expressing 16 parts by weight of oxygen
by O, we can formulate all the above statements by the chemical symbol
of water, H_{2}O. As only definite chemical compounds are denoted
by formulæ, having denoted the formula of a compound substance we
express by it the entire series of properties which go to make up our
conception of a definite compound, and at the same time the quantitative
composition of the substance by weight. Further, as we shall afterwards
see, formulæ express the volume of the gases contained in a substance.
Thus the formula of water shows that it contains two volumes of hydrogen
and one volume of oxygen. Besides which, we shall learn that the formula
expresses the density of the vapour of a compound, and on this many
properties of substances depend, and, as we shall learn, determine the
quantities of the bodies entering into reactions. This vapour density we
shall find also determines the quantity of a substance entering into a
reaction. Thus the letters H_{2}O tell the chemist the entire history of
the substance. This is an international language, which endows chemistry
with a simplicity, clearness, stability, and trustworthiness founded on
the investigation of the laws of nature.

  [47] This determination may be carried on in an apparatus like that
       mentioned in Note 13 of Chapter I.

  [48] We will proceed to describe Dumas' method and results. For this
       determination pure and dry copper oxide is necessary. Dumas took
       a sufficient quantity of copper oxide for the formation of 50
       grams of water in each determination. As the oxide of copper
       was weighed before and after the experiment, and as the amount
       of oxygen contained in water was determined by the difference
       between these weights, it was essential that no other substance
       besides the oxygen forming the water should be evolved from
       the oxide of copper during its ignition in hydrogen. It was
       necessary, also, that the hydrogen should be perfectly pure,
       and free not only from traces of moisture, but from any other
       impurities which might dissolve in the water or combine with the
       copper and form some other compound with it. The bulb containing
       the oxide of copper (fig. 26), which was heated to redness,
       should be quite free from air, as otherwise the oxygen in the
       air might, in combining with the hydrogen passing through the
       vessel, form water in addition to that formed by the oxygen
       of the oxide of copper. The water formed should be entirely
       absorbed in order to accurately determine its quantity. The
       hydrogen was evolved in the three-necked bottle. The sulphuric
       acid, for acting on the zinc, is poured through funnels into
       the middle neck. The hydrogen evolved in the Woulfe's bottle
       passes through [U] tubes, in which it is purified, to the bulb,
       where it comes into contact with the copper oxide, forms water,
       and reduces the oxide to metallic copper; the water formed is
       condensed in the second bulb, and any passing off is absorbed in
       the second set of [U] tubes. This is the general arrangement of
       the apparatus. The bulb with the copper oxide is weighed before
       and after the experiment. The loss in weight shows the quantity
       of oxygen which entered into the composition of the water formed,
       the weight of the latter being shown by the gain in weight of
       the absorbing apparatus. Knowing the amount of oxygen in the
       water formed, we also know the quantity of hydrogen contained in
       it, and consequently we determine the composition of water by
       weight. This is the essence of the determination. We will now
       turn to certain particulars. In one neck of the three-necked
       bottle a tube is placed dipping under mercury. This serves as
       a safety-valve to prevent the pressure inside the apparatus
       becoming too great from the rapid evolution of hydrogen. If the
       pressure rose to any considerable extent, the current of gases
       and vapours would be very rapid, and, as a consequence, the
       hydrogen would not be perfectly purified, or the water entirely
       absorbed in the tubes placed for this purpose. In the third neck
       of the Woulfe's bottle is a tube conducting the hydrogen to the
       purifying apparatus, consisting of eight [U] tubes, destined
       for the purification and testing of the hydrogen. The hydrogen,
       evolved by zinc and sulphuric acid, is purified by passing it
       first through a tube full of pieces of glass moistened with a
       solution of lead nitrate next through silver sulphate; the lead
       nitrate retains sulphurette hydrogen, and arseniuretted hydrogen
       is retained by the tube with silver sulphate. Caustic potash in
       the next [U] tube retains any acid which might come over. The
       two following tubes are filled with lumps of dry caustic potash
       in order to absorb any carbonic anhydride and moisture which the
       hydrogen might contain. The next two tubes, to remove the last
       traces of moisture, are filled with phosphoric anhydride, mixed
       with lumps of pumice-stone. They are immersed in a freezing
       mixture. The small [U] tube contains hygroscopic substances,
       and is weighed before the experiment: this is in order to know
       whether the hydrogen passing through still retains any moisture.
       If it does not, then the weight of this tube will not vary during
       the whole experiment, but if the hydrogen evolved still retains
       moisture, the tube will increase in weight. The copper oxide is
       placed in the bulb, which, previous to the experiment, is dried
       with the copper oxide for a long period of time. The air is then
       exhausted from it, in order to weigh the oxide of copper in a
       vacuum and to avoid the need of a correction for weighing in
       air. The bulb is made of infusible glass, that it may be able to
       withstand a lengthy (20 hours) exposure to a red heat without
       changing in form. The weighed bulb is only connected with the
       purifying apparatus after the hydrogen has passed through for
       a long time, and after experiment has shown that the hydrogen
       passing from the purifying apparatus is pure and does not
       contain any air. On passing from the condensing bulb the gas and
       vapour enter into an apparatus for absorbing the last traces of
       moisture. The first [U] tube contains pieces of ignited potash,
       the second and third tubes phosphoric anhydride or pumice-stone
       moistened with sulphuric acid. The last of the two is employed
       for determining whether all the moisture is absorbed, and is
       therefore weighed separately. The final tube only serves as a
       safety-tube for the whole apparatus, in order that the external
       moisture should not penetrate into it. The glass cylinder
       contains sulphuric acid, through which the excess of hydrogen
       passes; it enables the rate at which the hydrogen is evolved
       to be judged, and whether its amount should be decreased or
       increased.

       When the apparatus is fitted up it must be seen that all its
       parts are hermetically tight before commencing the experiment.
       When the previously weighed parts are connected together and the
       whole apparatus put into communication, then the bulb containing
       the copper oxide is heated with a spirit lamp (reduction does
       not take place without the aid of heat), and the reduction of
       the copper oxide then takes place, and water is formed. When
       nearly all the copper oxide is reduced the lamp is removed and
       the apparatus allowed to cool, the current of hydrogen being
       kept up all the time. When cool, the drawn-out end of the bulb
       is fused up, and the hydrogen remaining in it is exhausted, in
       order that the copper may be again weighed in a vacuum. The
       absorbing apparatus remains full of hydrogen, and would therefore
       present a less weight than if it were full of air, as it was
       before the experiment, and for this reason, having disconnected
       the copper oxide bulb, a current of dry air is passed through it
       until the gas passing from the glass cylinder is quite free from
       hydrogen. The condensing bulb and the two tubes next to it are
       then weighed, in order to determine the quantity of water formed.
       Dumas repeated this experiment many times. The average result was
       that water contains 1253·3 parts of hydrogen per 10,000 parts
       of oxygen. Making a correction for the amount of air contained
       in the sulphuric acid employed for producing the hydrogen,
       Dumas obtained the average figure 1251·5, between the extremes
       1247·2 and 1256·2. This proves that per 1 part of hydrogen water
       contains 7·9904 parts of oxygen, with a possible error of not
       more than 1/250, or 0·03, in the amount of oxygen per 1 part of
       hydrogen.

       Erdmann and Marchand, in eight determinations, found that per
       10,000 parts of oxygen water contains an average of 1,252 parts
       of hydrogen, with a difference of from 1,258·5 to 1,248·7; hence
       per 1 part of hydrogen there would be 7·9952 of oxygen, with an
       error of at least 0·05.

       Keiser (1888), in America by employing palladium hydride, and
       by introducing various fresh precautions for obtaining accurate
       results, found the composition of water to be 15·95 parts of
       oxygen per 2 of hydrogen.

       Certain of the latest determinations of the composition of
       water, as also those made by Dumas, always give less than 8, and
       on the average 7·98, of oxygen per 1 part of hydrogen. However,
       not one of these figures is to be entirely depended on, and for
       ordinary accuracy it may be considered that O = 16 when H = 1.




                               CHAPTER III

         OXYGEN AND THE CHIEF ASPECTS OF ITS SALINE COMBINATIONS


On the earth's surface there is no other element which is so widely
distributed as oxygen in its various compounds.[1] It makes up
eight-ninths of the weight of water, which occupies the greater part
of the earth's surface. Nearly all earthy substances and rocks consist
of compounds of oxygen with metals and other elements. Thus, the
greater part of sand is formed of silica, SiO_{2}, which contains 53
p.c. of oxygen; clay contains water, alumina (formed of aluminium and
oxygen), and silica. It may be considered that earthy substances and
rocks contain up to one-third of their weight of oxygen; animal and
vegetable substances are also very rich in oxygen. Without counting
the water present in them, plants contain up to 40, and animals up to
20 p.c. by weight of oxygen. Thus, oxygen compounds predominate on the
earth's surface. Besides this, a portion exists in a free state, and is
contained in admixture with nitrogen in the atmosphere, forming about
one-fourth of its mass, or one-fifth of its volume.

  [1] As regards the interior of the earth, it probably contains far less
      oxygen compounds than the surface, judging by the accumulated
      evidences of the earth's origin, of meteorites, of the earth's
      density, &c. (_see_ Chapter VIII., Note 58, and Chapter XXII.,
      Note 2).

Being so widely distributed in nature, oxygen plays a very important
part in it, for a number of the phenomena which take place before us
are mainly dependent on it. _Animals breathe_ air in order to obtain
only _oxygen_ from it, the oxygen entering into their respiratory organs
(the lungs of human beings and animals, the gills of fishes, and the
trachæ of insects); they, so to say, drink in air in order to absorb the
oxygen. The oxygen of the air (or dissolved in water) passes through the
membranes of the respiratory organs into the blood, is retained in it by
the blood corpuscles, is transmitted by their means to all parts of the
body, aids their transformations, bringing about chemical processes in
them, and chiefly extracting carbon from them in the form of carbonic
anhydride, the greater part of which passes into the blood, is dissolved
by it, and is thrown off by the lungs during the absorption of the
oxygen. Thus, in the process of respiration carbonic anhydride (and
water) is given off, and the oxygen of the air absorbed, by which means
the blood is changed from a red venous to a dark-red arterial blood. The
cessation of this process causes death, because then all those chemical
processes, and the consequent heat and work which the oxygen introduced
into the system brought about, cease. For this reason suffocation and
death ensue in a vacuum, or in a gas which does not contain free oxygen,
_i.e._ which does not support combustion. If an animal be placed in an
atmosphere of free oxygen, at first its movements are very active and a
general invigoration is remarked, but a reaction soon sets in, and death
may ensue. The oxygen of the air when it enters the lungs is diluted
with four volumes of nitrogen, which is not absorbed into the system,
so that the blood absorbs but a small quantity of oxygen from the air,
whilst in an atmosphere of pure oxygen a large quantity of oxygen
would be absorbed, and would produce a very rapid change of all parts
of the organism, and destroy it. From what has been said, it will be
understood that oxygen may be employed in respiration, at any rate for a
limited time, when the respiratory organs suffer under certain forms of
suffocation and impediment to breathing.[2]

  [2] It is evident that the partial pressure (_see_ Chapter I.) acts in
      respiration. The researches of Paul Bert showed this with
      particular clearness. Under a pressure of one-fifth of an
      atmosphere consisting of oxygen only, animals and human beings
      remain under the ordinary conditions of the partial pressure
      of oxygen, but organisms cannot support air rarefied to
      one-fifth, for then the partial pressure of the oxygen falls
      to one-twenty-fifth of an atmosphere. Even under a pressure of
      one-third of an atmosphere the regular life of human beings
      is impossible, by reason of the impossibility of respiration
      (because of the decrease of solubility of oxygen in the blood),
      owing to the small partial pressure of the oxygen, and not from
      any mechanical effect of the decrease of pressure. Paul Bert
      illustrated all this by many experiments, some of which he
      conducted on himself. This explains, among other things, the
      discomfort felt in the ascent of high mountains or in balloons
      when the height reached exceeds eight kilometres, and at pressures
      below 250 mm. (Chapter II., Note 23). It is evident that an
      artificial atmosphere has to be employed in the ascent to great
      heights, just as in submarine work. The cure by compressed and
      rarefied air which is practised in certain illnesses is based
      partly on the mechanical action of the change of pressure, and
      partly on the alteration in the partial pressure of the respired
      oxygen.

The combustion of organic substances--that is, substances which make
up the composition of plants and animals--proceeds in the same manner
as the combustion of many inorganic substances, such as sulphur,
phosphorus, iron, &c., from the combination of these substances with
oxygen, as was described in the Introduction. The decomposition,
rotting, and similar transformations of substances, which proceed
around us, are also very often dependent on the action of the oxygen
of the air, and also reduce it from a free to a combined state. The
majority of the compounds of oxygen are, like water, very stable, and
do not give up their oxygen under the ordinary conditions of nature. As
these processes are taking place everywhere, it might be expected that
the amount of free oxygen in the atmosphere should decrease, and this
decrease should proceed somewhat rapidly. This is, in fact, observed
where combustion or respiration proceeds in a closed space. Animals
suffocate in a closed space because in consuming the oxygen the air
remains unfit for respiration. In the same manner combustion, after a
time, ceases in a closed space, which may be proved by a very simple
experiment. An ignited substance--for instance, a piece of burning
sulphur--has only to be placed in a glass flask, which is then closed
with a stout cork to prevent the access of the external air; combustion
will proceed for a certain time, so long as the flask contains any
free oxygen, but it will cease when the oxygen of the enclosed air has
combined with the sulphur. From what has been said, it is evident that
regularity of combustion or respiration requires a constant renewal of
air--that is, that the burning substance or respiring animal should
have access to a fresh supply of oxygen. This is attained in dwellings
by having many windows, outlets, and ventilators, and by the current of
air produced by fires and stoves. As regards the air over the entire
earth's surface its amount of oxygen hardly decreases, because in nature
there is a process going on which renews the supply of free oxygen.
_Plants_, or rather their leaves, during daytime,[3] under the influence
of light, absorb carbonic anhydride CO_{2}, and _evolve free oxygen_.
Thus the loss of oxygen which occurs in consequence of the respiration
of animals and of combustion is made good by plants. If a leaf be placed
in a bell jar containing water, and carbonic anhydride (because this
gas is absorbed and oxygen evolved from it by plants) be passed into
the bell, and the whole apparatus placed in sunlight, then oxygen will
accumulate in the bell jar. This experiment was first made by Priestley
at the end of the last century. Thus the life of plants on the earth not
only serves for the formation of food for animals, but also for keeping
up a constant percentage of oxygen in the atmosphere. In the long period
of the life of the earth an equilibrium has been attained between the
processes absorbing and evolving oxygen, by which a definite quantity of
free oxygen is preserved in the entire mass of the atmosphere.[4]

  [3] At night, without the action of light, without the absorption
      of that energy which is required for the decomposition of carbonic
      anhydride into free oxygen and carbon (which is retained by the
      plants) they breathe like animals, absorbing oxygen and evolving
      carbonic anhydride. This process also goes on side by side with
      the reverse process in the daytime, but it is then far feebler
      than that which gives oxygen.

  [4] The earth's surface is equal to about 510 million square
      kilometres, and the mass of the air (at a pressure of 760 mm.)
      on each kilometre of surface is about 10-1/3 thousand millions
      of kilograms, or about 10-1/3 million tons; therefore the whole
      weight of the atmosphere is about 5,100 million million (= 51 ×
      10^{14}) tons. Consequently there are about 2 × 10^{15} tons of
      free oxygen in the earth's atmosphere. The innumerable series of
      processes which absorb a portion of this oxygen are compensated
      for by the plant processes. Assuming that 100 million tons of
      vegetable matter, containing 40 p.c. of carbon, formed from
      carbonic acid, are produced (and the same process proceeds in
      water) per year on the 100 million square kilometres of dry land
      (ten tons of roots, leaves, stems, &c., per hectare, or 1/100 of
      a square kilometre), we find that the plant life of the dry land
      gives about 100,000 tons of oxygen, which is an insignificant
      fraction of the entire mass of the oxygen of the air.

Oxygen was obtained as an independent gas in 1774 by Priestley in
England and in the same year by Scheele in Sweden, but its nature and
great importance were only perfectly elucidated by Lavoisier.

Free oxygen may be obtained by one or other method from all the
substances in which it occurs. Thus, for instance, the oxygen of many
substances may be transferred into water, from which, as we have already
seen, oxygen may be obtained.[5] We will first consider the methods of
extracting oxygen from air as being a substance everywhere distributed.
The separation of oxygen from it is, however, hampered by many
difficulties.

  [5] The extraction of oxygen from water may be effected by two
      processes: either by the decomposition of water into its
      constituent parts by the action of a galvanic current (Chapter
      II.), or by means of the removal of the hydrogen from water.
      But, as we have seen and already know, hydrogen enters into
      direct combination with very few substances, and then only under
      special circumstances; whilst oxygen, as we shall soon learn,
      combines with nearly all substances. Only gaseous chlorine (and,
      especially, fluorine) is capable of decomposing water, taking up
      the hydrogen from it, without combining with the oxygen. Chlorine
      is soluble in water, and if an aqueous solution of chlorine,
      so-called chlorine water, be poured into a flask, and this flask
      be inverted in a basin containing the same chlorine water, then we
      shall have an apparatus by means of which oxygen may be extracted
      from water. At the ordinary temperature, and in the dark, chlorine
      does not act on water, or only acts very feebly; but under the
      action of direct sunlight chlorine decomposes water, with the
      evolution of oxygen. The chlorine then combines with the hydrogen,
      and gives hydrochloric acid, which dissolves in the water, and
      therefore free oxygen only will be separated from the liquid, and
      it will only contain a small quantity of chlorine in admixture,
      which can be easily removed by passing the gas through a solution
      of caustic potash.

From air, which contains a _mixture_ of oxygen and nitrogen, the
nitrogen alone cannot be removed, because it has no inclination to
combine directly or readily with any substance; and although it does
combine with certain substances (boron, titanium), these substances
combine simultaneously with the oxygen of the atmosphere.[6] However,
oxygen may be separated from air by causing it to combine with
substances which may be easily decomposed by the action of heat,
and, in so doing, give up the oxygen absorbed--that is, by making
use of reversible reactions. Thus, for instance, the oxygen of the
atmosphere may be made to oxidise sulphurous anhydride, SO_{2} (by
passing directly over ignited spongy platinum), and to form sulphuric
anhydride, or sulphur trioxide, SO_{3}; and this substance (which is a
solid and volatile, and therefore easily separated from the nitrogen and
sulphurous anhydride), on further heating, gives oxygen and sulphurous
anhydride. Caustic soda or lime extracts (absorbs) the sulphurous
anhydride from this mixture, whilst the oxygen is not absorbed, and thus
it is isolated from the air. On a large scale in works, as we shall
afterwards see, sulphurous anhydride is transformed into hydrate of
sulphuric trioxide, or sulphuric acid, H_{2}SO_{4}; if this is allowed
to drop on to red-hot flagstones, water, sulphurous anhydride, and
oxygen are obtained. The oxygen is easily isolated from this mixture
by passing the gases over lime. The extraction of oxygen from oxide
of mercury (Priestley, Lavoisier), which is obtained from mercury and
the oxygen of the atmosphere, is also a reversible reaction by which
oxygen may be obtained from the atmosphere. So also, by passing dry air
through a red-hot tube containing barium oxide, it is made to combine
with the oxygen of the air. In this reaction the so-called barium
peroxide, BaO_{2}, is formed from the barium oxide, BaO, and at a higher
temperature the former evolves the absorbed oxygen, and leaves the
barium oxide originally taken.[7]

  [6] A difference in the physical properties of both gases cannot be
      here taken advantage of, because they are very similar in this
      respect. Thus the density of oxygen is 16 times and of nitrogen
      14 times greater than the density of hydrogen, and therefore
      porous vessels cannot be here employed--the difference between
      the times of their passage through a porous surface would be too
      insignificant.

      [Illustration: FIG. 27.--Graham's apparatus for the decomposition
      of air by pumping it through india-rubber.]

      Graham, however, succeeded in enriching air in oxygen by passing
      it through india-rubber. This may be done in the following
      way:--A common india-rubber cushion, E (Fig. 27), is taken, and
      its orifice hermetically connected with an air-pump, or, better
      still, a mercury aspirator (the Sprengel pump is designated by
      the letters A, C, B). When the aspirator (Chapter II., Note 16)
      has pumped out the air, which will be seen by the mercury running
      out in an almost uninterrupted stream, and from its standing
      approximately at the barometric height, then it may be clearly
      observed that gas passes through the india-rubber. This is also
      seen from the fact that bubbles of gas continually pass along
      with the mercury. A minus pressure may be constantly maintained
      in the cushion by pouring mercury into the funnel A, and screwing
      up the pinchcock C, so that the stream flowing from it is small,
      and then a portion of the air passing through the india-rubber
      will be carried along with the mercury. This air may be collected
      in the cylinder, R. Its composition proves to be about 42 volumes
      of oxygen with 57 volumes of nitrogen, and one volume of carbonic
      anhydride, whilst ordinary air contains only 21 volumes of oxygen
      in 100 volumes. A square metre of india-rubber surface (of the
      usual thickness) passes about 45 c.c. of such air per hour. This
      experiment clearly shows that india-rubber is permeable to gases.
      This may, by the way, be observed in common toy balloons filled
      with coal-gas. They fall after a day or two, not because there
      are holes in them, but because air penetrates into, and the gas
      from, their interior, through the surface of the india-rubber of
      which they are made. The rate of the passage of gases through
      india-rubber does not, as Mitchell and Graham showed, depend
      on their densities, and consequently its permeability is not
      determined by orifices. It more resembles dialysis--that is, the
      penetration of liquids through colloid surfaces. Equal volumes
      of gases penetrate through india-rubber in periods of time which
      are related to each other as follows:--carbonic anhydride, 100;
      hydrogen, 247; oxygen, 532; marsh gas, 633; carbonic oxide,
      1,220; nitrogen, 1,358. Hence nitrogen penetrates more slowly
      than oxygen, and carbonic anhydride more quickly than other
      gases. 2·556 volumes of oxygen and 13·585 volumes of carbonic
      anhydride penetrate in the same time as one volume of nitrogen.
      By multiplying these ratios by the amounts of these gases in air,
      we obtain figures which are in almost the same proportion as the
      volumes of the gases penetrating from air through india-rubber. If
      the process of dialysis be repeated on the air which has already
      passed through india-rubber, then a mixture containing 65 p.c. by
      volume of oxygen is obtained. It may be thought that the cause
      of this phenomenon is the absorption or occlusion (_see_ Chap.
      II., Note 37) of gases by india-rubber and the evolution of the
      gas dissolved in a vacuum; and, indeed, india-rubber does absorb
      gases, especially carbonic anhydride. Graham called the above
      method of the decomposition of air _atmolysis_.

  [7] The preparation of oxygen by this method, which is due to
      Boussingault, is conducted in a porcelain tube, which is placed
      in a stove heated by charcoal, so that its ends project beyond
      the stove. Barium oxide (which may be obtained by igniting barium
      nitrate, previously dried) is placed in the tube, one end of
      which is connected with a pair of bellows, or a gas-holder, for
      keeping up a current of air through it. The air is previously
      passed through a solution of caustic potash, to remove all
      traces of carbonic anhydride, and it is very carefully dried
      (for the hydrate BaH_{2}O_{2} does not give the peroxide). At a
      _dark-red heat_ (500-600°) the oxide of barium absorbs oxygen
      from the air, so that the gas leaving the tube consists almost
      entirely of nitrogen. When the absorption ceases, the air will
      pass through the tube unchanged, which may be recognised from the
      fact that it supports combustion. The barium oxide is converted
      into peroxide under these circumstances, and eleven parts of
      barium oxide absorb about one part of oxygen by weight. When the
      absorption ceases, one end of the tube is closed, a cork with a
      gas-conducting tube is fixed into the other end, and the heat of
      the stove is increased to a _bright-red heat_ (800°). At this
      temperature the barium peroxide gives up all that oxygen which it
      acquired at a dark-red heat--_i.e._ about one part by weight of
      oxygen is evolved from twelve parts of barium peroxide. After the
      evolution of the oxygen there remains the barium oxide which was
      originally taken, so that air may be again passed over it, and
      thus the preparation of oxygen from one and the same quantity of
      barium oxide may be repeated many times. Oxygen has been produced
      one hundred times from one mass of oxide by this method; all the
      necessary precautions being taken, as regards the temperature of
      the mass and the removal of moisture and carbonic acid from the
      air. Unless these precautions be taken, the mass of oxide soon
      spoils.

      As oxygen may become of considerable technical use, from its
      capacity for giving high temperatures and intense light in
      the combustion of substances, its preparation directly from
      air by practical methods forms a problem whose solution many
      investigators continue to work at up to the present day. The
      most practical methods are those of Tessié du Motay and Kassner.
      The first is based on the fact that a mixture of equal weights
      of manganese peroxide and caustic soda at an incipient red
      heat (about 350°) absorbs oxygen from air, with the separation
      of water, according to the equation MnO_{2} + 2NaHO + O =
      Na_{2}MnO_{4} + H_{2}O. If superheated steam, at a temperature
      of about 450°, be then passed through the mixture, the manganese
      peroxide and caustic soda originally taken are regenerated, and
      the oxygen held by them is evolved, according to the reverse
      equation Na_{2}MnO_{4} + H_{2}O = MnO_{2} + 2NaHO + O. This mode
      of preparing oxygen may be repeated for an infinite number of
      times. The oxygen in combining liberates water, and steam, acting
      on the resultant substance, evolves oxygen. Hence all that is
      required for the preparation of oxygen by this method is fuel
      and the alternate cutting off the supply of air and steam. In
      Kassner's process (1891) a mixture of oxide of lead and lime (PbO
      + 2CaO) is heated to redness in the presence of air, oxygen is
      then absorbed and calcium plumbate, Ca_{2}PbO_{4}, formed. The
      latter is of a chocolate colour, and on further heating evolves
      oxygen and gives the original mixture PbO + 2CaO--that is, the
      phenomenon is essentially the same as in Boussingault's process
      (with BaO), but according to Le Chatelier (1893) the dissociation
      tension of the oxygen evolved from Ca_{2}PbO_{4} is less than with
      BaO_{2} at equal temperatures; for instance, at 940°, 112 mm. of
      mercury for the first, and for the latter 210 mm. at 720°, and 670
      mm. at 790°, while for Ca_{2}PbO_{4} this tension is only reached
      at 1,080°. However, in Kassner's process the oxygen is absorbed
      more rapidly, and the influence of the presence of moisture and
      CO_{2} in the air is not so marked, so that this process, like
      that of Tessié du Motay, deserves consideration.

Oxygen is evolved with particular ease by a whole series of unstable
oxygen compounds, of which we shall proceed to take a general survey,
remarking that many of these reactions, although not all, belong to
the number of reversible reactions;[8] so that in order to obtain many
of these substances (for instance, potassium chlorate) rich in oxygen,
recourse must be had to indirect methods (see Introduction) with which
we shall become acquainted in the course of this book.

  [8] Even the decomposition of manganese peroxide is reversible, and it
      may be re-obtained from that suboxide (or its salts), which
      is formed in the evolution of oxygen (Chap. XI., Note 6). The
      compounds of chromic acid containing the trioxide CrO_{3} in
      evolving oxygen give chromium oxide, Cr_{2}O_{3}, but they re-form
      the salt of chromic acid when heated to redness in air with an
      alkali.

1. _The compounds of oxygen_ with certain metals, and especially
with the so-called noble metals--that is, mercury, silver, gold, and
platinum--having once been obtained, retain their oxygen at the ordinary
temperature, but part with it at a red heat. The compounds are solids,
generally amorphous and infusible, and are easily decomposed by heat
into the metal and oxygen. We have seen an example of this in speaking
of the decomposition of mercury oxide. Priestley, in 1774, obtained
pure oxygen for the first time by heating mercury oxide by means of a
burning-glass, and clearly showed its difference from air. He showed
its characteristic property of supporting combustion 'with remarkable
vigour,' and named it dephlogisticated air.

2. The substances called _peroxides_[9] evolve oxygen at a greater or
less heat (and also by the action of many acids). They usually contain
metals combined with a large quantity of oxygen. Peroxides are the
highest oxides of certain metals; those metals which form them generally
give several compounds with oxygen. Those of the lowest degrees of
oxidation, containing the least amount of oxygen, are generally
substances which are capable of easily reacting with acids--for
instance, with sulphuric acid. Such low oxides are called bases.
Peroxides contain more oxygen than the bases formed by the same metals.
For example, lead oxide contains 7·1 parts of oxygen in 100 parts, and
is basic, but lead peroxide contains 13·3 parts of oxygen in 100 parts.
_Manganese peroxide_ is a similar substance, and is a solid of a dark
colour, which occurs in nature. It is employed for technical purposes
under the name of black oxide of manganese (in German, 'Braunstein,' the
pyrolusite of the mineralogist). Peroxides are able to evolve oxygen
at a more or less elevated temperature. They do not then part with all
their oxygen, but with only a portion of it, and are converted into
a lower oxide or base. Thus, for example, lead peroxide, on heating,
gives oxygen and lead oxide. The decomposition of this peroxide proceeds
tolerably easily on heating, even in a glass vessel, but manganese
peroxide only evolves oxygen at a strong red heat, and therefore oxygen
can only be obtained from it in iron, or other metallic, or clay
vessels. This was formerly the method for obtaining oxygen. Manganese
peroxide only parts with one-third of its oxygen (according to the
equation 3MnO_{2} = Mn_{3}O_{4} + O_{2}), whilst two-thirds remain in
the solid substance which forms the residue after heating. Metallic
peroxides are also capable of evolving oxygen on heating with sulphuric
acid. They then evolve just that amount of oxygen which is in excess of
that necessary for the formation of the base, the latter reacting on the
sulphuric acid forming a compound (salt) with it. Thus barium peroxide,
when heated with sulphuric acid, forms oxygen and barium oxide, which
gives a compound with sulphuric acid termed barium sulphate (BaO_{2}
+ H_{2}SO_{4} = BaSO_{4} + H_{2}O + O).[9 bis] This reaction usually
proceeds with greater ease than the decomposition of peroxides by heat
alone. For the purposes of experiment powdered manganese peroxide is
usually taken and mixed with strong sulphuric acid in a flask, and the
apparatus set up as shown in Fig. 28. The gas which is evolved is passed
through a Woulfe's bottle containing a solution of caustic potash, to
purify it from carbonic anhydride and chlorine, which accompany the
evolution of oxygen from commercial manganese peroxide, and the gas is
not collected until a thin smouldering taper placed in front of the
escape orifice bursts into flame, which shows that the gas coming off
is oxygen. By this method of decomposition of the manganese peroxide by
sulphuric acid there is evolved, not, as in heating, one-third, but
one-half of the oxygen contained in the peroxide (MnO_{2} + H_{2}SO_{4}
= MnSO_{4} + H_{2}O + O)--that is, from 50 grams of peroxide about
7-1/5 grams, or about 5-1/2 litres, of oxygen,[10] whilst by heating
only about 3-1/2 litres are obtained. The chemists of Lavoisier's time
generally obtained oxygen by heating manganese peroxide. At the present
time more convenient methods are known.

[Illustration: FIG. 28.--Preparation of oxygen from manganese peroxide
and sulphuric acid. The gas evolved is passed through a Woulfe's bottle
containing caustic potash.]

  [9] We shall afterwards see that it is only substances like barium
      peroxide (which give hydrogen peroxide) which should be counted
      as true peroxides, and that MnO_{2}, PbO_{2}, &c., should be
      distinguished from them (they do not give hydrogen peroxide with
      acids), and therefore it is best to call them dioxides.

  [9 bis] Peroxide of barium also gives oxygen at the ordinary
      temperature in the presence of the solutions of many substances
      in a higher degree of oxidation. In this respect we may mention
      that Kassner (1890) proposes to obtain oxygen for laboratory
      purposes by mixing BaO_{2} with FeK_{3}(CN)_{6} (red prussiate of
      potash, Chapter XXII.): the reaction proceeds with the evolution
      of oxygen even on the addition of a very small quantity of
      water. In order to ensure a gradual evolution of gas the author
      proposes to introduce both substances into the reaction, little
      by little, instead of all at once, which may be done with the
      following arrangement (Gavaloffsky): finely powdered peroxide
      of barium is placed in an ordinary flask and sufficient water
      is added to fill the flask one-third full. The cork closing the
      flask has three holes; (1) for the gas-conducting tube; (2) for a
      rod to stir the BaO_{2}; and (3) for a glass rod terminating in a
      perforated glass vessel containing crystals of FeK_{3}(CN)_{6}.
      When it is desired to start the evolution of the oxygen, the
      vessel is lowered until it is immersed in the liquid in the flask,
      and the BaO_{2} is stirred with the other rod. The reaction
      proceeds according to the equation, BaO_{2} + 2FeK_{3}(CN)_{6}
      = FeK_{4}(CN)_{6} + FeK_{2}Ba(CN)_{6} + O_{2}. The double salt,
      FeBa_{2}(CN)_{6}, crystallises out from the mother liquor. To
      understand the course of the reaction, it must be remembered
      BaO_{2} is of a higher degree of oxidation, and that it parts
      with oxygen and gives the base BaO which enters into the complex
      salt FeK_{2}Ba(CN)_{6} = Fe(CN)_{2} + 2KCN + Ba(CN)_{2}, and this
      latter = BaO + 2HCN-H_{2}O. Moreover, FeK_{3}(CN)_{6} contains the
      salt Fe_{2}(CN)_{6} which also corresponds to the higher degree
      of oxidation of iron, Fe_{2}O_{3}, whilst after the reaction a
      salt is obtained which contains Fe(CN)_{2}, and corresponds to the
      lower degree of oxidation, FeO, so that (in the presence of water)
      oxygen is also set free on this side also, _i.e._ the reaction
      gives lower degrees of oxidation and oxygen.

  [10] Scheele, in 1785, discovered the method of obtaining oxygen by
       treating manganese peroxide with sulphuric acid.

3. A third source to which recourse may be had for obtaining oxygen is
represented in _acids_ and _salts_ containing much oxygen, which are
capable, by parting with a portion or all of their oxygen, of being
converted into other compounds (lower products of oxidation) which are
more difficultly decomposed. These acids and salts (like peroxides)
evolve oxygen either on heating alone, or only when in the presence
of some other substance. Sulphuric acid may be taken as an example of
an acid which is decomposed by the action of heat alone,[11] for it
breaks up at a red heat into water, sulphurous anhydride, and oxygen,
as was mentioned before. Priestley, in 1772, and Scheele, somewhat
later, obtained oxygen by heating nitre to a red heat. The best
examples of the formation of oxygen by the heating of salts is given
in _potassium chlorate_, or Berthollet's salt, so called after the
French chemist who discovered it. Potassium chlorate is a salt composed
of the elements potassium, chlorine, and oxygen, KClO_{3}. It occurs
as transparent colourless plates, is soluble in water, especially in
hot water, and resembles common table salt in some of its reactions
and physical properties; it melts on heating, and in melting begins
to decompose, evolving oxygen gas. This decomposition ends in all the
oxygen being evolved from the potassium chlorate, potassium chloride
being left as a residue, according to the equation KClO_{3} = KCl +
O_{3}.[12] This decomposition proceeds at a temperature which allows
of its being conducted in a glass vessel. However, in decomposing,
the molten potassium chlorate swells up and boils, and gradually
solidifies, so the evolution of the oxygen is not regular, and the
glass vessel may crack. In order to overcome this inconvenience, the
potassium chlorate is crushed and mixed with a powder of a substance
which is infusible, incapable of combining with the oxygen evolved,
and is a good conductor of heat. Usually it is mixed with manganese
peroxide.[13] The decomposition of the potassium chlorate is then
considerably facilitated, and proceeds at a lower temperature (because
the entire mass is then better heated, both externally and internally),
without swelling up, and this method is therefore more convenient than
the decomposition of the salt alone. This method for the preparation
of oxygen is very convenient; it is generally employed when a small
quantity of oxygen is required. Further, potassium chlorate is easily
obtained pure, and it evolves much oxygen. 100 grams of the salt give as
much as 39 grams, or 30 litres, of oxygen. This method is so simple and
easy,[14] that a course of practical chemistry is often commenced by the
preparation of oxygen by this method, and of hydrogen by the aid of zinc
and sulphuric acid, since by means of these gases many interesting and
striking experiments may be performed.[15]

  [11] All acids rich in oxygen, and especially those whose elements form
       lower oxides, evolve oxygen either directly at the ordinary
       temperature (for instance, ferric acid), or on heating (nitric,
       manganic, chromic, chloric, and others), or if basic lower oxides
       are formed from them, by heating with sulphuric acid. Thus
       the salts of chromic acid (for example, potassium dichromate,
       K_{2}Cr_{2}O_{7}) give oxygen with sulphuric acid; first
       potassium sulphate, K_{2}SO_{4}, is formed, and then the chromic
       acid set free gives a sulphuric acid salt of the lower oxide,
       Cr_{2}O_{3}.

  [12] This reaction is not reversible, and is exothermal--that is, it
       does not absorb heat, but, on the contrary, evolves 9,713
       calories per molecular weight KClO_{3}, equal to 122 parts of
       salt (according to the determination of Thomsen, who burnt
       hydrogen in a calorimeter either alone or with a definite
       quantity of potassium chlorate mixed with oxide of iron). It does
       not proceed at once, but first forms perchlorate, KClO_{4} (_see_
       Chlorine and Potassium). It is to be remarked that potassium
       chloride melts at 766°, potassium chlorate at 359°, and potassium
       perchlorate at 610°. (Concerning the decomposition of KClO_{3},
       _see_ Chapter II., Note 47.)

  [13] The peroxide does not evolve oxygen in this case. It may be
       replaced by many oxides--for instance, by oxide of iron. It is
       necessary to take the precaution that no combustible substances
       (such as bits of paper, splinters, sulphur, &c.) fall into the
       mixture, as they might cause an explosion.

  [14] The decomposition of a mixture of fused and well-crushed potassium
       chlorate with powdered manganese peroxide proceeds at so low a
       temperature (the salt does not melt) that it may be effected in
       an ordinary glass flask. The apparatus is arranged in the same
       manner as in the decomposition of mercury oxide (Introduction),
       or as shown in the last drawing. As the reaction is exothermal,
       the decomposition of potassium chlorate with the formation of
       oxygen may probably be accomplished, under certain conditions
       (for example, under contact action), at very low temperatures.
       Substances mixed with the potassium chlorate probably act
       partially in this manner.

  [15] Many other salts evolve oxygen by heat, like potassium chlorate,
       but they only part with it either at a very high temperature (for
       instance, common nitre) or else are unsuited for use on account
       of their cost (potassium manganate), or evolve impure oxygen at a
       high temperature (zinc sulphate at a red heat gives a mixture of
       sulphurous anhydride and oxygen), and are not therefore used in
       practice.

A solution of _bleaching powder_, which contains calcium hypochlorite,
CaCl_{2}O_{2}, evolves oxygen on gently heating when a small quantity
of certain oxides is added--for instance, cobalt oxide, which in this
case acts by contact (_see_ Introduction). When heated by itself, a
solution of bleaching powder does not evolve oxygen, but it oxidises
the cobalt oxide to a higher degree of oxidation; this higher oxide of
cobalt in contact with the bleaching powder decomposes into oxygen and
lower oxidation products, and the resultant lower oxide of cobalt with
bleaching powder again gives the higher oxide, which again gives up
its oxygen, and so on.[16] The calcium hypochlorite is here decomposed
according to the equation CaCl_{2}O_{2} = CaCl_{2} + O_{2}. In this
manner a small quantity of cobalt oxide[17] is sufficient for the
decomposition of an indefinitely large quantity of bleaching powder.

  [16] Such is, at present, the only possible method of explaining the
       phenomenon of contact action. In many cases, such as the present
       one, it is supported by observations based on facts. Thus,
       for instance, it is known, as regards oxygen, that often two
       substances rich in oxygen retain it so long as they are separate,
       but directly they come into contact free oxygen is evolved from
       both of them. Thus, an aqueous solution of hydrogen peroxide
       (containing twice as much oxygen as water) acts in this manner on
       silver oxide (containing silver and oxygen). This reaction takes
       place at the ordinary temperature, and the oxygen is evolved
       from both compounds. To this class of phenomena may be also
       referred the fact that a mixture of barium peroxide and potassium
       manganate with water and sulphuric acid evolves oxygen at the
       ordinary temperature (Note 9 bis). It would seem that the essence
       of phenomena of this kind is entirely and purely a property of
       contact; the distribution of the atoms is changed by contact,
       and if the equilibrium be unstable it is destroyed. This is more
       especially evident in the case of those substances which change
       exothermally--that is, for those reactions which are accompanied
       by an evolution of heat. The decomposition CaCl_{2}O_{2} =
       CaCl_{2} + O_{2} belongs to this class (like the decomposition of
       potassium chlorate).

  [17] Generally a solution of bleaching powder is alkaline (contains
       free lime), and therefore, a solution of cobalt chloride is added
       directly to it, by which means the oxide of cobalt required for
       the reaction is formed.

_The properties of oxygen._[18]--It is a permanent _gas_--that is,
it cannot be liquefied by pressure at the ordinary temperature, and
further, is only liquefied with difficulty (although more easily than
hydrogen) at temperatures below -120°, because this is its absolute
boiling point. As its critical pressure[19] is about 50 atmospheres, it
can be easily liquefied under pressures greater than 50 atmospheres at
temperatures below -120°. According to Dewar, the density of oxygen in
a critical state is 0·65 (water = 1), but, like all other substances
in this state,[20] it varies considerably in density with a change of
pressure and temperature, and therefore many investigators who made
their observations under high pressures give a greater density, as much
as 1·1. Liquefied oxygen is an exceedingly mobile transparent liquid,
with a faint blue tint and boiling (tension = 1 atmosphere) about -180°.
Oxygen, like all gases, is transparent, and like the majority of gases,
colourless. It has no smell or taste, which is evident from the fact of
its being a component of air. The weight of one litre of oxygen gas at
0° and 760 mm. pressure is 1·4298 gram; it is therefore slightly denser
than air. Its density in respect to air = 1·1056 and in respect to
hydrogen = 16.[21]

  [18] It must be remarked that in all the reactions above mentioned the
       formation of oxygen may be prevented by the admixture of
       substances capable of combining with it--for example, charcoal,
       many carbon (organic) compounds, sulphur, phosphorus, and various
       lower oxidation products, &c. These substances absorb the oxygen
       evolved, combine with it, and a compound containing oxygen is
       formed. Thus, if a mixture of potassium chlorate and charcoal be
       heated, no oxygen is obtained, but an explosion takes place from
       the rapid formation of gases resulting from the combination of
       the oxygen of the potassium chlorate with the charcoal and the
       evolution of gaseous CO_{2}.

       The oxygen obtained by any of the above-described methods is
       rarely pure. It generally contains aqueous vapour, carbonic
       anhydride, and very often small traces of chlorine. The oxygen
       may be freed from these impurities by passing it through a
       solution of caustic potash, and by drying it. If the potassium
       chlorate be dry and pure, it gives almost pure oxygen. However,
       if the oxygen be required for respiration in cases of sickness,
       it should be washed by passing it through a solution of caustic
       alkali and through water. The best way to obtain pure oxygen
       directly is to take potassium perchlorate (KClO_{4}), which can
       be well purified and then evolves pure oxygen on heating.

  [19] With regard to the absolute boiling point, critical pressure, and
       the critical state in general, _see_ Chapter II., Notes 29 and 34.

  [20] Judging from what has been said in Note 34 of the last chapter,
       and also from the results of direct observation, it is evident
       that all substances in a critical state have a large coefficient
       of expansion, and are very compressible.

  [21] As water consists of 1 volume of oxygen and 2 volumes of hydrogen,
       and contains 16 parts by weight of oxygen per 2 parts by weight
       of hydrogen, it therefore follows directly that oxygen is 16
       times denser than hydrogen. Conversely, the composition of water
       by weight may be deduced from the densities of hydrogen and
       oxygen, and the volumetric composition of water. This method of
       mutual and reciprocal correction strengthens the practical data
       of the exact sciences, whose conclusions require the greatest
       possible exactitude and variety of corrections.

       It must he observed that the specific heat of oxygen at constant
       pressure is 0·2175, consequently it is to the specific heat of
       hydrogen (3·409) as 1 is to 15·6. Hence, the specific heats are
       inversely proportional to the weights of equal volumes. This
       signifies that equal volumes of both gases have (nearly) equal
       specific heats--that is, they require an equal quantity of heat
       for raising their temperature by 1°. We shall afterwards consider
       the specific heat of different substances more fully in Chap. XIV.

       Oxygen, like the majority of difficultly-liquefiable gases, is
       but slightly soluble in water and other liquids. The solubility
       is given in Note 30, Chap. I. From this it is evident that water
       standing in air must absorb--_i.e._ dissolve--oxygen. This oxygen
       serves for the respiration of fishes. Fishes cannot exist in
       boiled water, because it does not contain the oxygen necessary
       for their respiration (_see_ Chap. I.)

[Illustration: FIG. 29.--Mode of burning sulphur, phosphorus, sodium,
&c., in oxygen.]

In its chemical properties oxygen is remarkable from the fact that
it very easily--and, in a chemical sense, vigorously--reacts on a
number of substances, forming oxygen compounds. However, only a few
substances and mixtures of substances (for example, phosphorus, copper
with ammonia, decomposing organic matter, aldehyde, pyrogallol with an
alkali, &c.) combine directly with oxygen at the ordinary temperature,
whilst many substances easily combine with oxygen at a red heat, and
often this combination presents a rapid chemical reaction accompanied by
the evolution of a large quantity of heat. Every reaction which takes
place rapidly, if it be accompanied by so great an evolution of heat
as to produce incandescence, is termed _combustion_. Thus combustion
ensues when many metals are plunged into chlorine, or oxide of sodium
or barium into carbonic anhydride, or when a spark falls on gunpowder.
A great many substances are combustible in oxygen, and, owing to its
presence, in air also. In order to start combustion it is generally
necessary[22] that the combustible substance should be brought to a
state of incandescence. The continuation of the process does not require
the aid of fresh external heat, because sufficient heat[23] is evolved
to raise the temperature of the remaining parts of the combustible
substance to the required degree. Examples of this are familiar to all
from every-day experience. Combustion proceeds in oxygen with greater
rapidity, and is accompanied by a more powerful incandescence, than in
ordinary air. This may be demonstrated by a number of very convincing
experiments. If a piece of charcoal, attached to a wire and previously
brought to red-heat, be plunged into a flask full of oxygen, it burns
rapidly at a white heat--_i.e._ it combines with the oxygen, forming a
gaseous product of combustion called carbonic anhydride, or carbonic
acid gas, CO_{2}. This is the same gas that is evolved in the act of
respiration, for charcoal is one of the substances which is obtained by
the decomposition of all organic substances which contain it, and in
the process of respiration part of the constituents of the body, so to
speak, slowly burn. If a piece of burning sulphur be placed in a small
cup attached to a wire and introduced into a flask full of oxygen, then
the sulphur, which burns in air with a very feeble flame, burns in the
oxygen with a violet flame, which, although pale, is much larger than
in air. If the sulphur be exchanged for a piece of phosphorus,[24]
then, unless the phosphorus be heated, it combines very slowly with
the oxygen; but, if heated, although on only one spot, it burns with
an exceedingly brilliant white flame. In order to heat the phosphorus
inside the flask, the simplest way is to bring a red-hot wire into
contact with it. Before the charcoal can burn, it must be brought to
a state of incandescence. Sulphur also will not burn under 100°,
whilst phosphorus inflames at 40°. Phosphorus which has been already
lighted in air cannot so well be introduced into the flask, because
it burns very rapidly and with a large flame in air. If a small lump
of metallic _sodium_ be put in a small cup made of lime,[25] melted,
and ignited,[26] it burns very feebly in air. But if burning sodium be
introduced into oxygen, the combustion is invigorated and is accompanied
by a brighter yellow flame. Metallic _magnesium_, which burns brightly
in air, continues to burn with still greater vigour in oxygen, forming
a white powder, which is a compound of magnesium with oxygen (magnesium
oxide; magnesia). A strip of _iron_ or steel does not burn in air, but
an iron wire or steel spring may be easily burnt in oxygen.[27] The
combustion of steel or iron in oxygen is not accompanied by a flame, but
sparks of oxide fly in all directions from the burning portions of the
iron.[28]

  [22] Certain substances (with which we shall afterwards become
       acquainted), however, ignite spontaneously in air; for example,
       impure phosphuretted hydrogen, silicon hydride, zinc ethyl, and
       pyrophorus (very finely divided iron, &c.)

  [23] If so little heat is evolved that the adjacent parts are not
       heated to the temperature of combustion, then combustion will
       cease.

  [24] The phosphorus must be dry; it is usually kept in water, as
       it oxidises in air. It should be cut under water, as otherwise
       the freshly-cut surface oxidises. It must be dried carefully and
       quickly by wrapping it in blotting-paper. If damp, it splutters
       on burning. A small piece should be taken, as otherwise the iron
       spoon will melt. In this and the other experiments on combustion,
       water should be poured over the bottom of the vessel containing
       the oxygen, to prevent it from cracking. The cork closing
       the vessel should not fit tightly, in order to allow for the
       expansion of the gas due to the heat of the combustion.

  [25] An iron cup will melt with sodium in oxygen.

  [26] In order to rapidly heat the lime crucible containing the sodium,
       it is heated in the flame of a blowpipe described in Chap. VIII.

  [27] In order to burn a watch spring, a piece of tinder (or paper
       soaked in a solution of nitre, and dried) is attached to one
       end. The tinder is lighted, and the spring is then plunged into
       the oxygen. The burning tinder heats the end of the spring, the
       heated part burns, and in so doing heats the further portions of
       the spring, which then burns completely if sufficient oxygen be
       present.

  [28] The sparks of rust are produced, owing to the fact that the volume
       of the oxide of iron is nearly twice that of the volume of the
       iron, and as the heat evolved is not sufficient to entirely melt
       the oxide or the iron, the particles must be torn off and fly
       about. Similar sparks are formed in the combustion of iron, in
       other cases also. We saw the combustion of iron filings in the
       Introduction. In the welding of iron small iron splinters fly off
       in all directions and burn in the air, as is seen from the fact
       that whilst flying through the air they remain red hot, and also
       because, on cooling, they are seen to be no longer iron, but a
       compound of it with oxygen. The same thing takes place when the
       hammer of a gun strikes against the flint. Small scales of steel
       are heated by the friction, and glow and burn in the air. The
       combustion of iron is still better seen by taking it as a very
       fine powder, such as is obtained by the decomposition of certain
       of its compounds--for instance, by heating Prussian blue, or by
       the reduction of its compounds with oxygen by hydrogen; when this
       fine powder is strewn in air, it burns by itself, even without
       being previously heated (it forms a pyrophorus). This obviously
       depends on the fact that the powder of iron presents a larger
       surface of contact with air than an equal weight in a compact
       form.

[Illustration: FIG. 30.--Mode of burning a steel spring in oxygen.]

In order to demonstrate by experiment the _combustion of hydrogen_ in
oxygen, a gas-conducting tube, bent so as to form a convenient jet,
is led from the vessel evolving hydrogen. The hydrogen is first set
light to in air, and then the gas-conducting tube is let down into a
flask containing oxygen. The combustion in oxygen will be similar to
that in air; the flame remains pale, notwithstanding the fact that its
temperature rises considerably. It is instructive to remark that oxygen
may burn in hydrogen, just as hydrogen in oxygen. In order to show the
combustion of oxygen in hydrogen, a tube bent vertically upwards and
ending in a fine orifice is attached to the stopcock of a gas-holder
full of oxygen. Two wires, placed at such a distance from each other as
to allow the passage of a constant series of sparks from a Ruhmkorff's
coil, are fixed in front of the orifice of the tube. This is in order to
ignite the oxygen, which may also be done by attaching tinder round the
orifice, and burning it. When the wires are arranged at the orifice of
the tube, and a series of sparks passes between them, then an inverted
(because of the lightness of the hydrogen) jar full of hydrogen is
placed over the gas-conducting tube. When the jar covers the orifice of
the gas-conducting tube (and not before, as otherwise an explosion might
take place) the cock of the gasometer is opened, and the oxygen flows
into the hydrogen and is set light to by the sparks. The flame obtained
is similar to that formed by the combustion of hydrogen in oxygen.[29]
From this it is evident that the flame is the locality where the oxygen
combines with the hydrogen, therefore a flame of burning oxygen can be
obtained as well as a flame of burning hydrogen.

  [29] The experiment may be conducted without the wires, if the hydrogen
       be lighted in the orifice of an inverted cylinder, and at
       the same time the cylinder be brought over the end of a
       gas-conducting tube connected with a gas-holder containing
       oxygen. Thomsen's method may be adopted for a lecture experiment.
       Two glass tubes, with platinum ends, are passed through orifices,
       about 1-1-1/2 centimetre apart, in a cork. One tube is connected
       with a gas-holder containing oxygen, and the other with a
       gas-holder full of hydrogen. Having turned on the gases, the
       hydrogen is lighted, and a common lamp glass, tapering towards
       the top, is placed over the cork. The hydrogen continues to burn
       inside the lamp glass, at the expense of the oxygen. If the
       current of oxygen be then decreased little by little, a point is
       reached when, owing to the insufficient supply of oxygen, the
       flame of the hydrogen increases in size, disappears for several
       moments, and then reappears at the tube supplying the oxygen. If
       the flow of oxygen be again increased, the flame reappears at the
       hydrogen tube. Thus the flame may be made to appear at one or the
       other tube at will, only the increase or decrease of the current
       of gas must take place by degrees and not suddenly. Further, air
       may be taken instead of oxygen, and ordinary coal-gas instead
       of hydrogen, and it will then be shown how air burns in an
       atmosphere of coal-gas, and it can easily be proved that the lamp
       glass is full of a gas combustible in air, because it may be
       lighted at the top.

If, instead of hydrogen, any other combustible gas be taken--for
example, ordinary coal gas--then the phenomenon of combustion will
be exactly the same, only a bright flame will be obtained, and the
products of combustion will be different. However, as coal gas contains
a considerable amount of free and combined hydrogen, it will also form a
considerable quantity of water in its combustion.

If hydrogen be mixed with oxygen in the proportion in which they form
water--_i.e._ if two volumes of hydrogen be taken for each volume of
oxygen--then the mixture will be the same as that obtained by the
decomposition of water by a galvanic current--detonating gas.

[Illustration: FIG. 31.--Cavendish's apparatus for exploding detonating
gas. The bell jar standing in the bath is filled with a mixture of two
volumes of hydrogen and one volume of oxygen, and the thick glass vessel
A is then screwed on to it. The air is first pumped out of this vessel,
so that when the stopcock C is opened, it becomes filled with detonating
gas. The stopcock is then re-closed, and the explosion produced by means
of a spark from a Leyden jar. After the explosion has taken place the
stopcock is again opened, and the water rises into the vessel A.]

We have already mentioned in the last chapter that the combination of
these gases, or their explosion, may be brought about by the action of
an electric spark, because the spark heats the space through which it
passes, and acts consequently in a manner similar to ignition by means
of contact with an incandescent or burning substance.[29 bis] Cavendish
made this experiment on the ignition of detonating gas, at the end of
the last century, in the apparatus shown in fig. 31. Ignition by the aid
of the electric spark is convenient, for the reason that it may then
be brought about in a closed vessel, and hence chemists still employ
this method when it is required to ignite a mixture of oxygen with a
combustible gas in a closed vessel. For this purpose, especially since
Bunsen's time,[30] an _eudiometer_ is employed. It consists of a thick
glass tube graduated along its length in millimetres (for indicating
the height of the mercury column), and calibrated for a definite volume
(weight of mercury). Two platinum wires are fused into the upper closed
end of the tube, as shown in fig. 32.[31] By the aid of the eudiometer
we may not only determine the volumetric composition of water,[32] and
the quantitative contents of oxygen in air,[33] but also make a number
of experiments explaining the phenomenon of combustion.

[Illustration: FIG. 32.--Eudiometer]

  [29 bis] In fact, instead of a spark a fine wire may be taken, and
       an electric current passed through it to bring it to a state of
       incandescence; in this case there will be no sparks, but the
       gases will inflame if the wire be fine enough to become red hot
       by the passage of the current.

  [30] Now, a great many other different forms of apparatus, sometimes
       designed for special purposes, are employed in the laboratory for
       the investigation of gases. Detailed descriptions of the methods
       of gas analysis, and of the apparatus employed, must be looked
       for in works on analytical and applied chemistry.

  [31] They must be sealed into the tube in such a manner as to leave
       no aperture between them and the glass. In order to test this,
       the eudiometer is filled with mercury, and its open end inverted
       into mercury. If there be the smallest orifice at the wires, the
       external air will enter into the cylinder and the mercury will
       fall, although not rapidly if the orifice be very fine.

  [32] The eudiometer is used for determining the composition of
       combustible gases. A detailed account of _gas analysis_ would be
       out of place in this work (_see_ Note 30), but, as an example,
       we will give a short description of the determination of the
       composition of water by the eudiometer.

       Pure and dry oxygen is first introduced into the eudiometer. When
       the eudiometer and the gas in it acquire the temperature of the
       surrounding atmosphere--which is recognised by the fact of the
       meniscus of the mercury not altering its position during a long
       period of time--then the heights at which the mercury stands in
       the eudiometer and in the bath are observed. The difference (in
       millimetres) gives the height of the column of mercury in the
       eudiometer. It must be reduced to the height at which the mercury
       would stand at 0° and deducted from the atmospheric pressure, in
       order to find the pressure under which the oxygen is measured
       (_see_ Chap. I. Note 29). The height of the mercury also shows
       the volume of the oxygen. The temperature of the surrounding
       atmosphere and the height of the barometric column must also
       be observed, in order to know the temperature of the oxygen
       and the atmospheric pressure. When the volume of the oxygen
       has been measured, pure and dry hydrogen is introduced into
       the eudiometer, and the volume of the gases in the eudiometer
       again measured. They are then exploded. This is done by a Leyden
       jar, whose outer coating is connected by a chain with one wire,
       so that a spark passes when the other wire, fused into the
       eudiometer, is touched by the terminal of the jar. Or else an
       electrophorus is used, or, better still, a Ruhmkorff's coil,
       which has the advantage of working equally well in damp or dry
       air, whilst a Leyden jar or electrical machine does not act in
       damp weather. Further, it is necessary to close the lower orifice
       of the eudiometer before the explosion (for this purpose the
       eudiometer, which is fixed in a stand, is firmly pressed down
       from above on to a piece of india-rubber placed at the bottom
       of the bath), as otherwise the mercury and gas would be thrown
       out of the apparatus by the explosion. It must also be remarked
       that to ensure complete combustion the proportion between the
       volumes of oxygen and hydrogen must not exceed twelve of hydrogen
       to one volume of oxygen, or fifteen volumes of oxygen to one
       volume of hydrogen, because no explosion will take place if one
       of the gases be in great excess. It is best to take a mixture
       of one volume of hydrogen with several volumes of oxygen. The
       combustion will then be complete. It is evident that water is
       formed, and that the volume (or tension) is diminished, so that
       on opening the end of the eudiometer the mercury will rise in
       it. But the tension of the aqueous vapour is now added to the
       tension of the gas remaining after the explosion. This must be
       taken into account (Chap. I. Note 1). If but little gas remain,
       the water which is formed will be sufficient for its saturation
       with aqueous vapour. This may be learnt from the fact that drops
       of water are visible on the sides of the eudiometer after the
       mercury has risen in it. If there be none, a certain quantity of
       water must be introduced into the eudiometer. Then the number of
       millimetres expressing the pressure of the vapour corresponding
       with the temperature of the experiment must be subtracted from
       the atmospheric pressure at which the remaining gas is measured,
       otherwise the result will be inaccurate (Chap. I. Note 1).

       This is essentially the method of the determination of the
       composition of water which was made for the first time by
       Gay-Lussac and Humboldt with sufficient accuracy. Their
       determinations led them to the conclusion that water consists of
       two volumes of hydrogen (more exactly 2·003, Le Duc 1892), and
       one volume of oxygen. Every time they took a greater quantity of
       oxygen, the gas remaining after the explosion was oxygen. When
       they took an excess of hydrogen, the remaining gas was hydrogen;
       and when the oxygen and hydrogen were taken in exactly the above
       proportion, neither one nor the other remained. The composition
       of water was thus definitely confirmed.

  [33] Concerning this application of the eudiometer, see the chapter
       on Nitrogen. It may be mentioned as illustrating the various uses
       of the eudiometer that Prof. Timeraseeff employed microscopically
       small eudiometers to analyse the bubbles of gas given off from
       the leaves of plants.

Thus, for example, it may be demonstrated, by the aid of the eudiometer,
that for the ignition of detonating gas, a _definite temperature_
is required. If the temperature be below that required, combination
will not take place, but if at any spot within the tube it rises to
the temperature of inflammation, then combination will ensue at that
spot, and evolve enough heat for the ignition of the adjacent portions
of the detonating mixture. If to 1 volume of detonating gas there be
added 10 volumes of oxygen, or 4 volumes of hydrogen, or 3 volumes of
carbonic anhydride, then we shall not obtain an explosion by passing a
spark through the diluted mixture. This depends on the fact that the
temperature falls with the dilution of the detonating gas by another
gas, because the heat evolved by the combination of the small quantity
of hydrogen and oxygen brought to incandescence by the spark is not
only transmitted to the water proceeding from the combination, but
also to the foreign substance mixed with the detonating gas.[34] The
necessity of a definite temperature for the ignition of detonating gas
is also seen from the fact that pure detonating gas explodes in the
presence of a red-hot iron wire, or of charcoal heated to 275°, but
with a lower degree of incandescence there is not any explosion. It may
also be brought about by rapid compression, when, as is known, heat
is evolved.[35] Experiments made in the eudiometer showed that the
ignition of detonating gas takes place at a temperature between 450°
and 560°.[36]

  [34] Thus 1/4 volume of carbonic oxide, an equal volume of marsh gas,
       two volumes of hydrogen chloride or of ammonia, and six volumes
       of nitrogen or twelve volumes of air added to one volume of
       detonating gas, prevent its explosion.

  [35] If the compression be brought about slowly, so that the heat
       evolved succeeds in passing to the surrounding space, then the
       combination of the oxygen and hydrogen does not take place, even
       when the mixture is compressed by 150 times; for the gases are
       not heated. If paper soaked with a solution of platinum (in aqua
       regia) and sal ammoniac be burnt, then the ash obtained contains
       very finely-divided platinum, and in this form it is best fitted
       for igniting hydrogen and detonating gas. Platinum wire requires
       to be heated, but platinum in so finely divided a state as it
       occurs in this ash inflames hydrogen, even at -20°. Many other
       metals, such as palladium (175°), iridium, and gold, act with a
       slight rise of temperature, like platinum; but mercury, at its
       boiling point, does not inflame detonating gas, although the
       slow formation of water then begins at 305°. All data of this
       kind show that the explosion of detonating gas presents one of
       the many cases of contact phenomena. This conclusion is further
       confirmed by the researches of V. Meyer (1892). He showed that
       only a very slow formation of steam begins at 448°, and that
       it only proceeds more rapidly at 518°. The temperature of the
       explosion of detonating gas, according to the same author,
       varies according as to whether the explosion is produced in open
       vessels or in closed tubes. In the first case the temperature
       of explosion lies between 530°-606°, and in the second between
       630°-730°. In general it may be remarked that the temperature of
       explosion of gaseous mixtures is always lower in closed vessels
       than when the detonating mixture flows freely through tubes.
       According to Freyer and V. Meyer, the following gases when mixed
       with the requisite amount of oxygen explode at the following
       temperatures:

       +----------------+---------------------+-------------------+
       |      --        | When flowing freely | In closed vessels |
       +----------------+---------------------+-------------------+
       | H_{2}          |     630°-730°       |     530°-606°     |
       | CH_{4}         |     650°-730°       |     606°-650°     |
       | C_{2}H_{6}     |     606°-650°       |     530°-606°     |
       | C_{2}H_{4}     |     606°-650°       |     530°-606°     |
       | CO             |     650°-730°       |     650°-730°     |
       | H_{2}S         |     315°-320°       |     250°-270°     |
       | H_{2} + Cl_{2} |     430°-440°       |     240°-270°     |
       +----------------+---------------------+-------------------+

       The velocity of the transmission of explosion in gaseous mixtures
       is as characteristic a quantity for gaseous systems as the
       velocity of the transmission of sound. Berthelot showed that this
       velocity depends neither upon the pressure nor upon the size of
       the tubes in which the gaseous mixture is contained, nor upon the
       material out of which the tube is made. Dixon (1891) determined
       the magnitude of these velocities for various mixtures, and his
       results proved very near to those previously given by Berthelot.
       For comparison we give the velocities expressed in metres per
       second:

                 +-----------------+-------+-----------+
                 |       --        | Dixon | Berthelot |
                 +-----------------+-------+-----------+
                 | H_{2} + O       | 2,821 |   2,810   |
                 | H_{2} + N_{2}O  | 2,305 |   2,284   |
                 | CH_{4} + 4O     | 2,322 |   2,287   |
                 | C_{2}H_{2} + 6O | 2,364 |   2,210   |
                 | C_{2}H_{2} + 5O | 2,391 |   2,482   |
                 | C_{2}N_{2} + 4O | 2,321 |   2,195   |
                 +-----------------+-------+-----------+

       The addition of oxygen to detonating gas lowers the velocity of
       the transmission of explosion almost as much as the introduction
       of nitrogen. An excess of hydrogen on the contrary raises the
       velocity of transmission. It is remarked that the explosion
       of mixtures of oxygen with marsh gas, ethylene and cyanogen
       is transmitted more quickly if the oxygen be taken in such a
       proportion that the carbon should burn to oxide of carbon, _i.e._
       the velocity of the explosion is less if the oxygen be taken in
       sufficient quantity to form carbonic anhydride. Observations upon
       liquid and solid explosives (Berthelot) show that in this case
       the velocity of transmission of explosion is dependent upon the
       material of the tube. Thus the explosion of liquid nitro-methyl
       ether in glass tubes travels at the rate (in dependence upon
       the diam., from 1 mm.-45 mm.) of from 1,890 to 2,482 metres,
       and in tubes of Britannia metal (3 mm. in diam) at the rate of
       1,230 metres. The harder the tube the greater the velocity of
       transmission of explosion. The following are the velocities for
       certain bodies:

                                                metres
                  Nitro-glycerine                1,300
                  Dynamite                       2,500
                  Nitro-mannite                  7,700
                  Picric acid                    6,500

       In conclusion we may add that Mallard and Le Chatelier (1882)
       observed that in the explosion of a mixture of 1 volume of
       detonating gas with _n_ volumes of an inert gas, the pressure is
       approximately equal to 9·2-0·9_n_ atmospheres.

  [36] From the very commencement of the promulgation of the idea of
       dissociation, it might have been imagined that reversible
       reactions of combination (the formation of H_{2} and O belongs
       to this number) commence at the same temperature as that at
       which dissociation begins. And in many cases this is so, but
       not always, as may be seen from the facts (1) that at 450-560°,
       when detonating gas explodes, the density of aqueous vapour not
       only does not vary (and it hardly varies at higher temperatures,
       probably because the amount of the products of dissociation is
       small), but there are not, as far as is yet known, any traces
       of dissociation; (2) that under the influence of contact the
       temperature at which combination takes place falls even to the
       ordinary temperature, when water and similar compounds naturally
       are not dissociated and, judging from the data communicated
       by D. P. Konovaloff (Introduction, Note 39) and others, it is
       impossible to escape the phenomena of contact; all vessels,
       whether of metal or glass, show the same influence as spongy
       platinum, although to a much less degree. The phenomena of
       contact, judging from a review of the data referring to it,
       must be especially sensitive in reactions which are powerfully
       exothermal, and the explosion of detonating gas is of this kind.

The combination of hydrogen with oxygen is accompanied by the evolution
of a very considerable amount of heat; according to the determinations
of _Favre_ and _Silbermann_,[37] 1 part by weight of hydrogen in
forming water evolves 34,462 units of heat. Many of the most recent
determinations are very close to this figure, so that it may be taken
that in the formation of 18 parts of water (H_{2}O) there are evolved 69
major calories, or 69,000 units of heat.[38] _If the specific heat of
aqueous vapour_ (0·48) _remained constant from the ordinary temperature
to that at which the combustion of detonating gas takes place_
(but there is now no doubt that it increases), were the combustion
concentrated at one point[39] (but it occurs in the whole region of a
flame), were there no loss from radiation and heat conduction, and _did
dissociation not take place_--that is, did not a state of equilibrium
between the hydrogen, oxygen, and water come about--_then it would be
possible to calculate the temperature of the flame of detonating gas_.
It would then be 8,000°.[40] In reality it is very much lower, but it
is nevertheless higher than the temperature attained in furnaces and
flames, and is as high as 2,000°. The explosion of detonating gas is
explained by this high temperature, because the aqueous vapour formed
must occupy a volume at least 5 times greater than that occupied by
the detonating gas at the ordinary temperature. Detonating gas emits a
sound, not only as a consequence of the commotion which occurs from the
rapid expansion of the heated vapour, but also because it is immediately
followed by a cooling effect, the conversion of the vapour into water,
and a rapid contraction.[41]

  [37] [Illustration: FIG. 33.--Favre and Silbermann's calorimeter for
       determining the heat evolved in combustion.]

       The amount of heat evolved in the combustion of a known weight
       (for instance, 1 gram) of a given substance is determined by the
       rise in temperature of water, to which the whole of the heat
       evolved in the combustion is transmitted. A _calorimeter_, for
       example that shown in fig. 33, is employed for this purpose.
       It consists of a thin (in order that it may absorb less heat),
       polished (that it should transmit a minimum of heat) metallic
       vessel, surrounded by down (_c_), or some other bad conductor
       of heat, and an outer metallic vessel. This is necessary in
       order that the least possible amount of heat should be lost
       from the vessels; nevertheless, there is always a certain loss,
       whose magnitude is determined by preliminary experiment (by
       taking warm water, and determining its fall in temperature after
       a definite period of time) as a correction for the results
       of observations. The water to which the heat of the burning
       substance is transmitted is poured into the vessel. The stirrer
       _g_ allows of all the layers of water being brought to the same
       temperature, and the thermometer serves for the determination
       of the temperature of the water. The heat evolved passes,
       naturally, not to the water only, but to all the parts of the
       apparatus. The quantity of water corresponding to the whole
       amount of those objects (the vessels, tubes, &c.) to which the
       heat is transmitted is previously determined, and in this manner
       another most important correction is made in the calorimetric
       determinations. The combustion itself is carried on in the vessel
       _a_. The ignited substance is introduced through the tube at the
       top, which closes tightly. In fig. 33 the apparatus is arranged
       for the combustion of a gas, introduced by a tube. The oxygen
       required for the combustion is led into _a_ by the tube _e_, and
       the products of combustion either remain in the vessel _a_ (if
       liquid or solid), or escape by the tube _f_ into an apparatus in
       which their quantity and properties can easily be determined.
       Thus the heat evolved in combustion passes to the walls of the
       vessel _a_, and to the gases which are formed in it, and these
       transmit it to the water of the calorimeter.

  [38] This quantity of heat corresponds with the formation of
       liquid water at the ordinary temperature from detonating gas at
       the same temperature. If the water be as vapour the heat evolved
       = 58 major calories; if as ice = 70·4 major calories. A portion
       of this heat is due to the fact that 2 vols. of hydrogen and 1
       vol. of oxygen give 2 vols. of aqueous vapour--that is to say,
       contraction ensues--and this evolves heat. This quantity of heat
       may be calculated, but it cannot be said how much is expended
       in the separation of the atoms of oxygen from each other, and,
       therefore, strictly speaking, we do not know the quantity of
       heat which is evolved in the reaction alone, although the number
       of units of heat evolved in the combustion of detonating gas is
       accurately known.

       The construction of the calorimeter and even the method of
       determination vary considerably in different cases. Since the
       beginning of the nineties, a large number of determinations
       of the heat of combustion have been conducted in closed bombs
       containing compressed oxygen. The greatest number of calorimetric
       determinations were made by Berthelot and Thomsen. They are
       given in their works _Essai de mécanique chimique fondée
       sur la thermochimie_, by M. Berthelot, 1879 (2 vols.), and
       _thermochemische Untersuchungen_, by J. Thomsen, 1886 (4 vols.)
       The most important methods of recent thermochemistry, and all
       the trustworthy results of experiment, are given in Prof. P. F.
       Louginin's _Description of the Different Modes of Determining
       the Heat of Combustion of Organic Compounds_, Moscow, 1894.
       The student must refer to works on theoretical and physical
       chemistry for a description of the elements and methods of
       _thermochemistry_, into the details of which it is impossible to
       enter in this work. One of the originators of thermochemistry,
       Hess, was a member of the St. Petersburg Academy of Sciences.
       Since 1870 a large amount of research has been carried out in
       this province of chemistry, especially in France and Germany,
       after the investigations of the French Academician, Berthelot,
       and Professor Thomsen, of Copenhagen. Among Russians, Beketoff,
       Louginin, Cheltzoff, Chroustchoff, and others are known by their
       thermochemical researches. The present epoch of thermochemistry
       must be considered rather as a collective one, wherein the
       material of facts is amassed, and the first consequences arising
       from them are noticed. In my opinion two essential circumstances
       prevent the possibility of deducing any exact consequences, of
       importance to chemical mechanics, from the immense store of
       thermochemical data already collected: (1) The majority of the
       determinations are conducted in weak aqueous solutions, and, the
       heat of solution being known, are referred to the substances
       in solution; yet there is much (Chapter I.) which leads to the
       conclusion that in solution water does not play the simple
       part of a diluting medium, but of itself acts independently
       in a chemical sense on the substance dissolved. (2) Physical
       and mechanical changes (decrease of volume, diffusion, and
       others) invariably proceed side by side with chemical changes,
       and for the present it is impossible, in a number of cases, to
       distinguish the thermal effect of the one and the other kind of
       change. It is evident that the one kind of change (chemical) is
       essentially inseparable and incomprehensible without the other
       (mechanical and physical); and therefore it seems to me that
       thermochemical data will only acquire their true meaning when
       the connection between the phenomena of both kinds (on the one
       hand chemical and atomic, and on the other hand mechanical and
       molecular or between entire masses) is explained more clearly
       and fully than is at present the case. As there is no doubt that
       the simple mechanical contact, or the action of heat alone,
       on substances sometimes causes an evident and always a latent
       (incipient) chemical change--that is, a different distribution
       or motion of the atoms in the molecules--it follows that purely
       chemical phenomena are inseparable from physical and mechanical
       phenomena. A mechanical change may be imagined without a physical
       change, and a physical without a chemical change, but it is
       impossible to imagine a chemical change without a physical and
       mechanical one, for without the latter we should not be able
       to recognise the former, and it is by their means that we are
       enabled to do so.

  [39] The flame, or locality where the combustion of gases and vapours
       takes place, is a complex phenomenon, 'an entire factory,' as
       Faraday says, and therefore we will consider flame in some detail
       in one of the following notes.

  [40] If 34,500 units of heat are evolved in the combustion of 1 part
       of hydrogen, and this heat is transmitted to the resulting 9
       parts by weight of aqueous vapour, then we find that, taking the
       specific heat of the latter as 0·475, each unit of heat raises
       the temperature of 1 part by weight of aqueous vapour 2°·1 and 9
       parts by weight (2·1 ÷ 9) O°·23; hence the 34,500 units of heat
       raise its temperature 7,935°. If detonating gas is converted
       into water in a closed space, then the aqueous vapour formed
       cannot expand, and therefore, in calculating the temperature of
       combustion, the specific heat at a constant volume must be taken
       into consideration; for aqueous vapour it is 0·36. This figure
       gives a still higher temperature for the flame. In reality it
       is much lower, but the results given by different observers are
       very contradictory (from 1,700° to 2,400°), the discrepancies
       depending on the fact that flames of different sizes are cooled
       by radiation to a different degree, but mainly on the fact that
       the methods and apparatus (pyrometers) for the determination
       of high temperatures, although they enable relative changes of
       temperature to be judged, are of little use for determining their
       absolute magnitude. By taking the temperature of the flame of
       detonating gas as 2,000°, I give, I think, the average of the
       most trustworthy determinations and calculations based upon the
       determination of the variation of the specific heat of aqueous
       vapour and other gases (_see_ Chapter XLI.)

  [41] It is evident that not only hydrogen, but every other combustible
       gas, will give an explosive mixture with oxygen. For this reason
       coal-gas mixed with air explodes when the mixture is ignited.
       The pressure obtained in the explosions serves as the _motive
       power of gas engines_. In this case advantage is taken, not only
       of the pressure produced by the explosion, but also of that
       contraction which takes place after the explosion. On this is
       based the construction of several motors, of which Lenoir's was
       formerly, and Otto's is now, the best known. The explosion is
       usually produced by coal-gas and air, but of late the vapours of
       combustible liquids (kerosene, benzene) are also being employed
       in place of gas (Chapter IX.) In Lenoir's engine a mixture of
       coal-gas and air is ignited by means of sparks from a Ruhmkorff's
       coil, but in the most recent machines the gases are ignited by
       the direct action of a gas jet, or by contact with the hot walls
       of a side tube.

[Illustration: FIG. 34.--Safety burner for detonating gas, described in
text.]

Mixtures of hydrogen and of various other gases with oxygen are taken
advantage of for obtaining high temperatures. By the aid of such high
temperatures metals like platinum may be melted on a large scale,
which cannot be performed in furnaces heated with charcoal and fed by
a current of air. The burner, shown in fig. 34, is constructed for
the application of detonating gas to the purpose. It consists of two
brass tubes, one fixed inside the other, as shown in the drawing. The
internal central tube C C conducts oxygen, and the outside, enveloping,
tube E' E' conducts hydrogen. Previous to their egress the gases
do not mix together, so that there can be no explosion inside the
apparatus. When this burner is in use C is connected with a gas-holder
containing oxygen, and E with a gas-holder containing hydrogen (or
sometimes coal-gas). The flow of the gases can be easily regulated by
the stopcocks O H. The flame is shortest and evolves the greatest heat
when the gases burning are in the proportion of 1 volume of oxygen to
2 volumes of hydrogen. The degree of heat may be easily judged from
the fact that a thin platinum wire placed in the flame of a properly
proportioned mixture easily melts. By placing the burner in the orifice
of a hollow piece of lime, a crucible A B is obtained in which the
platinum may be easily melted, even in large quantities if the current
of oxygen and hydrogen be sufficiently great (Deville). The flame
of detonating gas may also be used for illuminating purposes. It is
by itself very pale, but owing to its high temperature it may serve
for rendering infusible objects incandescent, and at the very high
temperature produced by the detonating gas the incandescent substance
gives a most intense light. For this purpose lime, magnesia, or oxide of
zirconium are used, as they are not fusible at the very high temperature
evolved by the detonating gas. A small cylinder of lime placed in
the flame of detonating gas, if regulated to the required point,
gives a very brilliant white light, which was at one time proposed
for illuminating lighthouses. At present in the majority of cases
the electric light, owing to its constancy and other advantages, has
replaced it for this purpose. The light produced by the incandescence of
lime in detonating gas is called the _Drummond light_ or _limelight_.

The above cases form examples of the combustion of elements in oxygen,
but exactly similar phenomena are observed in the _combustion of
compounds_. So, for instance, the solid, colourless, shiny substance,
naphthalene, C_{10}H_{8}, burns in the air with a smoky flame, whilst
in oxygen it continues to burn with a very brilliant flame. Alcohol,
oil, and other substances burn brilliantly in oxygen on conducting the
oxygen by a tube to the flame of lamps burning these substances. A high
temperature is thus evolved, which is sometimes taken advantage of in
chemical practice.

In order to understand why combustion in oxygen proceeds more rapidly,
and is accompanied by a more intense heat effect, than combustion in
air, it must be recollected that air is oxygen diluted with nitrogen,
which does not support combustion, and therefore fewer particles of
oxygen flow to the surface of a substance burning in air than when
burning in pure oxygen, besides which the reason of the intensity of
combustion in oxygen is the high temperature acquired by the substance
burning in it.[41 bis]

  [41 bis] Let us consider as an example the combustion of
       sulphur in air and in oxygen. If 1 gram of sulphur burns in air
       or oxygen it evolves in either case 2250 units of heat--_i.e._
       evolves sufficient heat for heating 2,250 grams of water 1°
       C. This heat is first of all transmitted to the sulphurous
       anhydride, SO_{2}, formed by the combination of sulphur with
       oxygen. In its combustion 1 gram of sulphur forms 2 grams of
       sulphurous anhydride--_i.e._ the sulphur combines with 1 gram of
       oxygen. In order that 1 gram of sulphur should have access to 1
       gram of oxygen in air, it is necessary that 3·4 grams of nitrogen
       should simultaneously reach the sulphur, because air contains
       seventy-seven parts of nitrogen (by weight) per twenty-three
       parts of oxygen. Thus in the combustion of 1 gram of sulphur,
       the 2,250 units of heat are transmitted to 2 grams of sulphurous
       oxide and to at least 3·4 grams of nitrogen. As 0·155 unit of
       heat is required to raise 1 gram of sulphurous anhydride 1°
       C., therefore 2 grams require 0·31 unit. So also 3·4 grams of
       nitrogen require 3·4 × 0·244 or 0·83 unit of heat, and therefore
       in order to raise both gases 1° C. 0·31 + 0·83 or 1·14 unit of
       heat is required; but as the combustion of the sulphur evolves
       2,250 units of heat, therefore the gases might be heated (if
       their specific heats remained constant) to 2250/1·14 or 1,974°
       C. That is, the maximum possible temperature of the flame of the
       sulphur burning in air will be 1,974° C. In the combustion of the
       sulphur in oxygen the heat evolved (2,250 units) can only pass to
       the 2 grams of sulphurous anhydride, and therefore the highest
       possible temperature of the flame of the sulphur in oxygen will
       be = 2250/0·31 or 7258°. In the same manner it may be calculated
       that the temperature of charcoal burning in air cannot exceed
       2,700°, while in oxygen it may attain 10,100° C. For this reason
       the temperature in oxygen will always be higher than in air,
       although (judging from what has been said respecting detonating
       gas) neither one temperature nor the other will ever approximate
       to the theoretical amount.

[Illustration: FIG. 35.--Faraday's experiment for investigating the
different parts of a candle flame.]

Among the phenomena accompanying the combustion of certain substances,
the _phenomenon of flame_ attracts attention. Sulphur, phosphorus,
sodium, magnesium, naphthalene, &c., burn like hydrogen with a flame,
whilst in the combustion of other substances no flame is observed, as,
for instance, in the combustion of iron and of charcoal. The appearance
of flame depends on the capacity of the combustible substance to yield
gases or vapours at the temperature of combustion. At the temperature
of combustion, sulphur, phosphorus, sodium, and naphthalene pass into
vapour, whilst wood, alcohol, oil, &c., are decomposed into gaseous
and vaporous substances. The combustion of gases and vapours forms
flames, and therefore _a flame is composed of the hot and incandescent
gases and vapours produced by combustion_. It may easily be proved that
the flames of such non-volatile substances as wood contain volatile
and combustible substances formed from them, by placing a tube in the
flame connected with an aspirator. Besides the products of combustion,
combustible gases and liquids, previously in the flame as vapours,
collect in the aspirator. For this experiment to succeed--_i.e._ in
order to really extract combustible gases and vapours from the flame
it is necessary that the suction tube should be placed _inside_ the
flame. The combustible gases and vapours can only remain unburnt inside
the flame, for at the surface of the flame they come into contact with
the oxygen of the air and burn.[42] Flames are of different degrees
of _brilliancy_, according to whether _solid_ incandescent particles
occur in the combustible gas or vapour, or not. Incandescent gases
and vapours emit but little light by themselves, and therefore give
a paler flame.[43] If a flame does not contain solid particles it is
transparent, pale, and emits but little light.[44] The flames of
burning alcohol, sulphur, and hydrogen are of this kind. A pale flame
may be rendered luminous by placing fine particles of solid matter in
it. Thus, if a very fine platinum wire be placed in the pale flame
of burning alcohol--or, better still, of hydrogen--the flame emits
a bright light. This is still better seen by sifting the powder of
an incombustible substance, such as fine sand, into the flame, or
by placing a bunch of asbestos threads in it. Every brilliant flame
always contains some kind of solid particles, or at least some very
dense vapour. The flame of sodium burning in oxygen has a brilliant
yellow colour, from the presence of particles of solid sodium oxide.
The flame of magnesium is brilliant from the fact that in burning
it forms solid magnesia, which becomes white hot, and similarly the
brilliancy of the Drummond light is due to the heat of the flame
raising the solid non-volatile lime to a state of incandescence.
The flames of a candle, wood, and similar substances are brilliant,
because they contain particles of charcoal or soot. It is not the flame
itself which is luminous, but the incandescent soot it contains. These
particles of charcoal which occur in flames may be easily observed
by introducing a cold object, like a knife, into the flame.[45] The
particles of charcoal burn at the outer surface of the flame if the
supply of air be sufficient, but if the supply of air--that is, of
oxygen--be insufficient for their combustion the flame smokes, because
the unconsumed particles of charcoal are carried off by the current of
air.[46]

  [42] Faraday proved this by a very convincing experiment on
       a candle flame. If one arm of a bent glass tube be placed in a
       candle flame above the wick in the dark portion of the flame,
       then the products of the partial combustion of the stearin will
       pass up the tube, condense in the other arm, and collect in a
       flask placed under it (fig. 35) as heavy white fumes which burn
       when lighted. If the tube be raised into the upper luminous
       portion of the flame, then a dense black smoke which will not
       inflame accumulates in the flask. Lastly, if the tube be let down
       until it touches the wick, then little but stearic acid condenses
       in the flask.

  [43] All transparent substances which transmit light with great ease
       (that is, which absorb but little light) are but little luminous
       when heated; so also substances which absorb but few heat rays,
       when heated transmit few rays of heat.

  [44] There is, however, no doubt but that very heavy dense vapours or
       gases under pressure (according to the experiments of Frankland)
       are luminous when heated, because, as they become denser they
       approach a liquid or solid state. Thus detonating gas when
       exploded under pressure gives a brilliant light.

  [45] If hydrogen gas be passed through a volatile liquid hydrocarbon--for
       instance, through benzene (the benzene may be poured directly
       into the vessel in which hydrogen is generated)--then its vapour
       burns with the hydrogen and gives a very bright flame, because
       the resultant particles of carbon (soot) become incandescent.
       Benzene, or platinum gauze, introduced into a hydrogen flame may
       be employed for illuminating purposes.

  [46] In _flames_ the separate parts may be distinguished with more or
       less distinctness. That portion of the flame whither the
       combustible vapours or gases flow, is not luminous because its
       temperature is still too low for the process of combustion to
       take place in it. This is the space which in a candle surrounds
       the wick, or in a gas jet is immediately above the orifice from
       which the gas escapes. In a candle the combustible vapours and
       gases which are formed by the action of heat on the melted
       tallow or stearin rise in the wick, and are heated by the high
       temperature of the flame. By the action of the heat, the solid or
       liquid substance is here, as in other cases, decomposed, forming
       products of dry distillation. These products occur in the central
       portion of the flame of a candle. The air travels to it from the
       outside, and is not able to intermix at once with the vapours
       and gases in all parts of the flame equally; consequently, in
       the outer portion of the flame the amount of oxygen will be
       greater than in the interior portions. But, owing to diffusion,
       the oxygen, of course mixed with nitrogen, flowing towards the
       combustible substance, does finally penetrate to the interior
       of the flame (when the combustion takes place in ordinary air).
       The combustible vapours and gases combine with this oxygen,
       evolve a considerable amount of heat, and bring about that state
       of incandescence which is so necessary both for keeping up the
       combustion and also for the uses to which the flame is applied.
       Passing from the colder envelope of air through the interior
       of the flame, to the source of the combustible vapours (for
       instance, the wick), we evidently first traverse layers of higher
       and higher temperature, and then portions which are less and
       less hot, in which the combustion is less complete, owing to the
       limited supply of oxygen.

       [Illustrationtion: FIG. 36.--In the candle flame the portion C
       contains the vapours and products of decomposition; in the bright
       zone A the combustion has commenced, and particles of carbon are
       emitted; and in the pale zone B the combustion is completed.]

       Thus unburnt products of the decomposition of organic substances
       occur in the interior of the flame. But there is always free
       hydrogen in the interior of the flame, even when oxygen is
       introduced there, or when a mixture of hydrogen and oxygen burns,
       because the temperature evolved in the combustion of hydrogen
       or the carbon of organic matter is so high that the products
       of combustion are themselves partially decomposed--that is,
       dissociated--at this temperature. Hence, in a flame a portion
       of the hydrogen and of the oxygen which might combine with
       the combustible substances must always be present in a free
       state. If a hydrocarbon burns, and we imagine that a portion
       of the hydrogen is in a free state, then a portion of the
       carbon must also occur in the same form in the flame, because,
       other conditions being unchanged, carbon burns after hydrogen,
       and this is actually observed in the combustion of various
       hydrocarbons. Charcoal, or the soot of a common flame, arises
       from the dissociation of organic substances contained in the
       flame. The majority of hydrocarbons, especially those containing
       much carbon--for instance, naphthalene--burn, even in oxygen,
       with separation of soot. In that portion of the flame where the
       hydrogen burns the carbon remains unburnt, or at least partly so.
       It is this free carbon which causes the brilliancy of the flame.
       That the interior of the flame contains a mixture which is still
       capable of combustion may be proved by the following experiment:
       A portion of the gases may be withdrawn by an aspirator from
       the central portion of the flame of carbonic oxide, which is
       combustible in air. For this purpose Deville passed water
       through a metallic tube having a fine lateral orifice, which is
       placed in the flame. As the water flows along the tube portions
       of the gases of the flame enter, and, passing along the tube
       alternately with cylinders of water, are carried away into an
       apparatus where they can be investigated. It appears that all
       portions of the flame obtained by the combustion of a mixture
       of carbonic oxide and oxygen contain a portion of this mixture
       still unburnt. The researches of Deville and Bunsen showed that
       in the explosion of a mixture of hydrogen and of carbonic oxide
       with oxygen in a closed space, complete combustion does not ever
       take place immediately. If two volumes of hydrogen and one volume
       of oxygen be confined in a closed space, then on explosion the
       pressure does not attain that magnitude which it would were there
       immediate and complete combustion. It may be calculated that the
       pressure should attain twenty-six atmospheres. In reality, it
       does not exceed nine and a half atmospheres.

       Hence the admixture of the products of combustion with an
       explosive mixture prevents the combustion of the remaining
       mass, although capable of burning. The admixture of carbonic
       anhydride prevents carbonic oxide from burning. The presence of
       any other foreign gas interferes in the same manner. This shows
       that every portion of a flame must contain combustible, burning,
       and already burnt substances--_i.e._ oxygen, carbon, carbonic
       oxide, hydrogen, hydrocarbons, carbonic anhydride, and water.
       Consequently, _it is impossible to attain instantaneous complete
       combustion_, and this is one of the reasons of the phenomenon
       of flame. A certain space is required, and the temperature must
       be unequal in different parts of it. In this space different
       quantities of the component parts are successively subjected
       to combustion, or are cooled under the influence of adjacent
       objects, and combustion only ends where the flame ends. If
       the combustion could be concentrated at one spot, then the
       temperature would be incomparably higher than it is under the
       actual circumstances.

       The various regions of the flame have formed the frequent
       subject of experimental research, and the experiments conducted
       by Smithells and Ingle (1892) are particularly instructive;
       they show that the reducing (interior) and oxidising (exterior)
       portions of the flame of a burning gas may be divided by taking a
       Bunsen burner and surrounding the flame of the gas burnt in it,
       by another wider tube (without the access of air to the annular
       space or allowing only a small current of air to pass), when
       a gaseous mixture, containing oxide of carbon and capable of
       further combustion, will issue from this enveloping tube, so that
       a second flame, corresponding to the exterior (oxidising) portion
       of an ordinary flame, may be obtained above the enveloping tube.
       This division of the flame into two portions is particularly
       clear when cyanogen C_{2}N_{2} is burnt, because the interior
       portion (where CO is chiefly formed according to the equation
       C_{2}N_{2} + O_{2} = 2CO + N_{2}, but a portion of the nitrogen
       is oxidised) is of a rose colour, while the exterior portion
       (where the CO burns into CO_{2} at the expense of a fresh
       quantity of oxygen and of the oxides of nitrogen proceeding from
       the interior portions) is of a bluish-grey colour.

The combination of various substances with oxygen may not present
any signs of combustion--that is, the temperature may rise but
inconsiderably. This may either proceed from the fact that the
reaction of the substance (for example, tin, mercury, lead at a
high temperature, or a mixture of pyrogallol with caustic potash at
the ordinary temperature) evolves but little heat, or that the heat
evolved is transmitted to good conductors of heat, like metals, or
that the combination with oxygen takes place so slowly that the heat
evolved succeeds in passing to the surrounding objects. Combustion
is only a particular, intense, and evident case of combination with
oxygen. Respiration is also an act of combination with oxygen; it
also serves, like combustion, for the development of heat by those
chemical processes which accompany it (the transformation of oxygen into
carbonic anhydride). Lavoisier enunciated this in the lucid expression,
'respiration is slow combustion.'

Reactions involving slow combination of substances with oxygen are
termed _oxidations_. Combination of this kind (and also combustion)
often results in the formation of acid substances, and hence the name
_oxygen_ (_Sauerstoff_). Combustion is only rapid oxidation. Phosphorus,
iron, and wine may be taken as examples of substances which slowly
oxidise in air at the ordinary temperature. If such a substance be
left in contact with a definite volume of air or oxygen, it absorbs
the oxygen little by little, as may be seen by the decrease in volume
of the gas. This slow oxidation is not often accompanied by a sensible
evolution of heat; an evolution of heat really does occur, only it
is not apparent to our senses owing to the small rise in temperature
which takes place; this is owing to the slow rate of the reaction and
to the transmission of the heat formed as radiant heat, &c. Thus, in
the oxidation of wine and its transformation into vinegar by the usual
method of preparation of the latter, the heat evolved cannot be observed
because it extends over several weeks, but in the so-called rapid
process of the manufacture of vinegar, when a large quantity of wine is
comparatively rapidly oxidised, the evolution of heat is quite apparent.

Such slow processes of oxidation are always taking place in nature by
the action of the atmosphere. Dead organisms and the substances obtained
from them--such as bodies of animals, wood, wool, grass, &c.--are
especially subject to this action. They _rot_ and _decompose_--that is,
their solid matter is transformed into gases, under the influence of
moisture and atmospheric oxygen, and generally under the influence of
other organisms, such as moulds, worms, micro-organisms (bacteria), and
the like. These are processes of slow combustion, of slow combination
with oxygen. It is well known that manure rots and develops heat,
that stacks of damp hay, damp flour, straw, &c., become heated and
are changed in the process.[47] In all these transformations the same
chief products of combustion are formed as those which are contained
in smoke; the carbon gives carbonic anhydride, and the hydrogen water.
Hence these processes require oxygen just like combustion. This is
the reason why the entire prevention of access of air hinders these
transformations,[48] and an increased supply of air accelerates them.
The mechanical treatment of arable lands by the plough, harrow, and
other similar means has not only the object of facilitating the spread
of roots in the ground, and of making the soil more permeable to water,
but it also serves to facilitate the access of the air to the component
parts of the soil; as a consequence of which the organic remains of soil
rot--so to speak, breathe air and evolve carbonic anhydride. One acre
of good garden land in the course of a summer evolves more than sixteen
tons of carbonic anhydride.

  [47] Cotton waste (used in factories for cleaning machines from
       lubricating oil) soaked in oil and lying in heaps is
       self-combustible, being oxidised by the air.

  [48] When it is desired to preserve a supply of vegetable and animal
       food, the access of the oxygen of the atmosphere (and also of
       the germs of organisms present in the air) is often prevented.
       With this object articles of food are often kept in hermetically
       closed vessels, from which the air has been withdrawn; vegetables
       are dried and soldered up while hot in tin boxes; sardines are
       immersed in oil, &c. The removal of water from substances is also
       sometimes resorted to with the same object (the drying of hay,
       corn, fruits), as also is saturation with substances which absorb
       oxygen (such as sulphurous anhydride), or which hinder the growth
       of organisms forming the first cause of putrefaction, as in
       processes of smoking, embalming, and in the keeping of fishes and
       other animal specimens in spirit, &c.

It is not only vegetable and animal substances which are subject to slow
oxidation in the presence of water. Some metals even rust under these
conditions. Copper very easily absorbs oxygen in the presence of acids.
Many metallic sulphides (for example, pyrites) are very easily oxidised
with access of air and moisture. Thus processes of slow oxidation
proceed throughout nature. However, there are many elements which do
not under any circumstances combine directly with gaseous oxygen;
nevertheless their compounds with oxygen may be obtained. Platinum,
gold, iridium, chlorine, and iodine are examples of such elements. In
this case recourse is had to a so-called _indirect method_--_i.e._ the
given substance is combined with another element, and by a method of
double decomposition this element is replaced by oxygen. Substances
which do not directly combine with oxygen, but form compounds with it
by an indirect method, often readily lose the oxygen which they had
absorbed by double decomposition or at the moment of its evolution.
Such, for example, are the compounds of oxygen with chlorine, nitrogen,
and platinum, which evolve oxygen on heating--that is, they may be
used as oxidising agents. In this respect _oxidising agents_, or those
compounds of oxygen which are employed in chemical and technical
practice for transferring oxygen to other substances, are especially
remarkable. The most important among these is nitric acid or _aqua
fortis_--a substance rich in oxygen, and capable of evolving it when
heated, which easily oxidises a great number of substances. Thus nearly
all metals and organic substances containing carbon and hydrogen are
more or less oxidised when heated with nitric acid. If strong nitric
acid be taken, and a piece of burning charcoal be immersed in the acid,
it continues to burn. Chromic acid acts like nitric acid; alcohol burns
when mixed with it. Although the action is not so marked, even water
may oxidise with its oxygen. Sodium is not oxidised in perfectly dry
oxygen at the ordinary temperature, but it burns very easily in water
and aqueous vapour. Charcoal can burn in carbonic anhydride--a product
of combustion--forming carbonic oxide. Magnesium burns in the same gas,
separating carbon from it. Speaking generally, combined oxygen can pass
from one compound to another.

The products of combustion or oxidation--and in general the definite
compounds of oxygen--are termed _oxides_. Some oxides are not capable of
combining with other oxides--or combine with only a few, and then with
the evolution of very little heat; others, on the contrary, enter into
combination with very many other oxides, and in general have remarkable
chemical energy. The oxides incapable of combining with others, or only
showing this quality in a small degree, are termed _indifferent oxides_.
Such are the peroxides, of which mention has before been made.

The class of oxides capable of entering into mutual combination we will
term _saline oxides_. They fall into two chief groups--at least, as
regards the most extreme members. The members of one group combine with
the members of the other group with particular ease. As representative
of one group may be taken the oxides of the metals, magnesium, sodium,
calcium, &c. Representatives of the other group are the oxides formed
by the non-metals, sulphur, phosphorus, carbon. Thus, if we take the
oxide of calcium, or lime, and bring it into contact with oxides of the
second group, combination very readily ensues. For instance, if we mix
calcium oxide with oxide of phosphorus they combine with great facility
and with the evolution of much heat. If we pass the vapour of sulphuric
anhydride, obtained by the combination of sulphurous oxide with oxygen,
over pieces of lime heated to redness, the sulphuric anhydride is
absorbed by the lime with the formation of a substance called calcium
sulphate. The oxides of the first kind, which contain metals, are
termed _basic oxides_ or _bases_. Lime is a familiar example of this
class. The oxides of the second group, which are capable of combining
with the bases, are termed _anhydrides of the acids_ or _acid oxides_.
Sulphuric anhydride, SO_{3}, may be taken as a type of the latter group.
It is a compound of sulphur with oxygen formed not directly but by the
addition of a fresh quantity of oxygen to sulphurous anhydride, SO_{2},
by passing it together with oxygen over incandescent spongy platinum.
Carbonic anhydride (often termed 'carbonic acid'), CO_{2}, phosphoric
anhydride, sulphurous anhydride, are all acid oxides, for they can
combine with such oxides as lime or calcium oxide, magnesia or magnesium
oxide, MgO, soda or sodium oxide, Na_{2}O, &c.

If a given element form but one basic oxide, it is termed the _oxide_;
for example, calcium oxide, magnesium oxide, potassium oxide. Some
indifferent oxides are also called 'oxides' if they have not the
properties of peroxides, and at the same time do not show the properties
of acid anhydrides--for example, carbonic oxide, of which mention
has already been made. If an element forms two basic oxides (or two
indifferent oxides not having the characteristics of a peroxide) then
that of the lower degree of oxidation is called a _suboxide_--that is,
suboxides contain less oxygen than oxides. Thus, when copper is heated
to redness in a furnace it increases in weight and absorbs oxygen,
until for 63 parts of copper there is absorbed not more than 8 parts
of oxygen by weight, forming a red mass, which is suboxide of copper;
but if the roasting be prolonged, and the draught of air increased,
63 parts of copper absorb 16 parts of oxygen, and form black oxide of
copper. Sometimes to distinguish between the degrees of oxidation a
change of suffix is made in the oxidised element, _-ic_ oxide denoting
the higher degree of oxidation, and _-ous_ oxide the lower degree.
Thus ferrous oxide and ferric oxide are the same as suboxide of iron
and oxide of iron. If an element forms one anhydride only, then it is
named by an adjective formed from the name of the element made to end in
_-ic_ and the word _anhydride_. When an element forms two anhydrides,
then the suffixes _-ous_ and _-ic_ are used to distinguish them: _-ous_
signifying less oxygen than _-ic_; for example, sulphurous and sulphuric
anhydrides.[49] When several oxides are formed from the same element,
the prefixes _mon_, _di_, _tri_, _tetra_ are used, thus: chlorine
monoxide, chlorine dioxide, chlorine trioxide, and chlorine tetroxide
or chloric anhydride.

  [49] It must be remarked that certain elements form oxides of all three
       kinds--_i.e._ indifferent, basic, and acid; for example,
       manganese forms manganous oxide, manganic oxide, peroxide of
       manganese, red oxide of manganese, and manganic anhydride,
       although some of them are not known in a free state but only
       in combination. The basic oxides contain less oxygen than the
       peroxides, and the peroxides less than the acid anhydrides.
       Thus they must be placed in the following general normal
       order with respect to the amount of oxygen entering into
       their composition--(1) basic oxides, suboxides, and oxides;
       (2) peroxides; (3) acid anhydrides. The majority of elements,
       however, do not give all three kinds of oxides, some giving only
       one degree of oxidation. It must further be remarked that there
       are oxides formed by the combination of acid anhydrides with
       basic oxides, or, in general, of oxides with oxides. For every
       oxide having a higher and a lower degree of oxidation, it might
       be said that the intermediate oxide was formed by the combination
       of the higher with the lower oxide. But this is not true in all
       cases--for instance, when the oxide under consideration forms
       a whole series of independent compounds--for oxides which are
       really formed by the combination of two other oxides do not give
       such independent compounds, but in many cases decompose into the
       higher and lower oxides.

The oxides themselves rarely undergo chemical transformations, and in
the few cases where they are subject to such changes a particularly
important part is played by their combinations with water. The majority
of, if not all, basic and acid oxides combine with water, either by a
direct or an indirect method forming _hydrates_--that is, compounds
which split up into water and an oxide of the same kind only. It is
well known that many substances are capable of combining with water.
Oxides possess this property in the highest degree. We have already
seen examples of this (Chapter I.) in the combination of lime, and
of sulphuric and phosphoric anhydrides, with water. The resulting
combinations are basic and acid hydrates. Acid hydrates are called
_acids_ because they have an acid taste when dissolved in water (or
saliva), for then only can they act on the palate. Vinegar, for example,
has an acid taste because it contains acetic acid dissolved in water.
Sulphuric acid, to which we have frequently referred, because it is the
acid of the greatest importance both in practical chemistry and for its
technical applications, is really a hydrate formed by the combination
of sulphuric anhydride with water. Besides their acid taste, dissolved
acids or acid hydrates have the property of changing the blue colour of
certain vegetable dyes to red. Of these dyes _litmus_ is particularly
remarkable and much used. It is the blue substance extracted from
certain lichens, and is used for dyeing tissues blue; it gives a blue
infusion with water. This infusion, on the addition of an acid, _changes
from blue to red_.[50]

  [50] Blotting or unsized paper, soaked in a solution of litmus, is
       usually employed for detecting the presence of acids. This paper
       is cut into strips, and is called _test paper_; when dipped into
       acid it immediately turns red. This is a most sensitive reaction,
       and may be employed for testing for the smallest traces of
       acids. If 10,000 parts by weight of water be mixed with 1 part
       of sulphuric acid, the coloration is distinct, and it is even
       perceptible on the addition of ten times more water. Certain
       precautions must, however, be taken in the preparation of such
       very sensitive litmus paper. Litmus is sold in lumps. Take, say,
       100 grams of it; powder it, and add it to cold pure water in
       a flask; shake and decant the water. Repeat this three times.
       This is done to wash away easily-soluble impurities, especially
       alkalis. Transfer the washed litmus (it is washed with absolute
       alcohol to remove the non-sensitive reddish colouring matter) to
       a flask, and pour in 600 c.c. of water, heat, and allow the hot
       infusion to remain for some hours in a warm place. Then filter,
       and divide the filtrate into two parts. Add a few drops of nitric
       acid to one portion, so that a faint red tinge is obtained, and
       then mix the two portions. Add spirit to the mixture, and keep
       it in a stoppered bottle (it soon spoils if left open to the
       air). This infusion may be employed directly; it reddens in the
       presence of acids, and turns blue in the presence of alkalis. If
       evaporated, a solid mass is obtained which is soluble in water,
       and may be kept unchanged for any length of time. The test paper
       may be prepared as follows:--Take a strong infusion of litmus,
       and soak blotting-paper with it; dry it, and cut it into strips,
       and use it as test-paper for acids. For the detection of alkalis,
       the paper must be soaked in a solution of litmus just reddened by
       a few drops of acid; if too much acid be taken, the paper will
       not be sensitive. Such acids as sulphuric acid colour litmus,
       and especially its infusion, a brick-red colour, whilst more
       feeble acids, such as carbonic, give a faint red-wine tinge.
       Test-paper of a yellow colour is also employed; it is dyed by
       an infusion of turmeric roots in spirit. In alkalis it turns
       brown, but regains its original hue in acids. Many blue and other
       vegetable colouring matters may be used for the detection of
       acids and alkalis; for example, infusions of cochineal, violets,
       log-wood, &c. Certain artificially prepared substances and dyes
       may also be employed. Thus rosolic acid, C_{20}H_{16}O_{3} and
       phenolphthaleïn, C_{20}H_{14}O_{4} (it is used in an alcoholic
       solution, and is not suitable for the detection of ammonia), are
       colourless in an acid, and red in an alkaline, solution. Cyanine
       is also colourless in the presence of acids, and gives a blue
       coloration with alkalis. Methyl-orange (yellow in an aqueous
       solution) is not altered by alkalis but becomes pink with acids
       (weak acids have no action), &c. These are very sensitive tests.
       Their behaviour in respect to various acids, alkalis, and salts
       sometimes give the means of distinguishing substances from each
       other.

Basic oxides, in combining with water, form hydrates, of which, however,
very few are soluble in water. Those which are soluble in water have
an alkaline taste like that of soap or of water in which wood ashes
have been boiled, and are called _alkalis_. Further, alkalis have the
property of restoring the blue colour to litmus which has been reddened
by the action of acids. The hydrates of the oxides of sodium and
potassium, NaHO and KHO, are examples of basic hydrates easily soluble
in water. They are true alkalis, and are termed _caustic_, because they
act very powerfully on the skin of animals and plants. Thus NaHO is
called 'caustic' soda.

The saline oxides are capable of combining together and with water.
Water itself is an oxide, and not an indifferent one, for it can, as we
have seen, combine with basic and acid oxides; it is a representative
of a whole series of saline oxides, _intermediate oxides_, capable of
combining with both basic and acid oxides. There are many such oxides,
which, like water, combine with basic and acid anhydrides--for instance,
the oxides of aluminium and tin, &c. From this it may be concluded that
all oxides might be placed, in respect to their capacity for combining
with one another, in one uninterrupted series, at one extremity of which
would stand those oxides which do not combine with the bases--that
is, the alkalis--while at the other end would be the acid oxides, and
in the interval those oxides which combine with one another and with
both the acid and basic oxides. The further apart the members of this
series are, the more stable are the compounds they form together, the
more energetically do they act on each other, the greater the quantity
of heat evolved in their reaction, and the more marked is their saline
chemical character.

We said above that basic and acid oxides combine together, but rarely
react on each other; this depends on the fact that the majority of them
are solids or gases--that is, they occur in the state least prone
to chemical reaction. The gaseo-elastic state is with difficulty
destroyed, because it necessitates overcoming the elasticity proper
to the gaseous particles. The solid state is characterised by the
immobility of its particles; whilst chemical action requires contact,
and hence a displacement and mobility. If solid oxides be heated, and
especially if they be melted, then reaction proceeds with great ease.
But such a change of state rarely occurs in nature or in practice.
Only in a few furnace processes is this the case. For example; in the
manufacture of glass, the oxides contained in it combine together in
a molten state. But when oxides combine with water, and especially
when they form hydrates soluble in water, then the mobility of their
particles increases to a considerable extent, and their reaction
is greatly facilitated. Reaction then takes place at the ordinary
temperature--easily and rapidly; so that this kind of reaction belongs
to the class of those which take place with unusual facility, and are,
therefore, very often taken advantage of in practice, and also have
been and are going on in nature at every step. We will now consider the
reactions of oxides in the state of hydrates, not losing sight of the
fact that water is itself an oxide with definite properties, and has,
therefore, no little influence on the course of those changes in which
it takes part.

If we take a definite quantity of an acid, and add an infusion of litmus
to it, it turns red; the addition of an alkaline solution does not
immediately alter the red colour of the litmus, but on adding more and
more of the alkaline solution a point is reached when the red colour
changes to violet, and then the further addition of a fresh quantity
of the alkaline solution changes the colour to blue. This change of
the colour of the litmus is a consequence of the formation of a new
compound. This reaction is termed the _saturation_ or _neutralisation_
of the acid by the base, or _vice versâ_. The solution in which the acid
properties of the acid are saturated by the alkaline properties of the
base is termed a _neutral_ solution. Such a solution, although derived
from the mixture of a base with an acid, does not exhibit either the
acid or basic reaction on litmus, yet it preserves many other signs of
the acid and alkali. It is observed that in such a definite admixture
of an acid with an alkali, besides the changes in the colour of litmus
there is a heating effect--_i.e._ an evolution of heat--which is alone
sufficient to prove that there was chemical action. And, indeed, if the
resultant violet solution be evaporated, there separates out, not the
acid or the alkali originally taken, but a substance which has neither
acid nor alkaline properties, but is usually solid and crystalline,
having a saline appearance; this is a _salt_ in the chemical sense of
the word. Hence a salt is derived from the reaction of an acid on an
alkali, in a certain definite proportion. The water here taken for
solution plays no other part than merely facilitating the progress
of the reaction. This is seen from the fact that the anhydrides of
the acids are able to combine with basic oxides, and give the same
salts as do the acids with the alkalis or hydrates. Hence, a salt is a
compound of definite quantities of an acid with an alkali. In the latter
reaction, water is separated out if the substance formed be the same
as is produced by the combination of anhydrous oxides together.[51]
Examples of the formation of salts from acids and bases are easily
observed, and are very often applied in practice. If we take, for
instance, insoluble magnesium oxide (magnesia) it is easily dissolved
in sulphuric acid, and on evaporation gives a saline substance, bitter,
like all the salts of magnesium, and familiar to all under the name of
Epsom salts, used as a purgative. If a solution of caustic soda--which
is obtained, as we saw, by the action of water on sodium oxide--be
poured into a flask in which charcoal has been burnt; or if carbonic
anhydride, which is produced under so many circumstances, be passed
through a solution of caustic soda, then sodium carbonate or soda,
Na_{2}CO_{3}, is obtained, of which we have spoken several times, and
which is prepared on a large scale and often used in manufactures. This
reaction is expressed by the equation, 2NaHO + CO_{2} = Na_{2}CO_{3} +
H_{2}O. Thus, the various bases and acids form an innumerable number of
different salts.[52] Salts constitute an example of definite chemical
compounds, and both in the history and practice of science are most
often cited as confirming the conception of definite chemical compounds.
Indeed, all the indications of a definite chemical combination are
clearly seen in the formation and properties of salts. Thus, they are
produced with a definite proportion of oxides, heat is evolved in their
formation,[53] and the chemical character of the oxides and many of the
physical properties become hidden in their salts. For example, when
gaseous carbonic anhydride combines with a base to form a solid salt,
the elasticity of the gas quite disappears in its passage into the
salt.[54]

  [51] That water really is separated in the reaction of acid on alkaline
       hydrates, may be shown by taking some other intermediate
       hydrate--for example, alumina--instead of water. Thus, if a
       solution of alumina in sulphuric acid be taken, it will have,
       like the acid, an acid reaction, and will therefore colour
       litmus red. If, on the other hand, a solution of alumina in an
       alkali--say, potash--be taken, it will have an alkaline reaction,
       and will turn red litmus blue. On adding the alkaline to the
       acid solution until neither an alkaline nor an acid reaction is
       produced, a salt is formed, consisting of sulphuric anhydride
       and potassium oxide. In this, as in the reaction of hydrates,
       an intermediate oxide is separated out--namely, alumina. Its
       separation will be very evident in this case, as alumina is
       insoluble in water.

  [52] The mutual interaction of hydrates, and their capacity of forming
       salts, may be taken advantage of for determining the character
       of those hydrates which are insoluble in water. Let us imagine
       that a given hydrate, whose chemical character is unknown, is
       insoluble in water. It is therefore impossible to test its
       reaction on litmus. It is then mixed with water, and an acid--for
       instance, sulphuric acid--is added to the mixture. If the hydrate
       taken be basic, reaction will take place, either directly or by
       the aid of heat, with the formation of a salt. In certain cases,
       the resultant salt is soluble in water, and this will at once
       show that combination has taken place between the insoluble basic
       hydrate and the acid, with the formation of a soluble saline
       substance. In those cases where the resultant salt is insoluble,
       still the water loses its acid reaction, and therefore it may he
       ascertained, by the addition of an acid, whether a given hydrate
       has a basic character, like the hydrates of oxide of copper,
       lead, &c. If the acid does not act on the given insoluble hydrate
       (at any temperature), then it has not a basic character, and it
       should be tested as to whether it has an acid character. This is
       done by taking an alkali, instead of the acid, and by observing
       whether the unknown hydrate then dissolves, or whether the
       alkaline reaction disappears. Thus it may he proved that hydrate
       of silica is acid, because it dissolves in alkalis and not in
       acids. If it be a case of an insoluble intermediate hydrate, then
       it will be observed to react on both the acid and alkali. Hydrate
       of alumina is an instance in question, which is soluble both in
       caustic potash and in sulphuric acid.

       The _degree of affinity_ or chemical _energy_ proper to oxides
       and their hydrates is very dissimilar; some extreme members
       of the series possess it to a great extent. When acting on
       each other they evolve a large quantity of heat, and when
       acting on intermediate hydrates they also evolve heat to a
       considerable degree, as we saw in the combination of lime and
       sulphuric anhydride with water. When extreme oxides combine
       they form stable salts, which are decomposed with difficulty,
       and often show characteristic properties. The compounds of the
       intermediate oxides with each other, or even with basic and acid
       oxides, present a very different case. However much alumina we
       may dissolve in sulphuric acid, we cannot saturate the acid
       properties of the sulphuric acid, the resulting solution will
       always have an acid reaction. So also, whatever quantity of
       alumina is dissolved in an alkali, the resulting solution will
       always present an alkaline reaction.

  [53] In order to give an idea of the quantity of heat evolved in the
       formation of salts I append a table of data for _very dilute
       aqueous solutions_ of acids and alkalis, according to the
       determinations of Berthelot and Thomsen. The figures are given
       in major calories--that is, in thousands of units of heat. For
       example, 49 grams of sulphuric acid, H_{2}SO_{4}, taken in a
       dilute aqueous solution, when mixed with such an amount of a weak
       solution of caustic soda, NaHO, that a neutral salt is formed
       (when all the hydrogen of the acid is replaced by the sodium),
       evolves 15,800 units of heat.

                              49 parts of      63 parts of
                              H_{2}SO_{4}        HNO_{3}

                   NaHO           15·8            13·7
                   KHO            15·7            13·8
                   NH_{3}         14·5            12·5
                   CaO            15·6            13·9
                   BaO            18·4            13·9
                   MgO            15·6            13·8
                   FeO            12·5            10·7 (?)
                   ZnO            11·7             9·8
                   Fe_{2}O_{3}     5·7             5·9

       These figures cannot be considered as the heat of neutralisation,
       because the water here plays an important part. Thus, for
       instance, sulphuric acid and caustic soda in dissolving in water
       evolve very much heat, and the resultant sodium sulphate very
       little; consequently, the amount of heat evolved in an anhydrous
       combination will be different from that evolved in a hydrated
       combination. Those acids which are not energetic in combining
       with the same quantity of alkalis required for the formation of
       normal salts of sulphuric or nitric acids always, however, give
       less heat. For instance, with caustic soda: carbonic acid gives
       10·2, hydrocyanic, 2·9, hydrogen sulphide, 3·9 major calories.
       And as feeble bases (for example, Fe_{2}O_{3}) also evolve less
       heat than those which are more powerful, so a certain general
       correlation between thermochemical data and the degree of
       affinity shows itself here, as in other cases (_see_ Chapter II.,
       Note 7); this does not, however, give any reason for measuring
       the affinity which binds the elements of salts by the heat
       of their formation in dilute solutions. This is very clearly
       demonstrated by the fact that water is able to decompose many
       salts, and is separated in their formation.

  [54] Carbonic anhydride evolves heat in dissolving in water. The
       solution easily dissociates and evolves carbonic anhydride,
       according to the law of Henry and Dalton (_see_ Chapter I.)
       In dissolving in caustic soda, it either gives a normal salt,
       Na_{2}CO_{3}, which does not evolve carbonic anhydride, or an
       acid salt, NaHCO_{3} which easily evolves carbonic anhydride when
       heated. The same gas, when dissolved in solutions of salts, acts
       in one or the other manner (_see_ Chapter II., Note 38). Here
       it is seen what a successive series of relations exists between
       compounds of a different order, between substances of different
       degrees of stability. By making a distinction between the
       phenomena of solutions and chemical compounds, we overlook those
       natural transitions which in reality exist.

Judging from the above, a salt is a compound of basic and acid oxides,
or the result of the action of hydrates of these classes on each other
with separation of water. But salts may be obtained by other methods. It
must not be forgotten that basic oxides are formed by metals, and acid
oxides usually by non-metals. But metals and non-metals are capable of
combining together, and a salt is frequently formed by the oxidation of
such a compound. For example, iron very easily combines with sulphur,
forming iron sulphide FeS (as we saw in the Introduction); this in
air, and especially moist air, absorbs oxygen, with the formation of
the same salt FeSO_{4}, that may be obtained by the combination of the
oxides of iron and sulphur, or of the hydrates of these oxides. Hence,
it cannot be said or supposed that a salt has the properties of the
oxides, or must necessarily contain two kinds of oxides in itself.
The derivation of salts from oxides is merely one of the methods of
their preparation. We saw, for instance, that in sulphuric acid it was
possible to replace the hydrogen by zinc, and that by this means zinc
sulphate was formed; so likewise the hydrogen in many other acids may be
replaced by zinc, iron, potassium, sodium, and a whole series of similar
metals, corresponding salts being obtained. The hydrogen of the acid,
in all these cases, is exchanged for a metal, and a salt is obtained
from the hydrate. Regarding a salt from this point of view, it may be
said that _a salt is an acid in which hydrogen is replaced by a metal_.
This definition shows that a salt and an acid are essentially compounds
of the same series, with the difference that the latter contains
hydrogen and the former a metal. Such a definition is more exact than
the first definition of salts, inasmuch as it likewise includes those
acids which do not contain oxygen, and, as we shall afterwards learn,
there is a series of such acids. Such elements as chlorine and bromine
form compounds with hydrogen in which the hydrogen may be replaced by
a metal, forming substances which, in their reactions and external
characters, resemble the salts formed from oxides. Table salt, NaCl, is
an example of this. It may be obtained by the replacement of hydrogen
in hydrochloric acid, HCl, by the metal sodium, just as sulphate of
sodium, Na_{2}SO_{4}, may be obtained by the replacement of hydrogen in
sulphuric acid, H_{2}SO_{4}, by sodium. The exterior appearance of the
resulting products, their neutral reaction, and even their saline taste,
show their resemblance to one another.

To the fundamental properties of salts yet another must be
added--namely, that they are more or less _decomposed by the action of a
galvanic current_. The results of this decomposition are very different
according to whether the salt be taken in a fused or dissolved state.
But the decomposition may generally be so represented, that the metal
appears at the electro-negative pole or cathode (like hydrogen in the
decomposition of water, or its mixture with sulphuric acid), and the
remaining parts of the salt appear at the electro-positive pole or anode
(where the oxygen of water appears). If, for instance, an electric
current acts on an aqueous solution of sodium sulphate, then the sodium
appears at the negative pole, and oxygen and the anhydride of sulphuric
acid at the positive pole. But in the solution itself the result is
different, for sodium, as we know, decomposes water with evolution of
hydrogen, forming caustic soda; consequently hydrogen will be evolved,
and caustic soda appear at the negative pole: while at the positive
pole the sulphuric anhydride immediately combines with water and forms
sulphuric acid, and therefore oxygen will be evolved and sulphuric acid
formed round this pole.[55] In other cases, when the metal separated is
not able to decompose water, it will be deposited in a free state. Thus,
for example, in the decomposition of copper sulphate, copper separates
out at the cathode, and oxygen and sulphuric acid appear at the anode,
and if a copper plate be attached to the positive pole, then the oxygen
evolved will oxidise the copper, and the oxide of copper will dissolve
and be deposited at the negative pole--that is, a transfer of copper
from the positive to the negative pole ensues. The galvanoplastic art
(electro-typing) is based on this principle.[56] Therefore the most
radical and general properties of salts (including also such salts as
table salt, which contain no oxygen) may be expressed by representing
the salt as composed of a metal M and a haloid X--that is, by expressing
the salt by MX. In common table salt the metal is sodium, and the haloid
an elementary body, chlorine. In sodium sulphate, Na_{2}SO_{4}, sodium
is again the metal, but the complex group, SO_{4}, is the haloid. In
sulphate of copper, CuSO_{4}, the metal is copper and the haloid the
same as in the preceding salt. Such a representation of salts expresses
with great simplicity the _capacity of every salt to enter into saline
double decompositions with other salts_; consisting in the mutual
replacement of the metals in the salts. This exchange of their metals
is the fundamental property of salts. In the case of two salts with
different metals and haloids, which are in solution or fusion, or in
any other manner brought into contact, the metals of these salts will
always partially or wholly exchange places. If we designate one salt
by MX, and the other by NY, then we either partially or wholly obtain
from them new salts, MY and NX. Thus we saw in the Introduction, that
on mixing solutions of table salt, NaCl, and silver nitrate, AgNO_{3},
a white insoluble precipitate of silver chloride, AgCl, is formed and
a new salt, sodium nitrate, NaNO_{3}, is obtained in solution. If the
metals of salts exchange places in reactions of double decomposition,
it is clear that metals themselves, taken in a separate state, are
able to act on salts, as zinc evolves hydrogen from acids, and as iron
separates copper from copper sulphate. When, to what extent, and which
metals displace each other, and how the metals are distributed between
the haloids, will be discussed in Chapter X., where we shall be guided
by those reflections and deductions which Berthollet introduced into the
science at the beginning of this century.

  [55] This kind of decomposition may be easily observed by pouring
       a solution of sodium sulphate into a U-shaped tube and
       inserting electrodes in the two branches. If the solution be
        with an infusion of litmus, it will easily be seen
       that it turns blue at the cathode, owing to the formation of
       sodium hydroxide, and red at the electro-positive pole, from the
       formation of sulphuric acid.

  [56] In other cases the decomposition of salts by the electric current
       may be accompanied by much more complex results. Thus, when the
       metal of the salt is capable of a higher degree of oxidation,
       such a higher oxide may be formed at the positive pole by the
       oxygen which is evolved there. This takes place, for instance,
       in the decomposition of salts of silver and manganese by the
       galvanic current, peroxides of these metals being formed. Thus in
       the electrolysis of a solution of KCl, KClO_{3} is formed, and
       of sulphuric acid (corresponding to SO_{3}) persulphuric acid,
       corresponding to S_{2}O_{7}. But all the phenomena as yet known
       may be expressed by the above law--that the current decomposes
       salts into metals, which appear at the negative pole, and into
       the remaining component parts, which appear at the positive pole.

According to the above observations, an acid is nothing more than a
salt of hydrogen. Water itself may be looked on as a salt in which
the hydrogen is combined with either oxygen or the aqueous radicle,
OH; water will then be HOH, and alkalis or basic hydrates, MOH. The
group OH, or the _aqueous radicle_, otherwise called _hydroxyl_, may
be looked on as a haloid like the chlorine in table salt, not only
because the element Cl and the group OH very often change places, and
combine with one and the same element, but also because free chlorine is
very similar in many properties and reactions to peroxide of hydrogen,
which is the same in composition as the aqueous radicle, as we shall
afterwards see in Chapter IV. Alkalis and basic hydrates are also
salts consisting of a metal and hydroxyl--for instance, caustic soda,
NaOH; this is therefore termed _sodium hydroxide_. According to this
view, _acid salts_ are those in which a portion only of the hydrogen is
replaced by a metal, and a portion of the hydrogen of the acid remains.
Thus sulphuric acid (H_{2}SO_{4}) not only gives the normal salt
Na_{2}SO_{4}, with sodium, but also an acid salt, NaHSO_{4}. A _basic
salt_ is one in which the metal is combined not only with the haloids
of acids, but also with the aqueous radicale of basic hydrates--for
example, bismuth gives not only a normal salt of nitric acid,
Bi(NO_{3})_{3}, but also basic salts like Bi(OH)_{2}(NO_{3}).

As basic and acid salts of the oxygen acids contain hydrogen and oxygen,
they are able to part with these as water and to give anhydro-salts,
which it is evident will be compounds of normal salts with anhydrides
of the acids or with bases. Thus the above-mentioned acid sodium
sulphate corresponds with the anhydro-salt, Na_{2}S_{2}O_{7}, equal
to 2NaHSO_{4}, less H_{2}O. The loss of water is here, and frequently
in other cases, brought about by heat alone, and therefore such salts
are frequently termed _pyro-salts_--for instance, the preceding is
sodium pyrosulphate (Na_{2}S_{2}O_{7}), or it may be regarded as the
normal salt Na_{2}SO_{4} + sulphuric anhydride, SO_{3}. _Double_ salts
are those which contain either two metals, KAl(SO_{4})_{2}, or two
haloids.[57]

  [57] The above-enunciated generalisation of the conception of salts as
       compounds of the metals (simple, or compound like ammonium,
       NH_{4}), with the haloids (simple, like chlorine, or compound,
       like cyanogen, CN, or the radical of sulphuric acid, SO_{4}),
       capable of entering into double saline decomposition, which
       is in accordance with the general data respecting salts, was
       only formed little by little after a succession of most varied
       propositions as to the chemical structure of salts.

       Salts belong to the class of substances which have been known
       since very early times, and have long been investigated in many
       directions. At first, however, no distinction was made between
       salts, acids, and bases. Glauber prepared many artificial salts
       during the latter half of the seventeenth century. Up to that
       time the majority of salts were obtained from natural sources,
       and that salt which we have referred to several times--namely,
       sodium sulphate--was named Glauber's salt after this chemist.
       Rouelle distinguished normal, acid, and basic salts, and showed
       their action on vegetable dyes, still he confounded many salts
       with acids (even now every acid salt ought to be regarded as an
       acid, because it contains hydrogen, which may be replaced by
       metals--that is, it is the hydrogen of an acid). Baumé disputed
       Rouelle's opinion concerning the subdivision of salts, contending
       that normal salts only are true salts, and that basic salts are
       simple mixtures of normal salts with bases and acid salts with
       acids, considering that washing alone could remove the base or
       acid from them. Rouelle, in the middle of the last century,
       however, rendered a great service to the study of salts and the
       diffusion of knowledge respecting this class of compounds in
       his attractive lectures. He, like the majority of the chemists
       of that period, did not employ the balance in his researches,
       but satisfied himself with purely qualitative data. The first
       quantitative researches on salts were carried on about this
       time by Wenzel, who was the director of the Freiburg mines, in
       Saxony. Wenzel studied the double decomposition of salts, and
       observed that in the double decomposition of neutral salts a
       neutral salt was always obtained. He proved, by a method of
       weighing, that this is due to the fact that the saturation of a
       given quantity of a base requires such relative quantities of
       different acids as are capable of saturating every other base.
       Having taken two neutral salts--for example, sodium sulphate and
       calcium nitrate--let us mix their solutions together. Double
       decomposition takes place, because calcium sulphate is formed,
       which is almost insoluble. However much we might add of each
       of the salts, the neutral reaction will still be preserved,
       consequently the neutral character of the salts is not destroyed
       by the interchange of metals; that is to say, that quantity of
       sulphuric acid which saturated the sodium is sufficient for the
       saturation of the calcium, and that amount of nitric acid which
       saturated the calcium is enough to saturate the sodium contained
       in combination with sulphuric acid in sodium sulphate. Wenzel
       was even convinced that matter does not disappear in nature, and
       on this principle he corrects, in his _Doctrine of Affinity_,
       the results of his experiments when he found that he obtained
       less than he had originally taken. Although Wenzel deduced the
       law of the double decomposition of salts quite correctly, he did
       not determine those quantities in which acids and bases act on
       each other. This was carried out at the end of the last century
       by Richter. He determined the quantities by weight of the bases
       which saturate acids and of the acids which saturate bases, and
       obtained comparatively correct results, although his conclusions
       were not correct, for he states that the quantity of a base
       saturating a given acid varies in arithmetical progression, and
       the quantity of an acid saturating a given base in geometrical
       progression. Richter studied the deposition of metals from their
       salts by other metals, and observed that the neutral reaction
       of the solution is not destroyed by this exchange. He also
       determined the quantities by weight of the metals replacing one
       another in salts. He showed that copper displaces silver from
       its salt, and that zinc displaces copper and a whole series of
       other metals. Those quantities of metals which were capable of
       replacing one another were termed equivalents.

       Richter's teaching found no followers, because, although he
       fully believed in the discoveries of Lavoisier, yet he still
       held to the phlogistic reasonings which rendered his expositions
       very obscure. The works of the Swedish savant Berzelius freed
       the facts discovered by Wenzel and Richter from the obscurity
       of former conceptions, and led to their being explained in
       accordance with Lavoisier's views, and in the sense of the law of
       multiple proportions which had already been discovered by Dalton.
       On applying to salts those conclusions which Berzelius arrived at
       by a whole series of researches of remarkable accuracy, we arrive
       at the following law of equivalents--_one part by weight of
       hydrogen in an acid is replaced by the corresponding equivalent
       weight of any metal_; and, therefore, when metals replace each
       other their weights are in the same ratio as their equivalents.
       Thus, for instance, one part by weight of hydrogen is replaced by
       23 parts of sodium, 39 parts of potassium, 12 parts of magnesium,
       20 parts of calcium, 28 parts of iron, 108 parts of silver, 33
       parts of zinc, &c.; and, therefore, if zinc replaces silver, then
       33 parts of zinc will take the place of 108 parts of silver, or
       33 parts of zinc will he substituted by 23 parts of sodium, &c.

       The doctrine of equivalents would be precise and simple did
       every metal only give one oxide or one salt. It is rendered
       complicated from the fact that many metals form several oxides,
       and consequently offer different equivalents in their different
       degrees of oxidation. For example, there are oxides containing
       iron in which its equivalent is 28--this is in the salts formed
       by the suboxide; and there is another series of salts in which
       the equivalent of iron equals 18-2/3--which contain less iron,
       and consequently more oxygen, and correspond with a higher degree
       of oxidation--ferric oxide. It is true that the former salts are
       easily formed by the direct action of metallic iron on acids, and
       the latter only by a further oxidation of the compound formed
       already; but this is not always so. In the case of copper,
       mercury, and tin, under different circumstances, salts are formed
       which correspond with different degrees of oxidation of these
       metals, and many metals have two equivalents in their different
       salts--that is, in salts corresponding with the different degrees
       of oxidation. Thus it is impossible to endow every metal with one
       definite equivalent weight. Hence the conception of equivalents,
       while playing an important part from an historical point of view,
       appears, with a fuller study of chemistry, to be but subordinate
       to a higher conception, with which we shall afterwards become
       acquainted.

       The fate of the theoretical views of chemistry was for a
       long time bound up with the history of salts. The clearest
       representation of this subject dates back to Lavoisier, and was
       systematically developed by Berzelius. This representation is
       called the _binary_ theory. All compounds, and especially salts,
       are represented as consisting of two parts. Salts are represented
       as compounds of a basic oxide (a base) and an acid (that is, an
       anhydride of an acid, then termed an acid), whilst hydrates are
       represented as compounds of anhydrous oxides with water. Such
       an expression was employed not only to denote the most usual
       method of formation of these substances (where it would be quite
       true), but also to express that internal distribution of the
       elements by which it was proposed to explain all the properties
       of these substances. Copper sulphate was supposed to contain
       two most intimate component parts--copper oxide and sulphuric
       anhydride. This is an hypothesis. It arose from the so-called
       _electro-chemical hypothesis_, which supposed the two component
       parts to be held in mutual union, because one component (the
       anhydride of the acid) has electro-negative properties, and the
       other (the base in salts) electro-positive. The two parts are
       attracted together, like substances having opposite electrical
       charges. But as the decomposition of salts in a state of fusion
       by an electric current always gives a metal, that representation
       of the constitution and decomposition of salts called the
       _hydrogen theory_ of acids is nearer the truth than that which
       considers salts as made up of a base and an anhydride of an acid.
       But the hydrogen theory of acids is also a binary hypothesis,
       and does not contradict the electro-chemical hypothesis, but
       is rather a modification of it. The binary theory dates from
       Rouelle and Lavoisier, the electro-chemical aspect was zealously
       developed by Berzelius, and the hydrogen theory of acids is due
       to Davy and Liebig.

       These hypothetical views simplified and generalised the study of
       a complicated subject, and served to support further arguments,
       but when salts were in question it was equally convenient to
       follow one or the other of these hypotheses. But these theories
       were brought to bear on all other substances, on all compound
       substances. Those holding the binary and electro-chemical
       hypotheses searched for two anti-polar component parts, and
       endeavoured to express the process of chemical reactions by
       electro-chemical and similar differences. If zinc replaces
       hydrogen, they concluded that it is more electro-positive
       than hydrogen, whilst they forgot that hydrogen may, under
       different circumstances, displace zinc--for instance, at a
       red heat. Chlorine and oxygen were considered as being of
       opposite polarity to hydrogen because they easily combine with
       it, nevertheless both are capable of replacing hydrogen, and,
       what is very characteristic, in the replacement of hydrogen
       by chlorine in carbon compounds not only does the chemical
       character often remain unaltered, but even the external form
       may remain unchanged, as Laurent and Dumas demonstrated. These
       considerations undermine the binary, and more especially the
       electro-chemical theory. An explanation of known reactions then
       began to be sought for not in the difference of the polarity of
       the different substances, but in the joint influences of all the
       elements on the properties of the compound formed. This is the
       reverse of the preceding hypothesis.

       This reversal was not, however, limited to the destruction of
       the tottering foundations of the preceding theory; it proposed a
       new doctrine, and laid the foundation for the modern course of
       our science. This doctrine may be termed the unitary theory--that
       is, it strictly acknowledges the joint influences of the elements
       in a compound substance, denies the existence of separate and
       contrary components in them, regards copper sulphate, for
       instance, as a strictly definite compound of copper, sulphur,
       and oxygen; then seeks for compounds which are analogous in
       their properties, and, placing them side by side, endeavours
       to express the influence of each element in determining the
       united properties of its compound. In the majority of cases it
       arrives at conclusions similar to those which are obtained by
       the above-mentioned hypotheses, but in certain special cases the
       conclusions of the unitary theory are in entire opposition to
       those of the binary theory and its corollaries. Cases of this
       kind are most often met with in the consideration of compounds of
       a more complex nature than salts, especially organic compounds
       containing hydrogen. But it is not in this change from an
       artificial to a natural system, important as it is, that the
       chief service and strength of the unitary doctrine lies. By a
       simple review of the vast store of data regarding the reactions
       of typical substances, it succeeded from its first appearance
       in establishing a new and important law, it introduced a new
       conception into science--namely, the conception of molecules,
       with which we shall soon become acquainted. The deduction of
       the law and of the conception of molecules has been verified by
       facts in a number of cases, and was the cause of the majority of
       chemists of our times deserting the binary theory and accepting
       the unitary theory, which forms the basis of the present work.
       Laurent and Gerhardt must be considered as the founders of this
       doctrine.

Inasmuch as oxygen compounds predominate in nature, it should be
expected from what has been said above, that salts, rather than acids
or bases, would occur most frequently in nature, for these latter would
always tend to combine forming salts, especially through the medium
of the all-pervading water. And, as a matter of fact, salts are found
everywhere in nature. They occur in animals and plants, although in
but small quantity, because, as forming the last stage of chemical
reaction, they are capable of only a few chemical transformations. And
organisms are bodies in which a series of uninterrupted, varied, and
active chemical transformations proceed, whilst salts, which only enter
into double decompositions between each other, are little prone to such
changes. But organisms always contain salts. Thus, for instance, bones
contain calcium phosphate, the juice of grapes potassium tartrate (cream
of tartar), certain lichens calcium oxalate, and the shells of mollusca
calcium carbonate, &c. As regards water and soil, portions of the earth
in which the chemical processes are less active, they are full of
salts. Thus the waters of the oceans, and all others (Chap. I.), abound
in salts, and in the soil, in the rocks of the earth's crust, in the
upheaved lavas, and in the falling meteorites the salts of silicic acid,
and especially its double salts, predominate. Saline substances also
make up the composition of those limestones which often form mountain
chains and whole thicknesses of the earth's strata, these consisting of
calcium carbonate, CaCO_{3}.

Thus we have seen oxygen in a free state and in various compounds
of different degrees of stability, from the unstable salts, like
Berthollet's salt and nitre, to the most stable silicon compounds, such
as exist in granite. We saw an entirely similar gradation of stability
in the compounds of water and of hydrogen. In all its aspects oxygen,
as an element, or single substance, remains the same however varied
its chemical states, just as a substance may appear in many different
physical states of aggregation. But our notion of the immense variety of
the chemical states in which oxygen can occur would not be completely
understood if we did not make ourselves acquainted with it in the form
in which it occurs in ozone and peroxide of hydrogen. In these it is
most active, its energy seems to have increased. They illustrate fresh
aspects of chemical correlations, and the variety of the forms in which
matter can appear stand out clearly. We will therefore consider these
two substances somewhat in detail.




                               CHAPTER IV

                OZONE AND HYDROGEN PEROXIDE--DALTON'S LAW


VAN MARUM, during the last century, observed that oxygen in a glass
tube, when subjected to the action of a series of electric sparks,
acquired a peculiar smell, and the property of combining with mercury
at the ordinary temperature. This was afterwards confirmed by a number
of fresh experiments. Even in the simple revolution of an electrical
machine, when electricity diffuses into the air or passes through
it, the peculiar and characteristic smell of ozone, proceeding from
the action of the electricity on the oxygen of the atmosphere, is
recognised. In 1840 Prof. Schönbein, of Basle, turned his attention to
this odoriferous substance, and showed that it is also formed, with the
oxygen evolved at the positive pole, in the decomposition of water by
the action of a galvanic current; in the oxidation of phosphorus in damp
air, and also in the oxidation of a number of substances, although it
is distinguished for its instability and capacity for oxidising other
substances. The characteristic smell of this substance gave it its name,
from the Greek [Greek: ozô], 'I emit an odour.' Schönbein pointed out
that _ozone_ is capable of oxidising many substances on which oxygen
does not act at the ordinary temperature. It will be sufficient to point
out for instance that it oxidises silver, mercury, charcoal, and iron
with great energy at the ordinary temperature. It might be thought that
ozone was some new compound substance, as it was at first supposed to
be; but careful observations made in this direction have long led to the
conclusion that ozone is nothing but oxygen altered in its properties.
This is most strikingly proved by the complete transformation of oxygen
containing ozone into ordinary oxygen when it is passed through a tube
heated to 250°. Further, at a low temperature pure oxygen gives ozone
when electric sparks are passed through it (Marignac and De la Rive).
Hence it is proved both by synthesis and analysis that ozone is that
same oxygen with which we are already acquainted, only endowed with
particular properties and in a particular state. However, by whatever
method it be obtained, the amount of it contained in the oxygen is
inconsiderable, generally only a few fractions per cent., rarely 2 per
cent., and only under very propitious circumstances as much as 20 per
cent. The reason of this must be looked for first in the fact that
_ozone in its formation from oxygen absorbs heat_. If any substance
be burnt in a calorimeter at the expense of ozonised oxygen, then
more heat is evolved than when it is burnt in ordinary oxygen, and
Berthelot showed that this difference is very large--namely, 29,600
heat units correspond with every forty-eight parts by weight of ozone.
This signifies that the transformation of forty-eight parts of oxygen
into ozone is accompanied by the absorption of this quantity of heat,
and that the reverse process evolves this quantity of heat. Therefore
the passage of ozone into oxygen should take place easily and fully
(as an exothermal reaction), like combustion; and this is proved by
the fact that at 250° ozone entirely disappears, forming oxygen. Any
rise of temperature may thus bring about the breaking up of ozone, and
as a rise of temperature takes place in the action of an electrical
discharge, there are in an electric discharge the conditions both for
the preparation of ozone and for its destruction. Hence it is clear
that the transformation of oxygen into ozone _as a reversible reaction_
has a limit when a state of equilibrium is arrived at between the
products of the two opposite reactions, that the phenomena of this
transformation accord with the phenomena of _dissociation_, and that
a fall of temperature should aid the formation of a large quantity of
ozone.[1] Further, it is evident, from what has been said, that the best
way of preparing ozone is not by electric sparks,[2] which raise the
temperature, but by the employment of a continual discharge or flow of
electricity--that is, by the action of a _silent discharge_.[3] For this
reason all _ozonisers_ (which are of most varied construction), or
forms of apparatus for the preparation of ozone from oxygen (or air) by
the action of electricity, now usually consist of sheets of metal--for
instance, tinfoil--a solution of sulphuric acid mixed with chromic acid,
&c. separated by thin glass surfaces placed at short distances from each
other, and between which the oxygen or air to be ozonised is introduced
and subjected to the action of a silent discharge.[4] Thus in Siemens'
apparatus (fig. 37) the exterior of the tube _a_ and the interior of
the tube _b c_ are coated with tinfoil and connected with the poles of
a source of electricity (with the terminals of a Ruhmkorff's coil). A
silent discharge passes through the thin walls of the glass cylinders
_a_ and _b c_ over all their surfaces, and consequently, if oxygen be
passed through the apparatus by the tube _d_, fused into the side of
_a_, it will be ozonised in the annular space between _a_ and _b c_. The
ozonised oxygen escapes by the tube _e_, and may be introduced into any
other apparatus.[5]

[Illustration: FIG. 37.--Siemens' apparatus for preparing ozone by means
of a silent discharge.]

  [1] This conclusion, deduced by me as far back as 1878 (_Moniteur
      Scientifique_) by conceiving the molecules of ozone (see later)
      as more complex than those of oxygen, and ozone as containing
      a greater quantity of heat than oxygen, has been proved
      experimentally by the researches of Mailfert (1880), who showed
      that the passage of a silent discharge through a litre of oxygen
      at 0° may form up to 14 milligrams of ozone, and at -30° up to
      60 milligrams; but best of all in the determinations of Chappuis
      and Hautefeuille (1880), who found that at a temperature of -25° a
      silent discharge converted 20 p.c. of oxygen into ozone, whilst
      at 20° it was impossible to obtain more than 12 p.c., and at 100°
      less than 2 p.c. of ozone was obtained.

  [2] A series of electric sparks may be obtained by an ordinary
      electrical machine, the electrophorus machines of Holtz and
      Teploff, &c., Leyden jars, Ruhmkorff coils, or similar means, when
      the opposite electricities are able to accumulate at the terminals
      of conductors, and a discharge of sufficient electrical intensity
      passes through the non-conductors air or oxygen.

  [3] A silent discharge is such a combination of opposite statical
      (potential) electricities as takes place (generally between large
      surfaces) regularly, without sparks, slowly, and quietly (as in
      the dispersion of electricity). The discharge is only luminous
      in the dark; there is no observable rise of temperature, and
      therefore a larger amount of ozone is formed. But, nevertheless,
      on continuing the passage of a silent discharge through ozone it
      is destroyed. For the action to be observable a large surface is
      necessary, and consequently a source of electricity at a high
      potential. For this reason the silent discharge is best produced
      by a Ruhmkorff coil, as the most convenient means of obtaining a
      considerable potential of statical electricity with the employment
      of the comparatively feeble current of a galvanic battery.

  [4] _v. Babo's apparatus_ was one of the first constructed for
      ozonising oxygen by means of a silent discharge (and it is still
      one of the best). It is composed of a number (twenty and more) of
      long, thin capillary glass tubes closed at one end. A platinum
      wire, extending along their whole length, is introduced into the
      other end of each tube, and this end is then fused up round the
      wire, the end of which protrudes outside the tube. The protruding
      ends of the wires are arranged alternately in two sides in such a
      manner that on one side there are ten closed ends and ten wires.
      A bunch of such tubes (forty should make a bunch of not more than
      1 c.m. diameter) is placed in a glass tube, and the ends of the
      wires are connected with two conductors, and are fused to the ends
      of the surrounding tube. The discharge of a Ruhmkorff coil is
      passed through these ends of the wires, and the dry air or oxygen
      to be ozonised is passed through the tube. If oxygen be passed
      through, ozone is obtained in large quantities, and free from
      oxides of nitrogen, which are partially formed when air is acted
      on. At low temperatures ozone is formed in large quantities. As
      ozone acts on corks and india-rubber, the apparatus should be made
      entirely of glass. With a powerful Ruhmkorff coil and forty tubes
      the ozonation is so powerful that the gas when passed through a
      solution of iodide of potassium not only sets the iodine free, but
      even oxidises it to potassium iodate, so that in five minutes the
      gas-conducting tube is choked up with crystals of the insoluble
      iodate.

  [5] In order to connect the ozoniser with any other apparatus it is
      impossible to make use of india-rubber, mercury, or cements, &c.,
      because they are themselves acted on by, and act on, ozone. All
      connections must, as was first proposed by Brodie, be hermetically
      closed by sulphuric acid, which is not acted on by ozone. Thus, a
      cork is passed over the vertical end of a tube, over which a wide
      tube passes so that the end of the first tube protrudes above the
      cork; mercury is first poured over the cork (to prevent its being
      acted on by the sulphuric acid), and then sulphuric acid is poured
      over the mercury. The protruding end of the first tube is covered
      by the lower end of a third tube immersed in the sulphuric acid.

_The properties of ozone_ obtained by such a method[6] distinguish it
in many respects from oxygen. Ozone very rapidly decolorises indigo,
litmus, and many other dyes by oxidising them. Silver is oxidised by it
at the ordinary temperature, whilst oxygen is not able to oxidise silver
even at high temperatures; a bright silver plate rapidly turns black
(from oxidation) in ozonised oxygen. It is rapidly absorbed by mercury,
forming oxide; it transforms the lower oxides into higher--for instance,
sulphurous anhydride into sulphuric, nitrous oxide into nitric,
arsenious anhydride (As_{2}O_{3}) into arsenic anhydride (As_{2}O_{5})
&c.[7] But what is especially characteristic in ozone is the decomposing
action it exerts on potassium iodide. Oxygen does not act on it, but
ozone passed into a solution of potassium iodide _liberates iodine_,
whilst the potassium is obtained as caustic potash, which remains in
solution, 2KI + H_{2}O + O = 2KHO + I_{2}. As the presence of minute
traces of free iodine may be discovered by means of starch paste,
with which it forms a very dark blue- substance, a mixture of
potassium iodide with starch paste will detect the presence of very
small traces of ozone.[8] Ozone is destroyed or converted into ordinary
oxygen not only by heat, but also by long keeping, especially in the
presence of alkalis, peroxide of manganese, chlorine, &c.

  [6] The method above described is the only one which has been well
      investigated. The admixture of nitrogen, or even of hydrogen,
      and especially of silicon fluoride, appears to aid the formation
      and preservation of ozone. Amongst other methods for preparing
      ozone we may mention the following: 1. In the action of oxygen on
      phosphorus at the ordinary temperature a portion of the oxygen
      is converted into ozone. At the ordinary temperature a stick of
      phosphorus, partially immersed in water and partially in air in
      a large glass vessel, causes the air to acquire the odour of
      ozone. It must further be remarked that if the air be left for
      long in contact with the phosphorus, or without the presence of
      water, the ozone formed is destroyed by the phosphorus. 2. By the
      action of sulphuric acid on peroxide of barium. If the latter
      be covered with strong sulphuric acid (the acid, if diluted
      with only one-tenth of water, does not give ozone), then at a
      low temperature the oxygen evolved contains ozone, and in much
      greater quantities than in that ozone is obtained by the action
      of electric sparks or phosphorus. 3. Ozone may also be obtained
      by decomposing strong sulphuric acid by potassium manganate
      especially with the addition of barium peroxide.

  [7] Ozone takes up the hydrogen from hydrochloric acid; chlorine is
      liberated, and can dissolve gold. Iodine is directly oxidised
      by ozone, but not by oxygen. Ammonia, NH_{3}, is oxidised by
      ozone into ammonium nitrite (and nitrate), 2NH_{3} + O_{3} =
      NH_{4}NO_{2} + H_{2}O, and therefore a drop of ammonia, on
      falling into the gas, gives a thick cloud of the salts formed.
      Ozone converts lead oxide into peroxide, and suboxide of
      thallium (which is colourless) into oxide (which is brown), so
      that this reaction is made use of for detecting the presence of
      ozone. Lead sulphide, PbS (black), is converted into sulphate,
      PbSO_{4} (colourless), by ozone. A neutral solution of manganese
      sulphate gives a precipitate of manganese peroxide, and an acid
      solution may be oxidised into permanganic acid, HMnO_{4}. With
      respect to the oxidising action of ozone on organic substances,
      it may be mentioned that with ether, C_{4}H_{10}O, ozone gives
      ethyl peroxide, which is capable of decomposing with explosion
      (according to Berthelot), and is decomposed by water into alcohol,
      2C_{2}H_{6}O, and hydrogen peroxide, H_{2}O_{2}.

  [8] This reaction is the one usually made use of for detecting the
      presence of ozone. In the majority of cases paper is soaked in
      solutions of potassium iodide and starch. Such _ozonometrical_
      or iodised starch-paper when damp turns blue in the presence
      of ozone, and the tint obtained varies considerably, according
      to the length of time it is exposed and to the amount of ozone
      present. The amount of ozone in a given gas may even to a certain
      degree he judged by the shade of colour acquired by the paper, if
      preliminary tests be made.

      Test-paper for ozone is prepared in the following manner:--One
      gram of neutral potassium iodide is dissolved in 100 grams of
      distilled water; 10 grams of starch are then shaken up in the
      solution, and the mixture is boiled until the starch is converted
      into a jelly. This jelly is then smeared over blotting-paper and
      left to dry. It must always he remembered, however, that the
      colour of iodised starch-paper is changed not only by the action
      of ozone, but of many other oxidisers; for example, by the oxides
      of nitrogen (especially N_{2}O_{3}) and hydrogen peroxide. Houzeau
      proposed soaking common litmus-paper with a solution of potassium
      iodide, which in the presence of iodine would turn blue, owing to
      the formation of KHO. In order to determine if the blue colour
      is not produced by an alkali (ammonia) in the gas, a portion of
      the paper is not soaked in the potassium iodide, but moistened
      with water; this portion will then also turn blue if ammonia be
      present. A reagent for distinguishing ozone from hydrogen peroxide
      with certainty is not known, and therefore these substances in
      very small quantities (for instance, in the atmosphere) may easily
      he confounded. Until recent years the mistake has frequently been
      made of ascribing the alteration of iodised starch-paper in the
      air to the presence of ozone; at the present time there is reason
      to believe that it is most often due to the presence of nitrous
      acid (Ilosva, 1889).

Hence _ozone_, although it has the same _composition as oxygen_, differs
from it in stability, and by the fact that it oxidises a number of
substances very energetically at the ordinary temperature. In this
respect ozone resembles the oxygen of certain unstable compounds, or
oxygen at the moment of its liberation.[8 bis]

  [8 bis] Fluorine (Chap. XI.), acting upon water at the ordinary
      temperature, takes up the hydrogen, and evolves the oxygen in the
      form of ozone (Moissan, 1889), and therefore the reaction must be
      expressed thus:--3H_{2}O + 3F_{2} = 6HF + O_{3}.

In ordinary oxygen and ozone we see an example of one and the same
substance, in this case an element, appearing in two states. This
indicates that the properties of a substance, and even of an element,
may vary without its composition varying. Very many such cases are
known. Such cases of a chemical transformation which determine a
difference in the properties of one and the same element are termed
cases of isomerism. The cause of isomerism evidently lies deep within
the essential conditions of a substance, and its investigation has
already led to a number of results of unexpected importance and of
immense scientific significance. It is easy to understand the difference
between substances containing different elements or the same elements
in different proportions. That a difference should exist in these cases
necessarily follows, if, as our knowledge compels us, we admit that
there is a radical difference in the simple bodies or elements. But
when the composition--_i.e._ the quality and quantity of the elements
in two substances is the same and yet their properties are different,
then it becomes clear that the conceptions of diverse elements and of
the varying composition of compounds, alone, are insufficient for the
expression of all the diversity of properties of matter in nature.
Something else, still more profound and internal than the composition of
substances, must, judging from isomerism, determine the properties and
transformation of substances.

On what are the isomerism of ozone and oxygen, and the peculiarities
of ozone, dependent? In what, besides the extra store of energy,
which is one of the peculiarities of ozone, resides the cause of its
difference from oxygen? These questions for long occupied the minds
of investigators, and were the motive for the most varied, exact, and
accurate researches, which were chiefly directed to the study of the
volumetric relations exhibited by ozone. In order to acquaint the reader
with the previous researches of this kind, I cite the following from a
memoir by Soret, in the 'Transactions of the French Academy of Sciences'
for 1866:

'Our present knowledge of the volumetric relations of ozone may be
expressed in the following manner:

'1. "Ordinary oxygen in changing into ozone under the action of
electricity shows a diminution in volume." This was discovered by
Andrews and Tait.

'2. "In acting on ozonised oxygen with potassium iodide and other
substances capable of being oxidised, we destroy the ozone, but the
volume of the gas remains unchanged." For the researches of Andrews,
Soret, v. Babo, and others showed that the proportion of ozonised oxygen
absorbed by the potassium iodide is equal to the original contraction of
volume of the oxygen--that is, in the absorption of the ozone the volume
of the gas remains unchanged. From this it might be imagined that ozone,
so to say, does not occupy any space--is indefinitely dense.

'3. "By the action of heat ozonised oxygen increases in volume, and is
transformed into ordinary oxygen. This increase in volume corresponds
with the quantity of ozonised oxygen which is given up to the potassium
iodide in its decomposition" (the same observers).

'4. These unquestionable experimental results lead to the conclusion
that ozone is denser than oxygen, and that in its oxidising action
it gives off that portion of its substance to which is due its extra
density distinguishing it from ordinary oxygen.'

If we imagine (says Weltzien) that _n_ volumes of ozone consist of _n_
volumes of oxygen combined with _m_ volumes of the same substance,
and that ozone in oxidising gives up _m_ volumes of oxygen and leaves
_n_ volumes of ordinary oxygen gas, then all the above facts can be
explained; otherwise it must be supposed that ozone is infinitely dense.
'In order to determine the density of ozone' (we again cite Soret)
'recourse cannot be had to the direct determination of the weight of
a given volume of the gas, because ozone cannot be obtained in a pure
state. It is always mixed with a very large quantity of oxygen. It was
necessary, therefore, to have recourse to such substances as would
absorb ozone without absorbing oxygen and without destroying the ozone.
Then the density might be deduced from the decrease of volume produced
in the gas by the action of this solvent in comparison with the quantity
of oxygen given up to potassium iodide. Advantage must also be taken of
the determination of the increase of volume produced by the action of
heat on ozone, if the volume occupied by the ozone before heating be
known.' Soret found two such substances, turpentine and oil of cinnamon.
'Ozone disappears in the presence of turpentine. This is accompanied
by the appearance of a dense vapour, which fills a vessel of small
capacity (0·14 litre) to such an extent that it is impenetrable to
direct solar-rays. On leaving the vessel at rest, it is observed that
the cloud of vapour settles; the clearing is first remarked at the upper
portion of the vessel, and the brilliant colours of the rainbow are
seen on the edge of a cloud of vapour.' Oil of cinnamon--that is, the
volatile or essential oil of the well-known spice, cinnamon--gives under
similar circumstances the same kind of vapours, but they are much less
voluminous. On measuring the gaseous volume before and after the action
of both volatile oils, a considerable decrease is remarked. On applying
all the necessary corrections (for the solubility of oxygen in the oily
liquids named above, for the tension of their vapour, for the change of
pressure, &c.) and making a series of comparative determinations, Soret
obtained the following result: two volumes of ozone capable of being
dissolved, when changed to ordinary (by heating a wire to a red-heat by
a galvanic current) increase by one volume. Hence it is evident that
in the formation of ozone three volumes of oxygen give two volumes of
ozone--that is, its density (referred to hydrogen) = 24.

The observations and determinations of Soret showed that ozone is
heavier than oxygen, and even than carbonic anhydride (because
ozonised oxygen passes through fine orifices more slowly than oxygen
and than its mixtures with carbonic anhydride), although lighter than
chlorine (it flows more rapidly through such orifices than chlorine),
and they indicated that _ozone is one and a half times denser than
oxygen_, which may be expressed by designating a molecule of oxygen
by O_{2} and of ozone by O_{3}, and hence ozone OO_{2} is comparable
with compound substances[9] formed by oxygen, as for instance CO_{2},
SO_{2}, NO_{2}, &c. This explains the chief differences between ozone
and oxygen and the cause of the isomerism, and at the same time leads
one to expect[10] that ozone, being a gas which is denser than oxygen,
would be liquefied much more easily. This was actually shown to be the
case in 1880, by Chappuis and Hautefeuille in their researches on the
_physical properties of ozone_. Its boiling point under a pressure of
760 mm. is about -106°, and consequently compressed and refrigerated
ozone when rapidly expanded forms drops, _i.e._ is liquefied. Liquid
and compressed[11] ozone is blue. In dissolving in water ozone partly
passes into oxygen. It explodes violently when suddenly compressed and
heated, changing into ordinary oxygen and evolving, like all explosive
substances,[12] that extra heat which distinguishes it from oxygen.

  [9] Ozone is, so to say, an oxide of oxygen, just as water is an oxide
      of hydrogen. Just as aqueous vapour is composed of two volumes of
      hydrogen and one volume of oxygen, which on combining condense
      into two volumes of aqueous vapour, so also two volumes of oxygen
      are combined in ozone with one volume of oxygen to give two
      volumes of ozone. In the action of ozone on different substances
      it is only that additional portion of its molecule by which it
      differs from ordinary oxygen that combines with other bodies, and
      that is why, under these circumstances, the volume of the ozonised
      oxygen does not change. Starting with two volumes of ozone,
      one-third of its weight is parted with, and two volumes of oxygen
      remain.

      The above observations of Soret on the capacity of turpentine
      for dissolving ozone, together with Schönbein's researches on
      the formation of ozone in the oxidation of turpentine and of
      similar volatile vegetable oils (entering into the composition
      of _perfumes_), also explain the action of this ethereal oil on
      a great many substances. It is known that turpentine oil, when
      mixed with many substances, promotes their oxidation. In this case
      it probably not only itself promotes the formation of ozone, but
      also dissolves ozone from the atmosphere, and thus acquires the
      property of oxidising many substances. It bleaches linen and cork,
      decolorises indigo, promotes the oxidation and hardening of boiled
      linseed oil, &c. These properties of turpentine oil are made use
      of in practice. Dirty linen and many stained materials are easily
      cleaned by turpentine, not only because it dissolves the grease,
      but also because it oxidises it. The admixture of turpentine with
      drying (boiled) oil, oil-colours, and lacs aids their rapid drying
      because it attracts ozone. Various oils occurring in plants,
      and entering into the composition of perfumes and certain scent
      extracts, also act as oxidisers. They act in the same manner as
      oil of turpentine and oil of cinnamon. This perhaps explains
      the refreshing influence they have in scents and other similar
      preparations, and also the salubrity of the air of pine forests.
      Water upon which a layer of turpentine oil has been poured
      acquires, when left standing in the light, the disinfecting and
      oxidising properties in general of ozonised turpentine (is this
      due to the formation of H_{2}O_{2}?).

  [10] The densest, most complex, and heaviest particles of matter
       should, under equal conditions, evidently be less capable of
       passing into a state of gaseous motion, should sooner attain a
       liquid state, and have a greater cohesive force.

  [11] The blue colour proper to ozone may be seen through a tube one
       metre long, filled with oxygen, containing 10 p.c. of ozone.
       The density of liquid ozone has not, so far as I am aware, been
       determined.

  [12] All explosive bodies and mixtures (gunpowder, detonating gas,
       &c.) evolve heat in exploding--that is, the reactions which
       accompany explosions are exothermal. In this manner ozone in
       decomposing evolves latent heat, although generally heat is
       absorbed in decomposition. This shows the meaning and cause of
       explosion.

Thus, judging by what has been said above, ozone should he formed in
nature not only in the many processes of oxidation which go on, but also
by the condensation of atmospheric oxygen. The significance of ozone
in nature has often arrested the attention of observers. There is a
series of ozonometrical observations which show the different amounts
of ozone in the air at different localities, at different times of the
year, and under different circumstances. But the observations made in
this direction cannot be considered as sufficiently exact, because the
methods in use for determining ozone were not quite accurate. It is
however indisputable[13] that the amount of ozone in the atmosphere is
subject to variation; that the air of dwellings contains no ozone (it
disappears in oxidising organic matter); that the air of fields and
forests always contains ozone, or substances (peroxide of hydrogen)
which act like it (on iodised starch paper &c.)[13 bis]; that the amount
of ozone increases after storms; and that miasms, &c., are destroyed by
ozonising the atmosphere. It easily oxidises organic substances, and
miasms are produced by organic substances and the germs of organisms,
all of which are easily changed and oxidised. Indeed, many miasms--for
instance, the volatile substance of decomposing organisms--are clearly
destroyed or changed not only by ozone, but also by many other
powerfully oxidising substances, such as chlorine water, potassium
permanganate, and the like.[14] All that is now known respecting
the presence of ozone in the air may be summed up in the following
words: A small quantity of an oxidising substance, resembling ozone
in its reactions, has undoubtedly been observed and determined in the
atmosphere, especially in fresh air, for instance after a storm, and it
is very likely that this substance contains a mixture of such oxidising
substances as ozone, peroxide of hydrogen, and the lower oxides of
nitrogen (especially nitrous acid and its ammonia salt) produced from
the elements of the atmosphere by oxidation and by the action of
electrical discharges.

  [13] In Paris it has been found that the further from the centre of the
       town the greater the amount of ozone in the air. The reason of
       this is evident: in a city there are many conditions for the
       destruction of ozone. This is why we distinguish country air
       as being fresh. In spring the air contains more ozone than in
       autumn; the air of fields more than the air of towns.

  [13 bis] The question of the presence of ozone in the air has not yet
       been fully elucidated, as those reactions by which ozone is
       generally detected are also common to nitrous acid (and its
       ammonia salt). Ilosvay de Ilosva (1889), in order to exclude
       the influence of such bodies, passed air through a 40 per cent.
       solution of caustic soda, and then through a 20 per cent.
       solution of sulphuric acid (these solutions do not destroy
       ozone), and tested the air thus purified for the presence of
       ozone. As no ozone was then detected the author concludes that
       all the effects which were formerly ascribed to ozone should
       be referred to nitrous acid. But this conclusion requires more
       careful verification, since the researches of Prof. Schönbein on
       the presence of peroxide of hydrogen in the atmosphere.

  [14] The oxidising action of ozone may be taken advantage of for
       technical purposes; for instance, for destroying colouring
       matters. It has even been employed for bleaching tissues and for
       the rapid preparation of vinegar, although these methods have not
       yet received wide application.

Thus in ozone we see (1) the capacity of elements (and it must be
all the more marked in compounds) of changing in properties without
altering in composition; this is termed isomerism;[15] (2) the capacity
of certain elements for condensing themselves into molecules of
different densities; this forms a special case of isomerism called
_polymerism_; (3) the capacity of oxygen for appearing in a still more
active and energetic chemical state than that in which it occurs in
ordinary gaseous oxygen; and (4) the formation of unstable equilibria,
or chemical states, which are illustrated both by the ease with which
ozone acts as an oxidiser and by its capacity for decomposing with
explosion.[16]

  [15] Isomerism in elements is termed _allotropism_.

  [16] A number of substances resemble ozone in one or other of these
       respects. Thus cyanogen, C_{2}N_{2}, nitrogen chloride, &c.,
       decompose with an explosion and evolution of heat. Nitrous
       anhydride, N_{2}O_{3}, forms a blue liquid like ozone, and in a
       number of cases oxidises like ozone.

_Hydrogen peroxide._--Many of those properties which we have seen in
ozone belong also to a peculiar substance containing oxygen and hydrogen
and called hydrogen peroxide or oxygenated water. This substance was
discovered in 1818 by Thénard. When heated it is decomposed into
water and oxygen, evolving as much oxygen as is contained in the
water remaining after the decomposition. That portion of oxygen by
which hydrogen peroxide differs from water behaves in a number of
cases just like the active oxygen in ozone, which distinguishes it
from ordinary oxygen. In H_{2}O_{2}, and in O_{3}, one atom of oxygen
acts as a powerful oxidiser, and on separating out it leaves H_{2}O
or O_{2}, which do not act so energetically, although they still
contain oxygen.[17] Both H_{2}O_{2} and O_{3} contain the oxygen in
a compressed state, so to speak, and when freed from pressure by the
forces (internal) of the elements in another substance, this oxygen is
easily evolved, and therefore acts as oxygen does at the moment of its
liberation. Both substances in decomposing, with the separation of a
portion of their oxygen, _evolve_ heat, whilst decomposition is usually
accompanied by an absorption of heat.

  [17] It is evident that there is a want of words here for distinguishing
       oxygen, O, as an ultimate _element_, from oxygen, O_{2}, as a
       _free element_. The latter should be termed oxygen gas, did not
       custom and the length of the expression render it inconvenient.

Hydrogen peroxide is formed under many circumstances by combustion and
oxidation, but in very limited quantities; thus, for instance, it is
sufficient to shake up zinc with sulphuric acid, or even with water,
to observe the formation of a certain quantity of hydrogen peroxide in
the water.[18] From this cause, probably, a series of diverse oxidation
processes are accomplished in nature, and according to Prof. Schöne of
Moscow, hydrogen peroxide occurs in the atmosphere, although in variable
and small quantities, and probably its formation is connected with
ozone, with which it has much in common. The usual mode of the formation
of hydrogen peroxide, and the method by which it may be indirectly
obtained,[19] is by the double decomposition of an acid and the
peroxides of certain metals, especially those of potassium, calcium,
and barium.[20] We saw when speaking of Oxygen (Chap. III.) that it is
only necessary to heat the anhydrous oxide of barium to a red heat in a
current of air or oxygen (or, better still, to heat it with potassium
chlorate, and then to wash away the potassium chloride formed) to obtain
peroxide of barium.[21] Barium peroxide gives hydrogen peroxide by the
action of acids in the cold.[22] The process of decomposition is very
clear in this case; the hydrogen of the acid replaces the barium of the
peroxide, a barium salt of the acid being formed, while the hydrogen
peroxide formed in the reaction remains in solution.[23]

  [18] Schönbein states that the formation of hydrogen peroxide is
       to be remarked in every oxidation in water or in the presence
       of aqueous vapour. According to Struve, hydrogen peroxide
       is contained in snow and in rain-water, and its formation,
       together with ozone and ammonium nitrate, is even probable in
       the processes of respiration and combustion. A solution of tin
       in mercury, or liquid tin amalgam, when shaken up in water
       containing sulphuric acid, produces hydrogen peroxide, whilst
       iron under the same circumstances does not give rise to its
       formation. The presence of small quantities of hydrogen peroxide
       in these and similar cases is recognised by many reactions.
       Amongst them, its action on _chromic acid_ in the presence of
       ether is very characteristic. Hydrogen peroxide converts the
       chromic acid into a higher oxide, Cr_{2}O_{7}, which is of a
       dark-blue colour and dissolves in ether. This ethereal solution
       is to a certain degree stable, and therefore the presence of
       hydrogen peroxide may be recognised by mixing the liquid to be
       tested with ether and adding several drops of a solution of
       chromic acid. On shaking the mixture the ether dissolves the
       higher oxide of chromium which is formed, and acquires a blue
       colour. The formation of hydrogen peroxide in the combustion and
       oxidation of substances containing or evolving hydrogen must
       be understood in the light of the conception, to be considered
       later, of molecules occupying equal volumes in a gaseous state.
       At the moment of its evolution a molecule H_{2} combines with
       a molecule O_{2}, and gives H_{2}O_{2}. As this substance is
       unstable, a large proportion of it is decomposed, a small amount
       only remaining unchanged. If it is obtained, water is easily
       formed from it; this reaction evolves heat, and the reverse
       action is not very probable. Direct determinations show that
       the reaction H_{2}O_{2} = H_{2}O + O evolves 22,000 heat units.
       From this it will be understood how easy is the decomposition
       of hydrogen peroxide, as well as the fact that a number of
       substances which are not directly oxidised by oxygen are oxidised
       by hydrogen peroxide and by ozone, which also evolves heat on
       decomposition. Such a representation of the origin of hydrogen
       peroxide has been developed by me since 1870. Recently (1890)
       Traube has pronounced a similar opinion, stating that Zn under
       the action of water and air gives, besides ZnH_{2}O_{2}, also
       H_{2}O_{2}.

  [19] The formation of hydrogen peroxide from barium peroxide by a
       method of double decomposition is an instance of a number of
       _indirect methods of preparation_. A substance A does not combine
       with B, but A B is obtained from A C in its action on B D (see
       Introduction) when C D is formed. Water does not combine with
       oxygen, but as a hydrate of acids it acts on the compound of
       oxygen with barium oxide, because this oxide gives a salt with an
       acid anhydride; or, what is the same thing, hydrogen with oxygen
       does not directly form hydrogen peroxide, but when combined with
       a haloid (for example, chlorine), under the action of barium
       peroxide, BaO_{2}, it leads to the formation of a salt of barium
       and H_{2}O_{2}. It is to be remarked that the passage of barium
       oxide, BaO, into the peroxide, BaO_{2}, is accompanied by the
       _evolution_ of 12,100 heat units per 16 parts of oxygen by weight
       combined, and the passage of H_{2}O into the peroxide H_{2}O_{2}
       does not proceed directly, because it would be accompanied by the
       _absorption_ of 22,000 units of heat by 16 parts by weight of
       oxygen combined. Barium peroxide, in acting on an acid, evidently
       evolves less heat than the oxide, and it is this difference of
       heat that is absorbed in the hydrogen peroxide. Its energy is
       obtained from that evolved in the formation of the salt of barium.

  [20] Peroxides of lead and manganese, and other analogous peroxides
       (see Chap. III., Note 9), do not give hydrogen peroxide under
       these conditions, but yield chlorine with hydrochloric acid.

  [21] The impure barium peroxide obtained in this manner may be easily
       purified. For this purpose it is dissolved in a dilute solution
       of nitric acid. A certain quantity of an insoluble residue always
       remains, from which the solution is separated by filtration.
       The solution will contain not only the compound of the barium
       peroxide, but also a compound of the barium oxide itself,
       a certain quantity of which always remains uncombined with
       oxygen. The acid compounds of the peroxide and oxide of barium
       are easily distinguishable by their stability. The peroxide
       gives an unstable compound, and the oxide a stable salt. By
       adding an aqueous solution of barium oxide to the resultant
       solution, the whole of the peroxide contained in the solution
       may be precipitated as a pure aqueous compound (Kouriloff, 1889,
       obtained the same result by adding an excess of BaO_{2}). The
       first portions of the precipitate will consist of impurities--for
       instance, oxide of iron. The barium peroxide then separates out,
       and is collected on a filter and washed; it forms a substance
       having a definite composition, BaO_{2},8H_{2}O, and is very
       pure. Pure hydrogen peroxide should always be prepared from such
       purified barium peroxide.

  [22] In the cold, strong sulphuric acid with barium peroxide gives
       ozone; when diluted with a certain amount of water it gives
       oxygen (see Note 6), and hydrogen peroxide is only obtained
       by the action of very weak sulphuric acid. Hydrochloric,
       hydrofluoric, carbonic, and hydrosilicofluoric acids, and others,
       when diluted with water also give hydrogen peroxide with barium
       peroxide. Professor Schöne, who very carefully investigated
       hydrogen peroxide, showed that it is formed by the action of many
       of the above-mentioned acids on barium peroxide. In preparing
       peroxide of hydrogen by means of sulphuric acid, the solution
       must be kept cold. A solution of maximum concentration may
       be obtained by successive treatments with sulphuric acid of
       increasing strength. In this manner a solution containing 2 to 3
       grams of pure peroxide in 100 c.c. of water may be obtained (V.
       Kouriloff).

  [23] With the majority of acids, that salt of barium which is formed
       remains in solution; thus, for instance, by employing
       hydrochloric acid, hydrogen peroxide and barium chloride remain
       in solution. Complicated processes would be required to obtain
       pure hydrogen peroxide from such a solution. It is much more
       convenient to take advantage of the action of carbonic anhydride
       on the pure hydrate of barium peroxide. For this purpose the
       hydrate is stirred up in water, and a rapid stream of carbonic
       anhydride is passed through the water. Barium carbonate,
       insoluble in water, is formed, and the hydrogen peroxide remains
       in solution, so that it may be separated from the carbonate
       by filtering only. On a large scale hydrofluosilicic acid is
       employed, its barium salt being also insoluble in water.

The reaction is expressed by the equation BaO_{2} + H_{2}SO_{4} =
H_{2}O_{2} + BaSO_{4}. It is best to take a weak cold solution of
sulphuric acid and to almost saturate it with barium peroxide, so that
a small excess of acid remains; insoluble barium sulphate is formed. A
more or less dilute aqueous solution of hydrogen peroxide is obtained.
This solution may be concentrated in a vacuum over sulphuric acid. In
this way the water may even be entirely evaporated from the solution of
the hydrogen peroxide; only in this case it is necessary to work at a
low temperature, and not to keep the peroxide for long in the rarefied
atmosphere, as otherwise it decomposes.[23 bis] A solution of peroxide
of hydrogen (mixed with the solution of a salt of sodium NaX) is used
for bleaching (especially silk and wool) on a large scale, and is now
usually prepared from peroxide of sodium Na_{2}O_{2} by the action of
acids. Na_{2}O_{2} + 2HX = 2NaX + H_{2}O_{2}[24].

  [23 bis] Hydrogen peroxide may be extracted from very dilute solutions
       by means of ether, which dissolves it, and when mixed with it
       the hydrogen peroxide may even be distilled. A solution of
       hydrogen peroxide in water may be strengthened by cooling it to
       a low temperature, when the water crystallises out--that is,
       is converted into ice--whilst the hydrogen peroxide remains in
       solution, as it only freezes at very low temperatures. It must be
       observed that hydrogen peroxide, in a strong solution in a pure
       state, is exceedingly unstable even at the ordinary temperature,
       and therefore it must be preserved in vessels always kept cold,
       as otherwise it evolves oxygen and forms water.

  [24] Peroxide of sodium (Chap. XII., Note 49) is prepared by burning
       sodium in dry air.

When pure, hydrogen peroxide is a colourless liquid, without smell, and
having a very unpleasant taste--such as belongs to the salts of many
metals--the so-called 'metallic' taste. Water stored in zinc vessels has
this taste, which is probably due to its containing hydrogen peroxide.
The tension of the vapour of hydrogen peroxide is less than that of
aqueous vapour; this enables its solutions to be concentrated in a
vacuum. The specific gravity of anhydrous hydrogen peroxide is 1·455.
Hydrogen peroxide decomposes, with the evolution of oxygen, when heated
even to 20°. But the more dilute its aqueous solution the more stable
it is. Very weak solutions may be distilled without decomposing the
hydrogen peroxide. It decolorises solutions of litmus and turmeric, and
acts in a similar manner on many colouring matters of organic origin
(for which reason it is employed for bleaching tissues).[24 bis]

  [24 bis] Peroxide of hydrogen should apparently find an industrial
       application in the arts, for instance, (1) as a bleaching agent,
       it having the important advantage over chloride of lime, SO_{2},
       &c., of not acting upon the material under treatment. It may be
       used for bleaching feathers, hair, silk, wool, wood, &c., it also
       removes stains of all kinds, such as wine, ink, and fruit stains;
       (2) it destroys bacteria like ozone without having any injurious
       effect upon the human body. It can also be used for washing all
       kinds of wounds, for purifying the air in the sick room, &c., and
       (3) as a preserving agent for potted meats, &c.

_Many substances decompose hydrogen peroxide_, forming water and oxygen,
without apparently suffering any change. In this case substances in a
state of fine division show a much quicker action than compact masses,
from which it is evident that the action is here based on contact
(_see_ Introduction). It is sufficient to bring hydrogen peroxide
into contact with charcoal, gold, the peroxide of manganese or lead,
the alkalis, metallic silver, and platinum, to bring about the above
decomposition.[25] Besides which, hydrogen peroxide forms water and
parts with its oxygen with great ease to a number of substances which
are capable of being oxidised or of combining with oxygen, and in this
respect is very like ozone and other _powerful oxidisers_.[26] To the
class of contact phenomena, which are so characteristic of hydrogen
peroxide as a substance which is unstable and easily decomposable with
the evolution of heat, must be referred the following--that in the
presence of many substances containing oxygen it evolves, not only its
own oxygen, but also that of the substances which are brought into
contact with it--that is, _it acts in a reducing manner_. It behaves
thus with ozone, the oxides of silver, mercury, gold and platinum,
and lead peroxide. The oxygen in these substances is not stable, and
therefore the feeble influence of contact is enough to destroy its
position. Hydrogen peroxide, especially in a concentrated form, in
contact with these substances, evolves an immense quantity of oxygen,
so that an explosion takes place and an exceedingly powerful evolution
of heat is observed if hydrogen peroxide in a concentrated form be made
to drop upon these substances in dry powder. Slow decomposition also
proceeds in dilute solutions.[27]

  [25] As the result of careful research, certain of the _catalytic_ or
       contact phenomena have been subjected to exact explanation, which
       shows the participation of a substance present in the process
       of a reaction, whilst, however, it does not alter the series
       of changes proceeding from mechanical actions only. Professor
       Schöne, of the Petroffsky Academy, has already explained a
       number of reactions of hydrogen peroxide which previously were
       not understood. Thus, for instance, he showed that with hydrogen
       peroxide, alkalis give peroxides of the alkaline metals, which
       combine with the remaining hydrogen peroxide, forming unstable
       compounds which are easily decomposed, and therefore alkalis
       evince a decomposing (catalytic) influence on solutions of
       hydrogen peroxide. Only acid solutions of hydrogen peroxide, and
       then only dilute ones, can be preserved well.

  [26] _Hydrogen peroxide_, as a substance containing much oxygen
       (namely, 16 parts to one part by weight of hydrogen), exhibits
       many _oxidising reactions_. Thus, it oxidises arsenic, converts
       lime into calcium peroxide, the oxides of zinc and copper into
       peroxides; it parts with its oxygen to many sulphides, converting
       them into sulphates, &c. So, for example, it converts black
       lead sulphide, PbS, into white lead sulphate, PbSO_{4}, copper
       sulphide into copper sulphate, and so on. The restoration of old
       oil paintings by hydrogen peroxide is based on this action. Oil
       colours are usually admixed with white lead, and in many cases
       the colour of oil-paints becomes darker in process of time. This
       is partly due to the sulphuretted hydrogen contained in the
       air, which acts on white lead, forming lead sulphide, which is
       black. The intermixture of the black colour darkens the rest.
       In cleaning a picture with a solution of hydrogen peroxide,
       the black lead sulphide is converted into white sulphate, and
       the colours brighten owing to the disappearance of the black
       substance which previously darkened them. Hydrogen peroxide
       oxidises with particular energy substances containing hydrogen
       and capable of easily parting with it to oxidising substances.
       Thus it decomposes hydriodic acid, setting the iodine free
       and converting the hydrogen it contains into water; it also
       decomposes sulphuretted hydrogen in exactly the same manner,
       setting the sulphur free. Starch paste with potassium iodide
       is not, however, directly  by peroxide of hydrogen in
       the entire absence of free acids; but the addition of a small
       quantity of iron sulphate (green vitriol) or of lead acetate to
       the mixture is enough to entirely blacken the paste. This is a
       very sensitive reagent (test) for peroxide of hydrogen, like the
       test with chromic acid and ether (_see_ Note 8).

  [27] To explain the phenomenon, an hypothesis has been put forward
       by Brodie, Clausius, and Schönbein which supposes ordinary oxygen
       to be an electrically neutral substance, composed, so to speak,
       of two electrically opposite kinds of oxygen--positive and
       negative. It is supposed that hydrogen peroxide contains one kind
       of such polar oxygen, whilst in the oxides of the above-named
       metals the oxygen is of opposite polarity. It is supposed that
       in the oxides of the metals the oxygen is electro-negative, and
       in hydrogen peroxide electro-positive, and that on the mutual
       contact of these substances ordinary neutral oxygen is evolved
       as a consequence of the mutual attraction of the oxygens of
       opposite polarity. Brodie admits the polarity of oxygen in
       combination, but not in an uncombined state, whilst Schönbein
       supposes uncombined oxygen to be polar also, considering ozone as
       electro-negative oxygen. The supposition that the oxygen of ozone
       is different from that of hydrogen peroxide is contradicted by
       the fact that in acting on barium peroxide strong sulphuric acid
       forms ozone, and dilute acid forms hydrogen peroxide.

Just as a whole series of metallic compounds, and especially the
oxides and their hydrates, correspond with water, so also there are
many substances analogous to hydrogen peroxide. Thus, for instance,
calcium peroxide is related to hydrogen peroxide in exactly the same
way as calcium oxide or lime is related to water. In both cases the
hydrogen is replaced by a metal--namely, by calcium.[27 bis] But it is
most important to remark that the nearest approach to the properties
of hydrogen peroxide is afforded by a non-metallic element, chlorine;
its action on colouring matters, its capacity for oxidising, and for
evolving oxygen from many oxides, is analogous to that exhibited by
hydrogen peroxide. Even the very formation of chlorine is closely
analogous to the formation of peroxide of hydrogen; chlorine is obtained
from manganese peroxide, MnO_{2}, and hydrochloric acid, HCl, and
hydrogen peroxide from barium peroxide, BaO_{2}, and the same acid.
The result in one case is essentially water, chlorine, and manganese
chloride; and in the other case barium chloride and hydrogen peroxide
are produced. Hence water + chlorine corresponds with hydrogen peroxide,
and the action of chlorine in the presence of water is analogous to the
action of hydrogen peroxide. This analogy between chlorine and hydrogen
peroxide is expressed in the conception of an aqueous radicle, which
(Chapter III.) has been already mentioned. _This aqueous radicle_ (or
hydroxyl) is that which is left from water if it be imagined as deprived
of half of its hydrogen. According to this method of expression, caustic
soda will be a compound of sodium with the aqueous radicle, because it
is formed from water with the evolution of half the hydrogen. This is
expressed by the following formulæ: water, H_{2}O, caustic soda, NaHO,
just as hydrochloric acid is HCl and sodium chloride NaCl. Hence the
aqueous radicle HO is a compound radicle, just as chlorine, Cl, is a
simple radicle. They both give hydrogen compounds, HHO, water, and HCl,
hydrochloric acid; sodium compounds, NaHO and NaCl, and a whole series
of analogous compounds. Free chlorine in this sense will be ClCl, and
hydrogen peroxide HOHO, which indeed expresses its composition, because
it contains twice as much oxygen as water does.[28]

  [27 bis] It should be mentioned that Schiloff (1893) on taking a 3 per
       cent. solution of H_{2}O_{2}, adding soda to it, and then
       extracting the peroxide of hydrogen from the mixture by shaking
       it with ether, obtained a 50 per cent. solution of H_{2}O_{2},
       which, although perfectly free from other acids, gave a
       distinctly acid reaction with litmus. And here attention should
       first of all be turned to the fact that the peroxides of the
       metals correspond to H_{2}O_{2}, like salts to an acid, for
       instance, Na_{2}O_{2} and BaO_{2}, &c. Furthermore, it must be
       remembered that O is an analogue of S (Chapters XV. and XX.),
       and sulphur gives H_{2}S, H_{2}SO_{3}, and H_{2}SO_{4}. And
       sulphurous acid, H_{2}SO_{3}, is unstable as a hydrate, and
       gives water and the anhydride SO_{2}. If the sulphur be replaced
       by oxygen, then instead of H_{2}SO_{3} and SO_{2}, we have
       H_{2}OO_{3} and OO_{2}. The latter is ozone, while the salt
       K_{2}O_{4} (peroxide of potassium) corresponds to the hydrate
       H_{2}O_{4} as to an acid. And between H_{2}O and H_{2}O_{4}
       there may exist intermediate acid compounds, the first of which
       would be H_{2}O_{2}, in which, from analogy to the sulphur
       compounds, one would expect acid properties. Besides which we
       may mention that for sulphur, besides H_{2}S (which is a feeble
       acid), H_{2}S_{2}, H_{2}S_{3}, H_{2}S_{5} are known. Thus in
       many respects H_{2}O_{2} offers points of resemblance to acid
       compounds, and as regards its qualitative (reactive) analogies,
       it not only resembles Na_{2}O_{2}, BaO_{2}, &c., but also
       persulphuric acid HSO_{4} (to which the anhydride S_{2}O_{7}
       corresponds) and Cu_{2}O_{7}, &c., which will be subsequently
       described.

  [28] Tamman and Carrara (1892) showed by determining the depression
       (fall of the temperature of the formation of ice, Chapters I.
       and VII.) that the molecule of peroxide of hydrogen contains
       H_{2}O_{2}, and not HO or H_{3}O_{3}.

Thus in ozone and hydrogen peroxide we see examples of very unstable,
easily decomposable (by time, spontaneously, and on contact) substances,
full of the energy necessary for change,[28 bis] capable of being
easily reconstituted (in this case decomposing with the evolution of
heat); they are therefore examples of _unstable chemical equilibria_.
If a substance exists, it signifies that it already presents a certain
form of equilibrium between those elements of which it is built up.
But chemical, like mechanical, equilibria exhibit different degrees of
stability or solidity.[29]

  [28 bis] The lower oxides of nitrogen and chlorine and the higher oxides
       of manganese are also formed with the absorption of heat, and
       therefore, like hydrogen peroxide, act in a powerfully oxidising
       manner, and are not formed by the same methods as the majority
       of other oxides. It is evident that, being endowed with a richer
       store of energy (acquired in combination or by absorption of
       heat), such substances, compared with others poorer in energy,
       will exhibit a greater diversity of cases of chemical action with
       other substances.

  [29] If the point of support of a body lies in a vertical line below
       the centre of gravity, it is in unstable equilibrium. If the
       centre of gravity lies below the point of support; the state of
       equilibrium is very stable, and a vibration may take place about
       this position of stable equilibrium, as in a pendulum or balance,
       when finally the body assumes a position of stable equilibrium.
       But if, keeping to the same mechanical example, the body be
       supported not on a point, in the geometrical sense of the word,
       but on a small plane, then the state of unstable equilibrium may
       be preserved, unless destroyed by external influences. Thus a
       man stands upright supported on the plane, or several points of
       the surfaces of his feet, having the centre of gravity above the
       points of support. Vibration is then possible, but it is limited,
       otherwise on passing outside the limit of possible equilibrium
       another more stable position is attained about which vibration
       becomes more possible. A prism immersed in water may have several
       more or less stable positions of equilibrium. The same is also
       true with the atoms in molecules. Some molecules present a state
       of more stable equilibrium than others. Hence from this simple
       comparison it will be at once evident that the stability of
       molecules may vary considerably, that one and the same elements,
       taken in the same number, may give isomerides of different
       stability, and, lastly, that there may exist states of equilibria
       which are so unstable, so ephemeral, that they will only arise
       under particularly special conditions--such, for example, as
       certain hydrates mentioned in the first chapter (_see_ Notes
       57, 67, and others). And if in one case the instability of a
       given state of equilibrium is expressed by its instability with
       a change of temperature or physical state, then in other cases
       it is expressed by the facility with which it decomposes under
       the influence of contact or of the chemical influence of other
       substances.

Besides this, hydrogen peroxide presents another side of the subject
which is not less important, and is much clearer and more general.

Hydrogen unites with oxygen in two degrees of oxidation: water or
hydrogen oxide, and oxygenated water or hydrogen peroxide; for a given
quantity of hydrogen, the peroxide contains twice as much oxygen as does
water. This is a fresh example confirming the correctness of the law
of multiple proportions, to which we have already referred in speaking
of the water of crystallisation of salts. We can now formulate this
law--_the law of multiple proportions_. _If two substances A and B
(either simple or compound), unite together to form several compounds,
A_{n}B_{m}, A_{q}B_{r} ..., then having expressed the compositions
of all these compounds in such a way that the quantity (by weight or
volume) of one of the component parts will be a constant quantity_ A, _it
will be observed that in all the compounds_ AB_{a}, AB_{b} _... the
quantities of the other component part,_ B, _will always be in
commensurable relation: generally in simple multiple proportion--that is,
that a : b ..., or m/n is to r/q as whole numbers, for instance as 2 : 3
or 3 : 4...._

The analysis of water shows that in 100 parts by weight it contains
11·112 parts by weight of hydrogen and 88·888 of oxygen, and the
analysis of peroxide of hydrogen shows that it contains 94·112 parts of
oxygen to 5·888 parts of hydrogen. In this the analysis is expressed,
as analyses generally are, in percentages; that is, it gives the amounts
of the elements in a hundred parts by weight of the substance. The
direct comparison of the percentage compositions of water and hydrogen
peroxide does not give any simple relation. But such a relation is
immediately apparent if we calculate the composition of water and of
hydrogen peroxide, having taken either the quantity of oxygen or the
quantity of hydrogen as a constant quantity--for instance, as unity.
The most simple proportions show that in water there are contained
eight parts of oxygen to one part of hydrogen, and in hydrogen peroxide
sixteen parts of oxygen to one part of hydrogen; or one-eighth part
of hydrogen in water and one-sixteenth part of hydrogen in hydrogen
peroxide to one part of oxygen. Naturally, the analysis does not give
these figures with absolute exactness--it gives them within a certain
degree of error--but they approximate, as the error diminishes, to
that limit which is here given. The comparison of the quantities of
hydrogen and oxygen in the two substances above named, taking one of the
components as a constant quantity, gives an example of the application
of the law of multiple proportions, because water contains eight parts
and hydrogen peroxide sixteen parts of oxygen to one part of hydrogen,
and these figures are commensurable and are in the simple proportion of
1 : 2.

An exactly similar multiple proportion is observed in the composition
of all other well-investigated definite chemical compounds,[30] and
therefore the law of multiple proportions is accepted in chemistry as
the starting point from which other considerations proceed.

  [30] When, for example, any element forms several oxides, they are
       subject to the law of multiple proportions. For a given quantity
       of the non-metal or metal the quantities of oxygen in the
       different degrees of oxidation will stand as 1 : 2, or as 1 : 3,
       or as 2 : 3, or as 2 : 7, and so on. Thus, for instance, copper
       combines with oxygen in at least two proportions, forming the
       oxides found in nature, and called the suboxide and the oxide of
       copper, Cu_{2}O and CuO; the oxide contains twice as much oxygen
       as the suboxide. Lead also presents two degrees of oxidation,
       the oxide and peroxide, and in the latter there is twice as
       much oxygen as in the former, PbO and PbO_{2}. When a base and
       an acid are capable of forming several kinds of salts, normal,
       acid, basic, and anhydro-, it is found that they also clearly
       exemplify the law of multiple proportions. This was demonstrated
       by Wollaston soon after the discovery of the law in question. We
       saw in the first chapter that salts show different degrees of
       combination with water of crystallisation, and that they obey the
       law of multiple proportions. And, more than this, the indefinite
       chemical compounds existing as solutions may, as we saw in the
       same chapter, be brought under the law of multiple proportions
       by the hypothesis that solutions are unstable hydrates formed
       according to the law of multiple proportions, but occurring in
       a state of dissociation. By means of this hypothesis the law of
       multiple proportions becomes still more general, and all the
       aspects of chemical compounds are subject to it. The direction
       of the whole contemporary state of chemistry was determined by
       the discoveries of Lavoisier and Dalton. By endeavouring to
       prove that in solutions we have nothing else than the liquid
       products of the dissociation of definite hydrates, it is my aim
       to bring also this category of indefinite compounds under the
       general principle enunciated by Dalton; just as astronomers have
       discovered a proof and not a negation of the laws of Newton in
       perturbations.

The law of multiple proportions was discovered at the beginning of this
century by John Dalton, of Manchester, in investigating the compounds of
carbon with hydrogen. It appeared that two gaseous compounds of these
substances--marsh gas, CH_{4}, and olefiant gas, C_{2}H_{4}, contain for
one and the same quantity of hydrogen, quantities of carbon which stand
in multiple proportion; namely, marsh gas contains relatively half as
much carbon as olefiant gas. Although the analysis of that time was not
exact, still the accuracy of this law, recognised by Dalton, was further
confirmed by more accurate investigations. On establishing the law of
multiple proportions, Dalton gave a hypothetical explanation for it.
This explanation is based on the atomic theory of matter. In fact, the
law of multiple proportions may be very easily understood by admitting
the atomic structure of matter.

The essence of the atomic theory is that matter is supposed to consist
of an agglomeration of small and indivisible parts--atoms--which do
not fill up the whole space occupied by a substance, but stand apart
from each other, as the sun, planets, and stars do not fill up the
whole space of the universe, but are at a distance from each other.
The form and properties of substances are determined by the position
of their atoms in space and by their state of motion, whilst the
reactions accomplished by substances are understood as redistributions
of the relative positions of atoms and changes in their motion. The
atomic representation of matter arose in very ancient times,[31] and
up to recent times was at variance with the dynamical hypothesis,
which considers matter as only a manifestation of forces. At the
present time, however, the majority of scientific men uphold the atomic
hypothesis, although the present conception of an atom is quite
different from that of the ancient philosophers. An atom at the present
day is regarded rather as an individual or unit which is indivisible by
physical[32] and chemical forces, whilst the atom of the ancients was
actually mechanically and geometrically indivisible. When Dalton (1804)
discovered the law of multiple proportions, he pronounced himself in
favour of the atomic doctrine, because it enables this law to be very
easily understood. If the divisibility of every element has a limit,
namely the atom, then the atoms of elements are the extreme limits of
all divisibility, and if they differ from each other in their nature,
the formation of a compound from elementary matter must consist in the
aggregation of several different atoms into one whole or system of
atoms, now termed _particles or molecules_. As atoms can only combine
in their entire masses, it is evident that not only the law of definite
composition, but also that of multiple proportions, must apply to the
combination of atoms with one another; for one atom of a substance
can combine with one, two, or three atoms of another substance, or in
general one, two, three atoms of one substance are able to combine
with one, two, or three atoms of another; this being the essence of
the law of multiple proportions. Chemical and physical data are very
well explained by the aid of the atomic theory. The displacement of one
element by another follows the law of equivalency. In this case one or
several atoms of a given element take the place of one or several atoms
of another element in its compounds. The atoms of different substances
can be mixed together in the same sense as sand can be mixed with clay.
They do not unite into one whole--_i.e._ there is not a perfect blending
in the one or other case, but only a juxtaposition, a homogeneous whole
being formed from individual parts. This is the first and most simple
method of applying the atomic theory to the explanation of chemical
phenomena.[33]

  [31] Leucippus, Democritus, and especially Lucretius, in the classical
       ages, represented matter as made up of atoms--that is, of parts
       incapable of further division. The geometrical impossibility of
       such an admission, as well as the conclusions which were deduced
       by the ancient atomists from their fundamental propositions,
       prevented other philosophers from following them, and the atomic
       doctrine, like very many others, lived, without being ratified
       by fact, in the imaginations of its followers. Between the
       present atomic theory and the doctrine of the above-named ancient
       philosophers there is naturally a remote historical connection,
       as between the doctrine of Pythagoras and Copernicus, but they
       are essentially different. For us the atom is indivisible, not
       in the geometrical abstract sense, but only in a physical and
       chemical sense. It would be better to call the atoms indivisible
       _individuals_. The Greek atom = the Latin individual, both
       according to the etymology and original sense of the words, but
       in course of time these two words have acquired a different
       meaning. The individual is mechanically and geometrically
       divisible, and only indivisible in a special sense. The earth,
       the sun, a man or a fly are individuals, although geometrically
       divisible. Thus the 'atoms' of contemporary science, indivisible
       in a chemical sense, form those units with which we are concerned
       in the investigation of the natural phenomena of matter, just
       as a man is an indivisible unit in the investigation of social
       relations, or as the stars, planets, and luminaries serve as
       units in astronomy. The formation of the vortex hypothesis,
       in which, as we shall afterwards see, atoms are entire whirls
       mechanically complex, although physico-chemically indivisible,
       clearly shows that the scientific men of our time in holding
       to the atomic theory have only borrowed the word and form of
       expression from the ancient philosophers, and not the essence
       of their atomic doctrine. It is erroneous to imagine that the
       contemporary conceptions of the atomists are nothing but the
       repetition of the metaphysical reasonings of the ancients. To
       show the true meaning of the atomism of the ancient philosophers,
       and the profound difference between their points of argument
       and those of contemporary men of science, I cite the following
       fundamental propositions of Democritus (B.C. 470-380) as the best
       expounder of the atomic doctrine of the ancients:--(1) Nothing
       can proceed from nothing, nothing that exists can disappear or
       be destroyed (and hence matter), and every change only consists
       of a combination or separation. (2) Nothing is accidental, there
       is a reason and necessity for everything. (3) All except atoms
       and vacua is reason and not existence. (4) The atoms, which are
       infinite in number and form, constitute the visible universe by
       their motion, impact, and consequent revolving motion. (5) The
       variety of objects depends only upon a difference in the number,
       form, and order of the atoms of which they are formed, and not
       upon a qualitative difference of their atoms, which only act
       upon each other by pressure and impact. (6) The spirit, like
       fire, consists of minute, spherical, smooth, and very mobile and
       all-penetrating atoms, whose motion forms the phenomenon of life.
       These Democritian, chiefly metaphysical, principles of atomism
       are so essentially different from the principles of the present
       atomic doctrine, which is exclusively applied to explaining
       the phenomena of the external world, that it may be useful to
       mention the essence of the atomic propositions of Boscovitch,
       a Slav who lived in the middle of the eighteenth century, and
       who is regarded as the founder of the modern atomic doctrines
       which, however, did not take hold upon the minds of scientific
       men, and were rarely applied prior to Dalton--_i.e._ until the
       beginning of the nineteenth century. The doctrine of Boscovitch
       was enunciated by him in 1758-1764 in his '_Philosophiæ naturalis
       theoria reducta ad unicam legem virium in natura existentium_.'
       Boscovitch considers matter to be composed of atoms, and the
       atoms to be the points or centres of forces (just as the stars
       and planets may be considered as points of space), acting between
       bodies and their parts. These forces vary with the distance, so
       that beyond a certain very small distance all atoms, and hence
       also their aggregates, are attracted according to Newton's law,
       but at less distances, there alternate wave-like spheres of
       gradually decreasing attraction and increasing (as the distance
       decreases) repulsion, until at last at a minimum distance only
       the repellent action remains. Atoms, therefore, cannot merge
       into each other. Consequently, the atoms are held at a certain
       distance from each other, and therefore occupy space. Boscovitch
       compares the sphere of repulsion surrounding the atoms to
       the spheres of action of firing of a detachment of soldiers.
       According to his doctrine, atoms are indestructible, do not merge
       into each other, have mass, are everlasting and mobile under the
       action of the forces proper to them. Maxwell rightly calls this
       hypothesis the 'extreme' among those existing to explain matter,
       but many aspects of Boscovitch's doctrine repeat themselves
       in the views of our day, with this essential difference, that
       instead of a mathematical point furnished with the properties of
       mass, the atoms are endowed with a corporality, just as the stars
       and planets are corporal, although in certain aspects of their
       interaction they may be regarded as mathematical points. In my
       opinion, the atomism of our day must first of all be regarded
       merely as a convenient method for the investigation of ponderable
       matter. As a geometrician in reasoning about curves represents
       them as formed of a succession of right lines, because such a
       method enables him to analyse the subject under investigation,
       so the scientific man applies the atomic theory as a method of
       analysing the phenomena of nature. Naturally there are people
       now, as in ancient times, and as there always will be, who apply
       reality to imagination, and therefore there are to be found
       atomists of extreme views; but it is not in their spirit that
       we should acknowledge the great services rendered by the atomic
       doctrine to all science, which, while it has been essentially
       independently developed, is, if it be desired to reduce all
       ideas to the doctrines of the ancients, a union of the ancient
       dynamical and atomic doctrines.

  [32] Dalton and many of his successors distinguished the atoms of
       elements and compounds, in which they clearly symbolised the
       difference of their opinion from the representations of the
       ancients. Now only the individuals of the elements, indivisible
       by physical and chemical forces, are termed atoms, and the
       individuals of compounds indivisible under physical changes are
       termed molecules; these are divisible into atoms by chemical
       forces.

  [33] In the present condition of science, either the atomic or the
       dynamical hypothesis is inevitably obliged to admit the existence
       of an invisible and imperceptible motion in matter, without
       which it is impossible to understand either light or heat, or
       gaseous pressure, or any of the mechanical, physical, or chemical
       phenomena. The ancients saw vital motion in animals only, but
       to us the smallest particle of matter, endued with _vis viva_,
       or energy in some degree or other, is incomprehensible without
       self-existent motion. Thus motion has become a conception
       inseparably knit with the conception of matter, and this has
       prepared the ground for the revival of the dynamical hypothesis
       of the constitution of matter. In the atomic theory there has
       arisen that generalising idea by which the world of atoms is
       constructed, like the universe of heavenly bodies, with its
       suns, planets, and meteors, endued with everlasting force of
       motion, forming molecules as the heavenly bodies form systems,
       like the solar system, which molecules are only relatively
       indivisible in the same way as the planets of the solar system
       are inseparable, and stable and lasting as the solar system
       is lasting. Such a representation, without necessitating the
       absolute indivisibility of atoms, expresses all that science can
       require for an hypothetical representation of the constitution
       of matter. In closer proximity to the dynamical hypothesis of
       the constitution of matter is the oft-times revived _vortex
       hypothesis_. Descartes first endeavoured to raise it; Helmholtz
       and Thomson (Lord Kelvin) gave it a fuller and more modern form;
       many scientific men applied it to physics and chemistry. The idea
       of vortex rings serves as the starting point of this hypothesis;
       these are familiar to all as the rings of tobacco smoke, and may
       be artificially obtained by giving a sharp blow to the sides
       of a cardboard box having a circular orifice and filled with
       smoke. Phosphuretted hydrogen, as we shall see later on, when
       bubbling from water always gives very perfect vortex rings in a
       still atmosphere. In such rings it is easy to observe a constant
       circular motion about their axes, and to notice the stability the
       rings possess in their motion of translation. This unchangeable
       mass, endued with a rapid internal motion, is likened to the
       atom. In a medium deprived of friction, such a ring, as is shown
       by theoretical considerations of the subject from a mechanical
       point of view, would be perpetual and unchangeable. The rings
       are capable of grouping together, and in combining, without
       being absolutely indivisible, remain indivisible. The vortex
       hypothesis has been established in our times, but it has not
       been fully developed; its application to chemical phenomena is
       not clear, although not impossible; it does not satisfy a doubt
       in respect to the nature of the space existing between the rings
       (just as it is not clear what exists between atoms, and between
       the planets), neither does it tell us what is the nature of the
       moving substance of the ring, and therefore for the present it
       only presents the germ of an hypothetical conception of the
       constitution of matter; consequently, I consider that it would
       be superfluous to speak of it in greater detail. However, the
       thoughts of investigators are now (and naturally will be in the
       future), as they were in the time of Dalton, often turned to the
       question of the limitation of the mechanical division of matter,
       and the atomists have searched for an answer in the most diverse
       spheres of nature. I select one of the methods attempted, which
       does not in any way refer to chemistry, in order to show how
       closely all the provinces of natural science are bound together.
       Wollaston proposed the investigation of the _atmosphere of the
       heavenly bodies_ as a means for confirming the existence of
       atoms. If the divisibility of matter be infinite, then air must
       extend throughout the entire space of the heavens as it extends
       all over the earth by its elasticity and diffusion. If the
       infinite divisibility of matter be admitted, it is impossible
       that any portion of the whole space of the universe can be
       entirely void of the component parts of our atmosphere. But if
       matter be divisible up to a certain limit only--namely, up to
       the atom--then there _can exist_ a heavenly body void of an
       atmosphere; and if such a body be discovered, it would serve as
       an important factor for the acceptation of the validity of the
       atomic doctrine. The moon has long been considered as such a
       luminary and this circumstance, especially from its proximity to
       the earth, has been cited as the best proof of the validity of
       the atomic doctrine. This proof is apparently (Poisson) deprived
       of some of its force from the possibility of the transformation
       of the component parts of our atmosphere into a solid or liquid
       state at immense heights above the earth's surface, where the
       temperature is exceedingly low; but a series of researches
       (Pouillet) has shown that the temperature of the heavenly
       space is comparatively not so very low, and is attainable by
       experimental means, so that at the low existing pressure the
       liquefaction of the gases of the atmosphere cannot he expected
       even on the moon. Therefore the absence of an atmosphere about
       the moon, if it were not subject to doubt, would be counted as a
       forcible proof of the atomic theory. As a proof of the absence of
       a lunar atmosphere, it is cited that the moon, in its independent
       motion between the stars, when eclipsing a star--that is, when
       passing between the eye and the star--does not show any signs
       of refraction at its edge; the image of the star does not alter
       its position in the heavens on approaching the moon's surface,
       consequently there is no atmosphere on the moon's surface capable
       of refracting the rays of light. Such is the conclusion by which
       the absence of a lunar atmosphere is acknowledged. But this
       conclusion is most feeble, and there are even facts in exact
       contradiction to it, by which the existence of a lunar atmosphere
       may be proved. The entire surface of the moon is covered with a
       number of mountains, having in the majority of cases the conical
       form natural to volcanoes. The volcanic character of the lunar
       mountains was confirmed in October 1866, when a change was
       observed in the form of one of them (the crater Linnea). These
       mountains must be on the edge of the lunar disc. Seen in profile,
       they screen one another and interfere with observations on the
       surface of the moon, so that when looking at the edge of the
       lunar disc we are obliged to make our observations not on the
       moon's surface, but at the summits of the lunar mountains. These
       mountains are higher than those on our earth, and consequently at
       their summits the lunar atmosphere must he exceedingly rarefied
       even if it possess an observable density at the surface. Knowing
       the mass of the moon to be eighty-two times less than the mass
       of the earth, we are able to determine approximately that our
       atmosphere at the moon's surface would be about twenty-eight
       times lighter than it is on the earth, and consequently at the
       very surface of the moon the refraction of light by the lunar
       atmosphere must he very slight, and at the heights of the lunar
       mountains it must be imperceptible, and would be lost within the
       limits of experimental error. Therefore the absence of refraction
       of light at the edge of the moon's disc cannot yet be urged in
       favour of the absence of a lunar atmosphere. There is even a
       series of observations obliging us to admit the existence of
       this atmosphere. These researches are due to Sir John Herschel.
       This is what he writes: 'It has often been remarked that during
       the eclipse of a star by the moon there occurs a peculiar
       optical illusion; it seems as if the star before disappearing
       passed over the edge of the moon and is seen through the lunar
       disc, sometimes for a rather long period of time. I myself have
       observed this phenomenon, and it has been witnessed by perfectly
       trustworthy observers. I ascribe it to optical illusion, but it
       must be admitted that the star might have been seen on the lunar
       disc through some deep ravine on the moon.' Geniller, in Belgium
       (1856), following the opinion of Cassini, Eiler, and others,
       gave an explanation of this phenomenon: he considers it due to
       the refraction of light in the valleys of the lunar mountains
       which occur on the edge of the lunar disc. In fact, although
       these valleys do not probably present the form of straight
       ravines, yet it may sometimes happen that the light of a star
       is so refracted that its image might he seen, notwithstanding
       the absence of a direct path for the light-rays. He then goes
       on to remark that the density of the lunar atmosphere must be
       variable in different parts, owing to the very long nights on
       the moon. On the dark, or non-illuminated portion, owing to
       these long nights, which last thirteen of our days and nights,
       there must be excessive cold, and hence a denser atmosphere,
       while, on the contrary, on the illuminated portion the atmosphere
       must be much more rarefied. This variation in the temperature
       of the different parts of the moon's surface explains also the
       absence of clouds, notwithstanding the possible presence of air
       and aqueous vapour, on the visible portion of the moon. The
       presence of an atmosphere round the sun and planets, judging from
       astronomical observations, may be considered as fully proved. On
       Jupiter and Mars even bands of clouds may be distinguished. Thus
       the atomic doctrine, admitting a finite mechanical divisibility
       only, must he, as yet at least, only accepted as a means, similar
       to that means which a mathematician employs when he breaks up
       a continuous curvilinear line into a number of straight lines.
       There is a simplicity of representation in atoms, but there is no
       absolute necessity to have recourse to them. The conception of
       the individuality of the parts of matter exhibited in chemical
       elements only is necessary and trustworthy.

A certain number of atoms _n_ of an element A in combining with several
atoms _m_ of another element B give a compound A_{_n_} B_{_m_}, each
molecule of which will contain the atoms of the elements A and B
in this ratio, and therefore the compound will present a _definite
composition_, expressed by the formula A_{_n_}B_{_m_}, where A and B
are the weights of the atoms and _n_ and _m_ their relative number. If
the same elements A and B, in addition to A_{_n_}B_{_m_}, also yield
another compound A_{_r_}B_{_q_}, then by expressing the composition of
the first compound by A_{_nr_}B_{_mr_} (and this is the same composition
as A_{_n_}B_{_m_}), and of the second compound by A_{_rn_}B_{_qn_},
we have the law of multiple proportions, because for a given quantity
of the first element, A_{_rn_}, there occur quantities of the second
element bearing the same ratio to each other as _mr_ is to _qn_; and
as _m_, _r_, _q_, and _n_ are whole numbers, their products are also
whole numbers, and this is expressed by the law of multiple proportion.
Consequently the atomic theory is in accordance with and evokes
the first laws of definite chemical compounds: the law of definite
composition and the law of multiple proportions.

So, also, is the relation of the atomic theory to the third law of
definite chemical compounds, the _law of reciprocal combining weights_,
which is as follows:--If a certain weight of a substance C combine with
a weight _a_ of a substance A, and with a weight _b_ of a substance B,
then, also, the substances A and B will combine together in quantities
_a_ and _b_ (or in multiples of them). This should be the case from
the conception of atoms. Let A, B, and C be the weights of the atoms
of the three substances, and for simplicity of reasoning suppose that
combination takes place between single atoms. It is evident that if
the substance gives AC and BC, then the substances A and B will give a
compound AB, or their multiple, A_{_n_}B_{_m_}. And so it is in reality
in nature.

Sulphur combines with hydrogen and with oxygen. Sulphuretted hydrogen
contains thirty-two parts by weight of sulphur to two parts by weight
of hydrogen; this is expressed by the formula H_{2}S. Sulphur dioxide,
SO_{2}, contains thirty-two parts of sulphur and thirty-two parts of
oxygen, and therefore we conclude, from the law of combining weights,
that oxygen and hydrogen will combine in the proportion of two parts
of hydrogen and thirty-two parts of oxygen, or multiple numbers of
them. And we have seen this to be the case. Hydrogen peroxide contains
thirty-two parts of oxygen, and water sixteen parts, to two parts
of hydrogen; and so it is in all other cases. This consequence of
the atomic theory is in accordance with nature, with the results of
analysis, and is one of the most important laws of chemistry. It is
a law, because it indicates the _relation between_ the weights of
substances entering into chemical combination. Further, it is an
eminently exact law, and not an approximate one. The law of combining
weights is a law of nature, and by no means an hypothesis, for even
if the entire theory of atoms be refuted, still the laws of multiple
proportions and of combining weights will remain, inasmuch as they deal
with facts. They may be guessed at from the sense of the atomic theory,
and historically the law of combining weights is intimately connected
with this theory; but they are not identical, but only connected, with
it. The law of combining weights is formulated with great ease, and is
an immediate consequence of the atomic theory; without it, it is even
difficult to understand. Data for its evolution existed previously, but
it was not formulated until those data were interpreted by the atomic
theory, an hypothesis which up to the present time has contradicted
neither experiment nor fact, and is useful and of general application.
Such is the nature of hypotheses. They are indispensable to science;
they bestow an order and simplicity which are difficultly attainable
without their aid. The whole history of science is a proof of this.
And therefore it may be truly said that it is better to hold to an
hypothesis which may afterwards prove untrue than to have none at all.
Hypotheses facilitate scientific work and render it consistent. In the
search for truth, like the plough of the husbandman, they help forward
the work of the labourer.




                                CHAPTER V

                            NITROGEN AND AIR


Gaseous _nitrogen_ forms about four-fifths (by volume) of the
atmosphere; consequently the air contains an exceedingly large
mass of it. Whilst entering in so considerable a quantity into the
composition of air, nitrogen does not seem to play any active part in
the atmosphere, the chemical action of which is mainly dependent on the
oxygen it contains. But this is not an entirely correct idea, because
animal life cannot exist in pure oxygen, in which animals pass into an
abnormal state and die; and the nitrogen of the air, although slowly,
forms diverse compounds, many of which play a most important part in
nature, especially in the life of organisms. However, neither plants[1]
nor animals directly absorb the nitrogen of the air, but take it up from
already prepared nitrogenous compounds; further, plants are nourished by
the nitrogenous substances contained in the soil and water, and animals
by the nitrogenous substances contained in plants and in other animals.
Atmospheric electricity is capable of aiding the passage of gaseous
nitrogen into nitrogenous compounds, as we shall afterwards see, and
the resultant substances are carried to the soil by rain, where they
serve for the nourishment of plants. Plentiful harvests, fine crops of
hay, vigorous growth of trees--other conditions being equal--are only
obtained when the soil contains _ready prepared nitrogenous compounds_,
consisting either of those which occur in air and water, or of the
residues of the decomposition of other plants or animals (as in manure).
The nitrogenous substances contained in animals have their origin in
those substances which are formed in plants. Thus the nitrogen of the
atmosphere is the origin of all the nitrogenous substances occurring in
animals and plants, although not directly so, but after first combining
with the other elements of air.

  [1] See Note 15 bis.

The nitrogenous compounds which enter into the composition of plants and
animals are of primary importance; no vegetable or animal cell--that
is, the elementary form of organism--exists without containing a
nitrogenous substance, and moreover organic life manifests itself
primarily in these nitrogenous substances. The germs, seeds, and
those parts by which cells multiply themselves abound in nitrogenous
substances; the sum total of the phenomena which are proper to organisms
depend primarily on the chemical properties of the nitrogenous
substances which enter into their composition. It will be sufficient,
for instance, to point out the fact that vegetable and animal organisms,
clearly distinguishable as such, are characterised by a different degree
of energy in their nature, and at the same time by a difference in the
amount of nitrogenous substances they contain. In plants, which compared
with animals possess but little activity, being incapable of independent
movement, &c., the amount of nitrogen is very much less than in animals,
whose tissues are almost exclusively formed of nitrogenous substances.
It is remarkable that the nitrogenous parts of plants, chiefly of the
lower orders, sometimes present both forms and properties which approach
to those of animal organisms; for example, the zoospores of sea-weeds,
or those parts by means of which the latter multiply themselves. These
zoospores on leaving the sea-weed in many respects resemble the lower
orders of animal life, having, like the latter, the property of moving.
They also approach the animal kingdom in their composition, their outer
coating containing nitrogenous matter. Directly the zoospore becomes
covered with that non-nitrogenous or cellular coating which is proper to
all the ordinary cells of plants, it loses all resemblance to an animal
organism and becomes a small plant. It may be thought from this that the
cause of the difference in the vital processes of animals and plants
is the different amount of nitrogenous substances they contain. The
nitrogenous substances which occur in plants and animals appertain to a
series of exceedingly complex and very changeable chemical compounds;
their elementary composition alone shows this; besides nitrogen, they
contain carbon, hydrogen, oxygen, and sulphur. Being distinguished by a
very great instability under many conditions in which other compounds
remain unchanged, these substances are fitted for those perpetual
changes which form the first condition of vital activity. These complex
and changeable nitrogenous substances of the organism are called
_proteïd substances_. The white of eggs is a familiar example of such a
substance. They are also contained in the flesh of animals, the curdy
elements of milk, the glutinous matter of wheaten flour, or so-called
gluten, which forms the chief component of macaroni, &c.

Nitrogen occurs in the earth's crust, in compounds either forming the
remains of plants and animals, or derived from the nitrogen of the
atmosphere as a consequence of its combination with the other component
parts of the air. It is not found in other forms in the earth's crust;
so that nitrogen must be considered, in contradistinction to oxygen,
as an element which is purely superficial, and does not extend to the
depths of the earth.[1 bis]

  [1 bis] The reason why there are no other nitrogenous substances within
      the earth's mass beyond those which have come there with the
      remains of organisms, and from the air with rain-water, must be
      looked for in two circumstances. In the first place, in the
      instability of many nitrogenous compounds, which are liable to
      break up with the formation of gaseous nitrogen; and in the second
      place in the fact that the salts of nitric acid, forming the
      product of the action of air on many nitrogenous and especially
      organic compounds, are very soluble in water, and on penetrating
      into the depths of the earth (with water) give up their oxygen. The
      result of the changes of the nitrogenous organic substances which
      fall into the earth is without doubt frequently, if not invariably,
      the formation of gaseous nitrogen. Thus the gas evolved from coal
      always contains much nitrogen (together with marsh gas, carbonic
      anhydride, and other gases).

_Nitrogen is liberated_ in a free state in the decomposition of the
_nitrogenous organic substances_ entering into the composition of
organisms--for instance, on their combustion. All organic substances
burn when heated to redness with oxygen (or substances readily yielding
it, such as oxide of copper); the oxygen combines with the carbon,
sulphur, and hydrogen, and the nitrogen is evolved in a free state,
because at a high temperature it does not form any stable compound,
but remains uncombined. Carbonic anhydride and water are formed from
the carbon and hydrogen respectively, and therefore to obtain pure
nitrogen it is necessary to remove the carbonic anhydride from the
gaseous products obtained. This may be done very easily by the action of
alkalis--for instance, caustic soda. The amount of nitrogen in organic
substances is determined by a method founded on this.

It is also very easy to obtain _nitrogen from air_, because oxygen
combines with many substances. Either phosphorus or metallic copper is
usually employed for removing the oxygen from air, but, naturally, a
number of other substances may also be used. If a small saucer on which
a piece of phosphorus is laid be placed on a cork floating on water, and
the phosphorus be lighted, and the whole covered with a glass bell jar,
then the air under the jar will be deprived of its oxygen, and nitrogen
only will remain, owing to which, on cooling, the water will rise to
a certain extent in the bell jar. The same object (procuring nitrogen
from air) is attained much more conveniently and perfectly by passing
air through a red-hot tube containing copper filings. At a red heat,
metallic copper combines with oxygen and gives a black powder of copper
oxide. If the layer of copper be sufficiently long and the current of
air slow, all the oxygen will be absorbed, and nitrogen alone will pass
from the tube.[2]

  [2] Copper (best as turnings, which present a large surface) absorbs
      oxygen, forming CuO, at the ordinary temperature in the presence
      of solutions of acids, or, better still, in the presence of a
      solution of ammonia, when it forms a bluish-violet solution of
      oxide of copper in ammonia. Nitrogen is very easily procured by
      this method. A flask filled with copper turnings is closed with a
      cork furnished with a funnel and stopcock. A solution of ammonia
      is poured into the funnel, and caused to drop slowly upon the
      copper. If at the same time a current of air be slowly passed
      through the flask (from a gasholder), then all the oxygen will
      be absorbed from it and the nitrogen will pass from the flask.
      It should be washed with water to retain any ammonia that may be
      carried off with it.

Nitrogen may also be procured from many of its _compounds with oxygen[3]
and hydrogen_,[4] but the best fitted for this purpose is a saline
mixture containing, on the one hand, a compound of nitrogen with oxygen,
termed nitrous anhydride, N_{2}O_{3}, and on the other hand, ammonia,
NH_{3}--that is, a compound of nitrogen with hydrogen. By heating such a
mixture, the oxygen of the nitrous anhydride combines with the hydrogen
of the ammonia, forming water, and gaseous nitrogen is evolved, 2NH_{3}
+ N_{2}O_{3} = 3H_{2}O + N_{4}. Nitrogen is procured by this method in
the following manner:--A solution of caustic potash is saturated with
nitrous anhydride, by which means potassium nitrite is formed. On the
other hand, a solution of hydrochloric acid saturated with ammonia is
prepared; a saline substance called sal-ammoniac, NH_{4}Cl, is thus
formed in the solution. The two solutions thus prepared are mixed
together and heated. Reaction takes place according to the equation
KNO_{2} + NH_{4}Cl = KCl + 2H_{2}O + N_{2}. This reaction proceeds in
virtue of the fact that potassium nitrite and ammonium chloride are
salts which, on interchanging their metals, give potassium chloride
and ammonium nitrite, NH_{4}NO_{2}, which breaks up into water and
nitrogen. This reaction does not take place without the aid of heat,
but it proceeds very easily at a moderate temperature. Of the resultant
substances, the nitrogen only is gaseous. Pure nitrogen may be obtained
by drying the resultant gas and passing it through a solution of
sulphuric acid (to absorb a certain quantity of ammonia which is evolved
in the reaction).[4 bis]

  [3] The oxygen compounds of nitrogen (for example, N_{2}O, NO, NO_{2})
      are decomposed at a red heat by themselves, and under the action
      of red-hot copper, iron, sodium, &c., they give up their oxygen
      to the metals, leaving the nitrogen free. According to Meyer and
      Langer (1885), nitrous oxide, N_{2}O, decomposes below 900°,
      although not completely.

  [4] Chlorine and bromine (in excess), as well as bleaching powder
      (hypochlorites), take up the hydrogen from ammonia, NH_{3},
      leaving nitrogen. Nitrogen is best procured from ammonia by the
      action of a solution of sodium hypobromite on solid sal-ammoniac.

  [4 bis] Lord Rayleigh in 1894, when determining the weight of a volume
      of carefully purified nitrogen by weighing it in one and the same
      globe, found that the gas obtained from air, by the action of
      incandescent copper (or iron or by removing the oxygen by ferrous
      oxide) was always 1/200 heavier than the nitrogen obtained from its
      compounds, for instance, from the oxide or suboxide of nitrogen,
      decomposed by incandescent pulverulent iron or from the ammonia
      salt of nitrous acid. For the nitrogen procured from air, he
      obtained, at 0° and 760·4 mm. pressure, a weight = 2·310 grms.,
      while for the nitrogen obtained from its compounds, 2·299 grms.
      This difference of about 1/200 could not be explained by the
      nitrogen not having been well purified, or by inaccuracy of
      experiment, and was the means for the remarkable discovery of the
      presence of a heavy gas in air, which will be mentioned in Note
      16 bis.

Nitrogen is a gaseous substance which does not differ much in physical
properties from air; its density, referred to hydrogen, is approximately
equal to 14--that is, it is slightly lighter than air, its density
referred to air being 0·972; one litre of nitrogen weighs 1·257 gram.
Nitrogen mixed with oxygen, which is slightly heavier than air, forms
air. It is a gas which, like oxygen and hydrogen, is liquefied with
difficulty, and is but little soluble in water and other liquids. Its
absolute boiling point[5] is about -140°; above this temperature it is
not liquefiable by pressure, and at lower temperatures it remains a
gas at a pressure of 50 atmospheres. Liquid nitrogen boils at -193°,
so that it may be employed as a source of great cold. At about -203°,
in vaporising under a decrease of pressure, nitrogen solidifies into
a colourless snow-like mass. Nitrogen does not burn,[5 bis] does not
support combustion, is not absorbed by any of the reagents used in gas
analysis, at least at the ordinary temperature--in a word, it presents
a whole series of negative chemical properties; this is expressed by
saying that this element has no energy for combination. Although it
is capable of forming compounds both with oxygen and hydrogen as well
as with carbon, yet these compounds are only formed under particular
circumstances, to which we will directly turn our attention. At a red
heat nitrogen combines with boron, titanium, and silicon, barium,
magnesium, &c., forming very stable nitrogenous compounds,[6] whose
properties are entirely different from those of nitrogen with hydrogen,
oxygen and carbon. However, the combination of nitrogen with carbon,
although it does not take place directly between the elements at a
red heat, yet proceeds with comparative ease by heating a mixture of
charcoal with an alkaline carbonate, especially potassium carbonate or
barium carbonate, to redness, carbo-nitrides or cyanides of the metals
being formed; for instance, K_{2}CO_{3} + 4C + N_{2} = 2KCN + 3CO.[7]

  [5] See Chapter II. Note 29.

  [5 bis] See Note 11 bis.

  [6] The combination of boron with nitrogen is accompanied by the
      evolution of sufficient heat to raise the mass to redness;
      titanium combines so easily with nitrogen that it is difficult
      to obtain it free from that element; magnesium easily absorbs
      nitrogen at a red heat. It is a remarkable and instructive fact
      that these compounds of nitrogen are very stable and non-volatile.
      Carbon (C = 12) with nitrogen gives cyanogen, C_{2}N_{2}, which
      is gaseous and very unstable, and whose molecule is not large,
      whilst boron (B = 11) forms a nitrogenous compound which is solid,
      non-volatile, and very stable. Its composition, BN, is similar
      to that of cyanogen, but its molecular weight, B_{n}N_{n}, is
      probably greater. Its composition, like that of N_{2}Mg_{3},
      NNa_{3}, N_{2}Hg_{3} and of many of the metallic nitrides,
      corresponds to ammonia with the substitution of all its hydrogen
      by a metal. In my opinion, a detailed study of the transformations
      of the nitrides now known, should lead to the discovery of many
      facts in the history of nitrogen.

  [7] This reaction, so far as is known, does not proceed beyond a
      certain limit, probably because cyanogen, CN, itself breaks up
      into carbon and nitrogen.

Nitrogen is found with oxygen in the air, but they do not readily
combine. Cavendish, however, in the last century, showed that _nitrogen
combines with oxygen under the influence of a series of electric
sparks_. Electric sparks in passing through a moist[8] mixture
of nitrogen and oxygen cause these elements to combine, forming
reddish-brown fumes of oxides of nitrogen,[9] which form nitric
acid,[10] NHO_{3}. The presence of the latter is easily recognised,
not only from its reddening litmus paper, but also from its acting as
a powerful oxidiser even of mercury. Conditions similar to these occur
in nature, during a thunderstorm or in other electrical discharges
which take place in the atmosphere; whence it may be taken for granted
that air and rain-water always contain traces of nitric and nitrous
acids.[11] Besides which Crookes (1892) showed that under certain
circumstances and when electricity of high potential[11 bis] passes
through the air, the combination of nitrogen with oxygen is accompanied
by the formation of a true flame. This was also observed previously
(1880) during the passage of electrical discharges through the air.

  [8] Frémy and Becquerel took dry air, and observed the formation of
      brown vapours of oxides of nitrogen on the passage of sparks.

  [9] If a mixture of one volume of nitrogen and fourteen volumes of
      hydrogen be burnt, then water and a considerable quantity of
      nitric acid are formed. It may be partly due to this that a
      certain quantity of nitric acid is produced in the slow oxidation
      of nitrogenous substances in an excess of air. This is especially
      facilitated by the presence of an alkali with which the nitric
      acid formed can combine. If a galvanic current be passed through
      water containing the nitrogen and oxygen of the air in solution,
      then the hydrogen and oxygen set free combine with the nitrogen,
      forming ammonia and nitric acid.

      When copper is oxidised at the expense of the air at the ordinary
      temperature in the presence of ammonia, oxygen is absorbed, not
      only for combination with the copper, but also for the formation
      of nitric acid.

      The combination of nitrogen with oxygen, even, for example, by
      the action of electric sparks, is not accompanied by an explosion
      or rapid combination, as in the action of a spark on a mixture
      of oxygen and hydrogen. This is explained by the fact that heat
      is not evolved in the combination of nitrogen with oxygen, but
      is absorbed--an expenditure of energy is required, there is no
      evolution of energy. In fact, there will not be the transmission
      of heat from particle to particle which occurs in the explosion
      of detonating gas. Each spark will aid the formation of a certain
      quantity of the compound of oxygen and nitrogen, but will not
      excite the same in the neighbouring particles. In other words, the
      combination of hydrogen with oxygen is an exothermal reaction, and
      the combination of nitrogen with oxygen an endothermal reaction.

      A condition particularly favourable for the oxidation of nitrogen
      is the explosion of detonating gas and air if the former be
      _in excess_. If a mixture of two volumes of detonating gas and
      one volume of air be exploded, then one-tenth of the air is
      converted into nitric acid, and consequently after the explosion
      has taken place there remain only nine-tenths of the volume of
      air originally taken. If a large proportion of air be taken--for
      instance, four volumes of air to two volumes of detonating
      gas--then the temperature of the explosion is lowered, the
      volume of air taken remains unchanged, and no nitric acid is
      formed. This gives a rule to be observed in making use of the
      eudiometer--namely that to weaken the force of the explosion
      not less than an equal volume of air should be added to the
      explosive mixture. On the other hand a large excess must not be
      taken as no explosion would then ensue (_see_ Chapter III. Note
      34). Probably in the future means will be found for obtaining
      compounds of nitrogen on a large industrial scale by the aid of
      electric discharges, and by making use of the inexhaustible mass
      of nitrogen in the atmosphere.

  [10] In reality nitric oxide, NO, is first formed, but with oxygen and
       water it gives (brown fumes) nitrous anhydride, which, as we
       shall afterwards learn, in the presence of water and oxygen gives
       nitric acid.

  [11] The nitric acid contained in the soil, river water (Chapter I.,
       Note 2), wells, &c., proceeds (like carbonic anhydride) from the
       oxidation of organic compounds which have fallen into water,
       soil, &c.

  [11 bis] Crookes employed a current of 15 ampères and 65 volts, and
       passed it through an induction coil with 330 vibrations per
       second, and obtained a flame between the poles placed at a
       distance of 46 mm. which after the appearance of the arc and flame
       could be increased to 200 mm. A platinum wire fused in the flame.

Further observations showed that under the influence of electrical
discharges,[12] silent as well as with sparks, nitrogen is able to enter
into many reactions with hydrogen and with many hydrocarbons; although
these reactions cannot be effected by exposure to a red heat. Thus,
for instance, a series of electric sparks passed through a mixture of
nitrogen and hydrogen causes them to combine and _form ammonia_[13] or
nitrogen hydride, NH_{3}, composed of one volume of nitrogen and three
volumes of hydrogen. This combination is limited to the formation of
6 per cent. of ammonia, because ammonia is decomposed, although not
entirely (94/100) by electric sparks. This signifies that under the
influence of an electrical discharge the reaction NH_{3} = N + 3H is
reversible, consequently it is a dissociation, and in it a state of
equilibrium is arrived at. The equilibrium may be destroyed by the
addition of gaseous hydrochloric acid, HCl, because with ammonia it
forms a solid saline compound, sal-ammoniac, NH_{4}Cl, which (being
formed from a gaseous mixture of 3H, N, and HCl) fixes the ammonia.
The remaining mass of nitrogen and hydrogen, under the action of the
sparks, again forms ammonia, and in this manner _solid sal-ammoniac is
obtained to the end by the action of a series of electric sparks on a
mixture of gaseous_ N, H_{3}, _and_ HCl.[14] Berthelot (1876) showed
that under the action of a silent discharge many non-nitrogenous organic
substances (benzene, C_{6}H_{6}, cellulose in the form of paper, resin,
glucose, C_{6}H_{10}O_{5}, and others) absorb nitrogen and form complex
nitrogenous compounds, which are capable, like albuminous substances, of
evolving their nitrogen as ammonia when heated with alkalis.[15]

  [12] This property of nitrogen, which under normal conditions is
       inactive, leads to the idea that under the influence of an
       electric discharge gaseous nitrogen changes in its properties;
       if not permanently like oxygen (electrolysed oxygen or ozone
       does not react on nitrogen, according to Berthelot), it may be
       temporarily at the moment of the action of the discharge, just
       as some substances under the action of heat are permanently
       affected (that is, when once changed remain so--for instance,
       white phosphorus passes into red, &c.), whilst others are only
       temporarily altered (the dissociation of S_{6} into S_{2} or
       of sal-ammoniac into ammonia and hydrochloric acid). Such a
       proposition is favoured by the fact that nitrogen gives two kinds
       of spectra, with which we shall afterwards become acquainted. It
       may be that the molecules N_{2} then give less complex molecules,
       N containing one atom, or form a complex molecule N_{3}, like
       oxygen in passing into ozone. Probably under a silent discharge
       the molecules of oxygen, O_{2}, are partly decomposed and the
       individual atoms O combine with O_{2}, forming ozone, O_{3}.

  [13] This reaction, discovered by Chabrié and investigated by Thénard,
       was only rightly understood when Deville applied the principles
       of dissociation to it.

  [14] The action of nitrogen on acetylene (Berthelot) resembles this
       reaction. A mixture of these gases under the influence of a
       silent discharge gives hydrocyanic acid, C_{2}H_{2} + N_{2} =
       2CNH. This reaction cannot proceed beyond a certain limit because
       it is reversible.

  [15] Berthelot successfully employed electricity of even feeble
       potential in these experiments, which fact led him to think that
       in nature, where the action of electricity takes place very
       frequently, a part of the complex nitrogenous substances may
       proceed from the gaseous nitrogen of the air by this method.

       As the nitrogenous substances of organisms play a very important
       part in them (organic life cannot exist without them), and as
       the nitrogenous substances introduced into the soil are capable
       of invigorating its crops (of course in the presence of the
       other nourishing principles required by plants), the question
       of the means of converting the atmospheric nitrogen into the
       nitrogenous compounds of the soil, or into _assimilable nitrogen_
       capable of being absorbed by plants and of forming complex
       (albuminous) substances in them, is one of great theoretical and
       practical interest. The artificial (technical) conversion of the
       atmospheric nitrogen into nitrogenous compounds, notwithstanding
       repeated attempts, cannot yet be considered as fulfilled in a
       practical remunerative manner although its possibility is already
       evident. Electricity will probably aid in solving this very
       important practical problem. When the theoretical side of the
       question is further advanced, then without doubt an advantageous
       means will be found for the manufacture of nitrogenous substances
       from the nitrogen of the air; and this is needed, before all,
       for the agriculturist, to whom nitrogenous fertilisers form an
       expensive item, and are more important than all other manures.

       One thousand tons of farmyard manure do not generally contain
       more than four tons of nitrogen in the form of complex
       nitrogenous substances, and this amount of nitrogen is contained
       in twenty tons of ammonium sulphate, therefore the effect of
       a mass of farmyard manure in respect to the introduction of
       nitrogen may be produced by small quantities of artificial
       nitrogenous fertilisers (_see_ Note 15 bis).

By such indirect methods does the gaseous nitrogen of the atmosphere
yield its primary compounds, in which form it enters into plants, and
is elaborated in them into complex albuminous substances.[15 bis] But,
starting from a given compound of nitrogen with hydrogen or oxygen,
we may, without the aid of organisms, obtain, as will afterwards be
partially indicated, most diverse and complex nitrogenous substances,
which cannot by any means be formed directly from gaseous nitrogen. In
this we see an example not only of the difference between an element in
the free state and an intrinsic element, but also of those circuitous
or _indirect methods_ by which substances are formed in nature. The
discovery, prognostication, and, in general, the study of such indirect
methods of the preparation and formation of substances forms one of the
existing problems of chemistry. From the fact that A does not act at all
on B, it must not be concluded that a compound AB is not to be formed.
The substances A and B contain atoms which occur in AB, but their state
or the nature of their motion may not be at all that which is required
for the formation of AB, and in this substance the chemical state of
the elements may be as different as the state of the atoms of oxygen in
ozone and in water. Thus free nitrogen is inactive; but in its compounds
it very easily enters into changes and is distinguished by great
activity. An acquaintance with the compounds of nitrogen confirms this.
But, before entering on this subject, let us consider air as a mass
containing free nitrogen.

  [15 bis] Although the numerous, and as far as possible accurate and
       varied researches made in the physiology of plants have proved
       that the higher forms of plants are not capable of directly
       absorbing the nitrogen of the atmosphere and converting it into
       complex albuminous substances, still it has been long and
       repeatedly observed that the amount of nitrogenous substances in
       the soil is increased by the cultivation of plants of the bean
       (leguminous) family such as pea, acacia, &c. A closer study of
       these plants has shown that this is connected with the formation
       of peculiar nodular swellings in their roots caused by the growth
       of peculiar micro-organisms (bacteria) which cohabit the soil with
       the roots, and are capable of absorbing nitrogen from the air,
       _i.e._ of converting it into assimilated nitrogen. This branch of
       plant physiology, which forms another proof of the important part
       played by micro-organisms in nature, cannot be discussed in this
       work, but it should be mentioned, since it is of great theoretical
       and practical interest, and, moreover, phenomena of this kind,
       which have recently been discovered, promise to explain, to some
       extent at least, certain of the complex problems concerning the
       development of life on the earth.

Judging from what has been already stated, it will be evident that
_atmospheric air_[16] contains a mixture of several gases and vapours.
Some of them are met with in it in nearly constant proportions, whilst
others, on the contrary, are very variable in their amount. The chief
component parts of air, placed in the order of their relative amounts,
are the following: nitrogen,[16 bis] oxygen, aqueous vapour, carbonic
anhydride, nitric acid, salts of ammonia, oxides of nitrogen, and also
ozone, hydrogen peroxide, and complex organic nitrogenous substances.
Besides these, air generally contains water, as spray, drops, and snow,
and particles of solids, perhaps of cosmic origin in certain instances,
but in the majority of cases proceeding from the mechanical translation
of solid particles from one locality to another by the wind. These
small solid and liquid particles (having a large surface in proportion
to their weight) are suspended in air as solid matter is suspended in
turbid water; they often settle on the surface of the earth, but the
air is never entirely free from them because they are never in a state
of complete rest. Then, air not unfrequently contains incidental traces
of various substances as everyone knows by experience. These incidental
substances sometimes belong to the order of those which act injuriously,
the germs of lower organisms--for instance of moulds--and the class of
carriers of infectious diseases.

  [16] Under the name of atmospheric air the chemist and physicist
       understand ordinary air containing nitrogen and oxygen only,
       notwithstanding that the other component parts of air have a very
       important influence on the living matter of the earth's surface.
       That air is so represented in science is based on the fact that
       only the two components above-named are met with in air in a
       constant quantity, whilst the others are variable. The solid
       impurities may be separated from air required for chemical or
       physical research by simple filtration through a long layer of
       cotton-wool placed in a tube. Organic impurities are removed by
       passing the air through a solution of potassium permanganate. The
       carbonic anhydride contained in air is absorbed by alkalis--best
       of all, soda-lime, which in a dry state in porous lumps absorbs
       it with exceeding rapidity and completeness. Aqueous vapour
       is removed by passing the air over calcium chloride, strong
       sulphuric acid, or phosphoric anhydride. Air thus purified is
       accepted as containing only nitrogen and oxygen, although in
       reality it still contains a certain quantity of hydrogen and
       hydrocarbons, from which it may be purified by passing over
       copper oxide heated to redness. The copper oxide then oxidises
       the hydrogen and hydrocarbons--it burns them, forming water and
       carbonic anhydride, which may be removed as above described. When
       it is said that in the determination of the density of gases the
       weight of air is taken as unity, it is understood to be such air,
       containing only nitrogen and oxygen.

  [16 bis] Thanks to the remarkable discovery made in the summer of 1894
       by Lord Rayleigh and Prof. Ramsay, the well-known component
       elements of air must now he supplemented by 1 p.c. (by volume) of
       a heavy gas (density about 19, H = 1), inactive like nitrogen,
       which was discovered in the researches made by Lord Rayleigh on
       the density of nitrogen as mentioned in note 4 bis. Up to the
       present time this gas has been always determined together with
       nitrogen, because it combines with neither the hydrogen in the
       eudiometer nor with the copper in the gravimetric method of
       determining the composition of air, and therefore has always
       remained with the nitrogen. It has been possible to separate it
       from nitrogen since magnesium absorbs nitrogen at a red heat,
       while this gas remains unabsorbed, and was found to have a density
       nearly one and a half time greater than that of nitrogen (is it
       not a polymer of nitrogen, N_{3}?). It is now known also that this
       gas gives a luminous spectrum, which contains the bright blue line
       observed in the spectrum of nitrogen. Owing to the fact that it is
       an exceedingly inert substance, even more so than nitrogen, it has
       been termed Argon. Further reference will be made to it in the
       Appendix.

In the air of the various countries of the earth, at different
longitudes and at different altitudes above its surface, on the ocean
or on the dry land--in a word, in the air of most diverse localities
of the earth--the oxygen and nitrogen are found everywhere to be in a
constant ratio. This is, moreover, self-evident from the fact that the
air constantly diffuses (intermixes by virtue of the internal motion of
the gaseous particles) and is also put into motion and intermixed by
the wind, by which processes it is equalised in its composition over
the entire surface of the earth. In those localities where the air is
subject to change, and is in a more or less enclosed space, or, at any
rate, in an unventilated space, it may alter very considerably in its
composition. For this reason the air in dwellings, cellars, and wells,
in which there are substances absorbing oxygen, contains less of this
gas, whilst the air on the surface of standing water, which abounds in
the lower orders of plant life evolving oxygen, contains an excess of
this gas.[17] The constant composition of air over the whole surface of
the earth has been proved by a number of most careful researches.[18]

  [17] As a further proof of the fact that certain circumstances may
       change the composition of air, it will be enough to point out
       that the air contained in the cavities of glaciers contains only
       up to 10 p.c. of oxygen. This depends on the fact that at low
       temperatures oxygen is much more soluble in snow-water and snow
       than nitrogen. When shaken up with water the composition of air
       should change, because the water dissolves unequal quantities of
       oxygen and nitrogen. We have already seen (Chapter I.) that the
       air boiled off from water saturated at about 0° contains about
       thirty-five volumes of oxygen and sixty-five volumes of nitrogen,
       and we have considered the reason of this.

  [18] The analysis of air by weight conducted by Dumas and Boussingault
       in Paris, which they repeated many times between April 27 and
       September 22, 1841, under various conditions of weather, showed
       that the amount by weight of oxygen only varies between 22·89
       p.c. and 23·08 p.c., the average amount being 23·07 p.c. Brunner,
       at Bern in Switzerland, and Bravais, at Faulhorn in the Bernese
       Alps, at a height of two kilometres above the level of the sea,
       Marignac at Geneva, Lewy at Copenhagen, and Stas at Brussels,
       have analysed the air by the same methods, and found that its
       composition does not exceed the limits determined for Paris.
       The most recent determinations (with an accuracy of ±0·05 p.c.)
       confirm the conclusion that the composition of the atmosphere is
       constant.

       As there are some grounds (which will be mentioned shortly) for
       considering that the composition of the air at great altitudes
       is slightly different from that at attainable heights--namely,
       that it is richer in the lighter nitrogen--several fragmentary
       observations made at Munich (Jolly, 1880) gave reason for
       thinking that in the upward currents (that is in the region of
       minimum barometric pressure or at the centres of meteorological
       cyclones) the air is richer in oxygen than in the descending
       currents of air (in the regions of anticyclones or of barometric
       maxima); but more carefully conducted observations showed
       this supposition to be incorrect. Improved methods for the
       analysis of air have shown that certain slight variations in
       its composition do actually occur, but in the first place they
       depend on incidental local influences (on the passage of the air
       over mountains and large surfaces of water, regions of forest
       and herbage, and the like), and in the second place are limited
       to quantities which are scarcely distinguishable from possible
       errors in the analyses. The researches made by Kreisler in
       Germany (1885) are particularly convincing.

       The considerations which lead to the supposition that the
       atmosphere at great altitudes contains less oxygen than at the
       surface of the earth are based on the law of partial pressures
       (Chapter I.) According to this law, the equilibrium of the
       oxygen in the strata of the atmosphere is not dependent on
       the equilibrium of the nitrogen, and the variation in the
       densities of both gases with the height is determined by the
       pressure of each gas separately. Details of the calculations
       and considerations here involved are contained in my work _On
       Barometric Levellings_, 1876, p. 48.

       On the basis of the law of partial pressure and of hypsometrical
       formulæ, expressing the laws of the variation of pressures at
       different altitudes, the conclusion may be deduced that at the
       upper strata of the atmosphere the proportion of the nitrogen
       with respect to the oxygen increases, but the increase will not
       exceed a fraction per cent., even at altitudes of four and a half
       to six miles, the greatest height within the reach of men either
       by climbing mountains or by means of balloons. This conclusion is
       confirmed by the analyses of air collected by Welch in England
       during his aëronautic ascents.

_The analysis of air_ is effected by converting the oxygen into a
non-gaseous compound, so as to separate it from the air. The original
volume of the air is first measured, and then the volume of the
remaining nitrogen. The quantity of oxygen is calculated either from
the difference between these volumes or by the weight of the oxygen
compound formed. All the volumetric measurements have to be corrected
for pressure, temperature, and moisture (Chapters I. and II.) The medium
employed for converting the oxygen into a non-gaseous substance should
enable it to be taken up from the nitrogen to the very end without
evolving any gaseous substance. So, for instance,[19] a mixture of
pyrogallol, C_{6}H_{6}O_{3}, with a solution of a caustic alkali absorbs
oxygen with great ease at the ordinary temperature (the solution turns
black), but it is unsuited for accurate analysis because it requires
an aqueous solution of an alkali, and it alters the composition of
the air by acting on it as a solvent.[20] However, for approximate
determinations this simple method gives results which are entirely
satisfactory.

  [19] The complete absorption of the oxygen may be attained by
       introducing moist phosphorus into a definite volume of air; the
       occurrence of this is recognised by the fact of the phosphorus
       becoming non-luminous in the dark. The amount of oxygen may be
       determined by measuring the volume of nitrogen remaining. This
       method however cannot give accurate results, owing to a portion
       of the air being dissolved in the water, to the combination
       of some of the nitrogen with oxygen and to the necessity of
       introducing and withdrawing the phosphorus, which cannot be
       accomplished without introducing bubbles of air.

  [20] For rapid and approximate analyses (technical and hygienic), such
       a mixture is very suitable for determining the amount of oxygen
       in mixtures of gases from which the substances absorbed by
       alkalis have first been removed. According to certain observers,
       this mixture evolves a certain (small) quantity of carbonic oxide
       after absorbing oxygen.

The determinations in a eudiometer (Chapter III.) give more exact
results, if all the necessary corrections for changes of pressure,
temperature, and moisture be taken into account. This determination
is carried out essentially as follows:--A certain amount of air is
introduced into the eudiometer, and its volume is determined. About
an equal volume of dry hydrogen is then passed into the eudiometer,
and the volume again determined. The mixture is then exploded, in the
way described for the determination of the composition of water. The
remaining volume of the gaseous mixture is again measured; it will be
less than the second of the previously measured volumes. Out of three
volumes which have disappeared, one belonged to the oxygen and two to
the hydrogen, consequently one-third of the loss of volume indicates the
amount of oxygen contained in the air.[21]

  [21] Details of eudiometrical analysis must, as was pointed out in
       Chap. III., Note 32, be looked for in works on analytical
       chemistry. The same remark applies to the other analytical
       methods mentioned in this work. They are only described for the
       purpose of showing the diversity of the methods of chemical
       research.

The most complete method for the analysis of air, and one which is
accompanied by the least amount of error, consists in the direct
weighing, as far as is possible, of the oxygen, nitrogen, water,
and carbonic anhydride contained in it. For this purpose the air is
first passed through an apparatus for retaining the moisture and
carbonic anhydride (which will be considered presently), and is then
led through a tube which contains shavings of metallic copper and has
been previously weighed. A long layer of such copper heated to redness
absorbs all the oxygen from the air, and leaves pure nitrogen, whose
weight must be determined. This is done by collecting it in a weighed
and exhausted globe, while the amount by weight of oxygen is shown by
the increase in weight of the tube with the copper after the experiment.

[Illustration: FIG. 38.--Dumas and Boussingault's apparatus for the
analysis of air by weight. The globe B contains 10-15 litres. The air is
first pumped out of it, and it is weighed empty. The tube T connected
with it is filled with copper, and is weighed empty of air. It is heated
in a charcoal furnace. When the copper has become red-hot, the stopcock
_r_ (near R) is slightly opened, and the air passes through the vessels
L, containing a solution of potash, _f_, containing solutions and pieces
of caustic potash, which remove the carbonic anhydride from the air,
and then through _o_ and _t_, containing sulphuric acid (which has
been previously boiled to expel dissolved air) and pumice-stone, which
removes the moisture from the air. The pure air then gives up its oxygen
to the copper in T. When the air passes into T the stopcock R of the
globe B is opened, and it becomes filled with nitrogen. When the air
ceases to flow in, the stopcocks are closed, and the globe B and tube T
weighed. The nitrogen is then pumped out of the tube and it is weighed
again. The increase in weight of the tube shows the amount of oxygen,
and the difference of the second and third weighings of the tube, with
the increase in weight of the globe, gives the weight of the nitrogen.]

Air free from moisture and carbonic anhydride[22] contains 20·95 to
20·88[23] parts by volume of oxygen; the mean amount of oxygen will
therefore be 20·92 ± 0·05 per cent. Taking the density of air = 1 and of
oxygen = 1·105 and nitrogen 0·972 the composition of air by weight will
be 23·12 per cent. of oxygen and 76·88 per cent. of nitrogen.[24]

  [22] Air free from carbonic anhydride indicates after explosion the
       presence of a small quantity of carbonic anhydride, as De
       Saussure remarked, and air free from moisture, after being
       passed over red-hot copper oxide, appears invariably to contain
       a small quantity of water, as Boussingault has observed. These
       observations lead to the assumption that air always contains a
       certain quantity of gaseous hydrocarbons, like marsh gas, which,
       as we shall afterwards learn, is evolved from the earth, marshes,
       &c. Its amount, however, does not exceed a few hundredths per
       cent.

  [23] The analyses of air are accompanied by errors, and there are
       variations of composition attaining hundredths per cent.; the
       average normal composition of air is therefore only correct to
       the first decimal place.

  [24] These figures express the mean composition of air from an average
       of the most accurate determinations; they are accurate within
       ±0·05 p.c.

[Illustration: FIG. 39.--Apparatus for the absorption and washing
of gases, known as Liebig's bulbs. The gas enters _m_, presses on
the absorptive liquid, and passes from m into _b_, _c_, _d_, and _e_
consecutively, and escapes through _f_.]

[Illustration: FIG. 40.--Geisler's potash bulbs. The gas enters at
_a_, and passes through a solution of potash in the lower bulbs, where
the carbonic anhydride is absorbed, and the gas escapes from _b_. The
lower bulbs are arranged in a triangle, so that the apparatus can stand
without support.]

The possibility of the composition of air being altered by the mere
action of a solvent very clearly shows that the component parts of air
are in a state of mixture, in which any gases may occur; they do not
in this case form a definite compound, although the composition of the
atmosphere does appear constant under ordinary conditions. The fact that
its composition varies under different conditions confirms the truth
of this conclusion, and therefore the constancy of the composition of
air must not be considered as in any way dependent on the nature of the
gases entering into its composition, but only as proceeding from cosmic
phenomena co-operating towards this constancy. It must be admitted,
therefore, that the processes evolving oxygen, and chiefly the processes
of the respiration of plants, are of equal force with those processes
which absorb oxygen over the entire surface of the earth.[25]

  [25] In Chapter III., Note 4, an approximate calculation is made for
       the determination of the balance of oxygen in the entire
       atmosphere; it may therefore he supposed that the composition of
       air will vary from time to time, the relation between vegetation
       and the oxygen absorbing processes changes; but as the atmosphere
       of the earth can hardly have a definite limit and we have
       already seen (Chapter IV., Note 33) that there are observations
       confirming this, it follows that our atmosphere should vary in
       its component parts with the entire heavenly space, and therefore
       it must he supposed that any variation in the composition by
       weight of the air can only take place exceedingly slowly, and in
       a manner imperceptible by experiment.

[Illustration: FIG. 41.--Tube for the absorption of carbonic acid. A
plug of cotton wool is placed in the bulb to prevent the powder of
soda-lime being carried off by the gas. The tube contains soda-lime and
chloride of calcium.]

Air always contains more or less moisture[26] and _carbonic anhydride_
produced by the respiration of animals and the combustion of carbon and
carboniferous compounds. The latter shows the properties of an acid
anhydride. In order to determine the amount of carbonic anhydride in
air, substances are employed which absorb it--namely, alkalis either in
solution or solid. A solution of caustic potash, KHO, is poured into
light glass vessels, through which the air is passed, and the amount
of carbonic anhydride is determined by the increase in weight of the
vessel. But it is best to take a solid porous alkaline mass such as
soda-lime.[27] With a slow current of air a layer of soda-lime 20 cm.
in length is sufficient to completely deprive 1 cubic metre of air of
the carbonic anhydride it contains. A series of tubes containing calcium
chloride for absorbing the moisture[28] is placed before the apparatus
for the absorption of the carbonic anhydride, and a measured mass of
air is passed through the whole apparatus by means of an aspirator. In
this manner the determination of the moisture is combined with the
absorption of the carbonic anhydride. The arrangement shown in fig. 38
is such a combination.

  [26] The amount of moisture contained in the air is considered in
       greater detail in the study of physics and meteorology and the
       subject has been mentioned above, in Chapter I., Note 1, where
       the methods of absorbing moisture from gases were pointed out.

  [27] Soda-lime is prepared in the following manner:--Unslaked lime
       is finely powdered and mixed with a slightly warmed and very
       strong solution of caustic soda. The mixing should be done in
       an iron dish, and the materials should be well stirred together
       until the lime begins to slake. When the mass becomes hot, it
       boils, swells up, and solidifies, forming a porous mass very rich
       in alkali and capable of rapidly absorbing carbonic anhydride. A
       lump of caustic soda or potash presents a much smaller surface
       for absorption and therefore acts much less rapidly. It is
       necessary to place an apparatus for absorbing water after the
       apparatus for absorbing the carbonic anhydride, because the
       alkali in absorbing the latter gives off water.

  [28] It is evident that the calcium chloride employed for absorbing the
       water should be free from lime or other alkalis in order that it
       may not retain carbonic anhydride. Such calcium chloride may be
       prepared in the following manner: A perfectly neutral solution of
       calcium chloride is prepared from lime and hydrochloric acid; it
       is then carefully evaporated first on a water-bath and then on a
       sand-bath. When the solution attains a certain strength a scum is
       formed, which solidifies at the surface. This scum is collected,
       and will be found to be free from caustic alkalis. It is
       necessary in any case to test it before use, as otherwise a large
       error may be introduced into the results, owing to the presence
       of free alkali (lime). It is best to pass carbonic anhydride
       through the tube containing the calcium chloride for some time
       before the experiment, in order to saturate any free alkali that
       may remain from the decomposition of a portion of the calcium
       chloride by water, CaCl_{2} + 2H_{2}O = CaOH_{2}O + 2HCl.

The amount of carbonic anhydride[29] in free air is incomparably more
constant than the amount of moisture. The average amount in 100 volumes
of dry air is approximately 0·03 volume--that is, 10,000 volumes of air
contain about three volumes of carbonic anhydride, most frequently about
2·95 volumes. As the specific gravity of carbonic anhydride referred to
air = 1·52, it follows that 100 parts by weight of air contain 0·045
part by weight of carbonic anhydride. This quantity varies according
to the time of year (more in winter), the altitude above the level of
the sea (less at high altitudes), the proximity to forests and fields
(less) or cities (greater), &c. But the variation is small and rarely
exceeds the limits of 2-1/2 to 4 ten-thousandths by volume.[30] As there
are many natural local influences which either increase the amount of
carbonic anhydride in the air (respiration, combustion, decomposition,
volcanic eruptions, &c.), or diminish it (absorption by plants and
water), the reason of the great constancy in the amount of this gas in
the air must be looked for, in the first place, in the fact that the
wind mixes the air of various localities together, and, in the second
place, in the fact that the waters of the ocean, holding carbonic acid
in solution,[31] form an immense reservoir for regulating the amount of
this gas in the atmosphere. Immediately the partial pressure of the
carbonic anhydride in the air decreases, the water evolves it, and when
the partial pressure increases, it absorbs it, and thus nature supplies
the conditions for a natural state of moving equilibrium in this as in
so many other instances.[32]

  [29] Recourse is had to special methods when the determination only
       takes note of the carbonic anhydride of the air. For instance,
       it is absorbed by an alkali which does not contain carbonates
       (by a solution of baryta or caustic soda mixed with baryta),
       and then the carbonic anhydride is expelled by an excess of an
       acid, and its amount determined by the volume given off. A rapid
       method of determining CO_{2} (for hygienic purposes) is given by
       the fall of tension produced by the introduction of an alkali
       (the air having been either brought to dryness or saturated
       with moisture). Dr. Schidloffsky's apparatus is based upon this
       principle. The question as to the amount of carbonic anhydride
       present in the air has been submitted to many voluminous and
       exact researches, especially those of Reiset, Schloesing, Müntz,
       and Aubin, who showed that the amount is not subject to such
       variations as at first announced on the basis of incomplete and
       insufficiently accurate determinations.

  [30] It is a different case in enclosed spaces, in dwellings, cellars,
       wells, caves, and mines, where the renewal of air is impeded.
       Under these circumstances large quantities of carbonic anhydride
       may accumulate. In cities, where there are many conditions for
       the evolution of carbonic anhydride (respiration, decomposition,
       combustion), its amount is greater than in free air, yet even
       in still weather the difference does not often exceed one
       ten-thousandth (that is, rarely attains 4 instead of 2·9 vols. in
       10000 vols. of air).

  [31] In the sea as well as in fresh water, carbonic acid occurs in
       two forms, directly dissolved in the water, and combined with
       lime as calcium bicarbonate (hard waters sometimes contain very
       much carbonic acid in this form). The tension of the carbonic
       anhydride in the first form varies with the temperature, and
       its amount with the partial pressure, and that in the form of
       acid salts is under the same conditions, for direct experiments
       have shown a similar dependence in this case, although the
       quantitative relations are different in the two cases.

  [32] In studying the phenomena of nature the conclusion is arrived at
       that the universally reigning state of mobile equilibrium forms
       the chief reason for that harmonious order which impresses all
       observers. It not unfrequently happens that we do not see the
       causes regulating the order and harmony; in the particular
       instance of carbonic anhydride, it is a striking circumstance
       that in the first instance a search was made for an harmonious
       and strict uniformity, and in incidental (insufficiently
       accurate and fragmentary) observations conditions were even
       found for concluding it to be absent. When, later, the rule of
       this uniformity was confirmed, then the causes regulating such
       order were also discovered. The researches of Schloesing were of
       this character. Deville's idea of the dissociation of the acid
       carbonates of sea-water is suggested in them. In many other cases
       also, a correct interpretation can only follow from a detailed
       investigation.

Besides nitrogen, oxygen, moisture, and carbonic acid, all the other
substances occurring in air are found in infinitesimally small
quantities by weight, and therefore the _weight of a cubic measure of
air_ depends, to a sensible degree, on the above-named components alone.
We have already mentioned that at 0° and 760 mm. pressure the weight
of a cubic litre of air is 1·293 gram. This weight varies with the
acceleration of gravity, _g_, so that if _g_ be expressed in metres the
weight of a litre of air, _e_ = _g_ × 0·131844 gram. For St. Petersburg
_g_ is about 9·8188, and therefore _e_ is about 1·2946,[33] the air
being understood to be dry and free from carbonic anhydride. Taking
the amount of the latter as 0·03 per 100 volumes, we obtain a greater
weight; for example, for St. Petersburg _e_ = 1·2948 instead of 1·2946
gram. The weight of one litre of moist air in which the tension[34] of
the aqueous vapour (partial pressure) = _f_ mm., at a pressure (total)
of air of H millimetres, at a temperature _t_, will be (_i.e._, if
at 0° and 760 mm. the weight of dry air = _e_) equal to _e_/(1 +
0·00367_t_) × (H - 0·38_f_)/760. For instance, if H = 730 mm., _t_ = 20°,
and _f_ = 10 mm. (the moisture is then slightly below 60 p.c.), the
weight of a litre of air at St. Petersburg = 1·1527 gram.[35]

  [33] The difference of the weight of a litre of dry air (free from
       carbonic anhydride) at 0° and 760 mm., at different longitudes
       and altitudes, depends on the fact that the force of gravity
       varies under these conditions, and with it the pressure of the
       barometrical column also varies. This is treated in detail
       in my works _On the Elasticity of Gases_ and _On Barometric
       Levellings_, and 'The Publications of the Weights and Measures
       Department' (_Journal of the Russian Physico-Chemical Society_,
       1894).

       In reality the weight is not measured in absolute units of weight
       (in pressure--refer to works on mechanics and physics), but in
       relative units (grams, scale weights) whose mass is invariable,
       and therefore the variation of the weight of the weights itself
       with the change of gravity must not be here taken into account,
       for we are here dealing with weights proportional to masses,
       since with a change of locality the weight of the weights varies
       as the weight of a given volume of air does. In other words: the
       mass of a substance always remains constant, but the pressure
       produced by it varies with the acceleration of gravity: the gram,
       pound, and other units of weight are really units of mass.

  [34] The tension of the aqueous vapour in the air is determined by
       hygrometers and other similar methods. It may also be determined
       by analysis (_see_ Chapter I., Note 1).

  [35] For rapid calculation the weight of a litre of air (in a room) in
       St. Petersburg, may under these conditions (H, _t_, and _f_)
       be obtained by the formula _e_ = 1·20671 + 0·0016 [H_{1}-755 +
       2·6(18°-_t_°)] where H_{1} = H-0·38_f_. In determining the weight
       of small and heavy objects (crucibles, &c. in analysis, and in
       determining the specific gravities of liquids, &c.) _a correction
       may be introduced for the loss of weight_ in the air of the
       room, by taking the weight of a litre of air displaced as 1·2
       gram, and consequently 0·0012 gram for every cubic centimetre.
       But if gases or, in general, large vessels are weighed, and the
       weighings require to be accurate, it is necessary to take into
       account all the data for the determination of the density of the
       air (_t_, H, and _f_), because sensitive balances can determine
       the possible variations of the weight of air, as in the case of a
       litre the weight of air varies in centigrams, even at a constant
       temperature, with variations of H and _f_. Some time ago (1859) I
       proposed the following method and applied it for this purpose. A
       large light and closed vessel is taken, and its volume and weight
       in a vacuum are accurately determined, and verified from time
       to time. On weighing it we obtain the weight in air of a given
       density, and by subtracting this weight from its absolute weight
       and dividing by its volume we obtain the density of the air.

The presence of ammonia, a compound of nitrogen and hydrogen, in the
air, is indicated by the fact that all acids exposed to the air absorb
ammonia from it after a time. De Saussure observed that aluminium
sulphate is converted by air into a double sulphate of ammonium and
aluminium, or the so-called ammonia alum. Quantitative determinations
have shown that the amount of ammonia[36] contained in air varies at
different periods. However, it may be accepted that 100 cubic metres of
air do not contain less than 1 or more than 5 milligrams of ammonia.
It is remarkable that mountain air contains more ammonia than the air
of valleys. The air in those places where animal substances undergoing
change are accumulated, and especially that of stables, generally
contains a much greater quantity of this gas. This is the reason of the
peculiar pungent smell noticed in such places. Moreover ammonia, as we
shall learn in the following chapter, combines with acids, and should
therefore be found in air in the form of such combinations, since air
contains carbonic and nitric acids.

  [36] Schloesing studied the equilibrium of the ammonia of the
       atmosphere and of the rivers, seas, &c., and showed that the
       amount of the gas is interchangeable between them. The ratio
       between the amount of ammonia in a cubic metre of air and a litre
       of water at 0° = 0·004, at 10° = 0·010, at 25° = 0·040 to 1, and
       therefore in nature there is a state of equilibrium in the amount
       of ammonia in the atmosphere and waters.

The presence of nitric acid in air is proved without doubt by the fact
that rain-water contains an appreciable amount of it.

Further (as already mentioned in Chapter IV.), air contains ozone
and hydrogen peroxide and nitrous acid (and its ammonia salt), _i.e._
substances having a direct oxidising action (for instance, upon iodized
starch-paper), but they are present in very small quantities.[37]

  [37] Whilst formed in the air these oxidising substances (N_{2}O_{3},
       ozone and hydrogen peroxide) at the same time rapidly disappear
       from it by oxidising those substances which are capable of
       being oxidised. Owing to this instability their amounts vary
       considerably, and, as would be expected, they are met with to an
       appreciable amount in pure air, whilst their amount decreases
       to zero in the air of cities, and especially in dwellings where
       there is a maximum of substances capable of oxidisation and a
       minimum of conditions for the formation of such bodies. There
       is a causal connection between the amount of these substances
       present in the air and its purity--that is, the amount of foreign
       residues of organic origin liable to oxidation present in the
       air. Where there is much of such residues their amount must be
       small. When they are present the amount of organic substances
       must be small, as otherwise they would be destroyed. For this
       reason efforts have been made to apply ozone for purifying the
       air by evolving it by artificial means in the atmosphere; for
       instance, by passing a series of electrical sparks through
       the ventilating pipes conveying air into a building. Air thus
       ozonised destroys by oxidation--that is, brings about the
       combustion of--the organic residues present in the air, and
       thus will serve for purifying it. For these reasons the air of
       cities contains less ozone and such like oxidising agents than
       country air. This forms the distinguishing feature of country
       air. However, animal life cannot exist in air containing a
       comparatively large amount of ozone.

Besides substances in a gaseous or vaporous state,[38] there is always
found a more or less considerable quantity of substances which are
not known in a state of vapour. These substances are present in the
air as _dust_. If a linen surface, moistened with an acid, be placed
in perfectly pure air, then the washings are found to contain sodium,
calcium, iron, and potassium.[39] Linen moistened with an alkali absorbs
carbonic, sulphuric, phosphoric, and hydrochloric acids. Further, the
presence of organic substances in air has been proved by a similar
experiment. If a glass globe be filled with ice and placed in a room
where are a number of people, then the presence of organic substances,
like albuminous substances, may be proved in the water which condenses
on the surface of the globe. It may be that the miasmas causing
infection in marshy localities, hospitals, and in certain epidemic
illnesses proceed from the presence of such substances in the air (and
especially in water, which contains many micro-organisms), as well as
from the presence of germs of lower organisms in the air as a minute
dust. Pasteur proved the existence of such germs in the air by the
following experiment:--He placed gun-cotton (pyroxylin), which has the
appearance of ordinary cotton, in a glass tube. Gun-cotton is soluble
in a mixture of ether and alcohol, forming the so-called collodion.
A current of air was passed through the tube for a long period of
time, and the gun-cotton was then dissolved in a mixture of ether and
alcohol. An insoluble residue was thus obtained which actually contained
the germs of organisms, as was shown by microscopical observations,
and by their capacity to develop into organisms (mould, &c.) under
favourable conditions. The presence of these germs determines the
property of air of bringing about the processes of putrefaction and
fermentation--that is the fundamental alteration of organic substances,
which is accompanied by an entire change in their properties. The
appearance of lower organisms, both vegetable and animal, is always to
be remarked in these processes. Thus, for instance, in the process of
fermentation, when, for example, wine is procured from the sweet juice
of grapes, a sediment separates out which is known under the name of
lees, and contains peculiar yeast organisms. Germs are required before
these organisms can appear.[40] They are floating in the air, and fall
into the sweet fermentable liquid from it. Finding themselves under
favourable conditions, the germs develop into organisms; they are
nourished at the expense of the organic substance, and during growth
change and destroy it, and bring about fermentation and putrefaction.
This is why, for instance, the juice of the grape when contained in the
skin of the fruit, which allows access of the air but is impenetrable to
the germs, does not ferment, does not alter so long as the skin remains
intact. This is also the reason why animal substances when kept from the
access of air may be preserved for a great length of time. Preserved
foods for long sea voyages are kept in this way.[41] Hence it is
evident that however infinitesimal the quantity of germs carried in the
atmosphere may be, still they have an immense significance in nature.[42]

  [38] Amongst them we may mention iodine and alcohol, C_{2}H_{6}O, which
       Müntz found to be always present in air, the soil, and water,
       although in minute traces only.

  [39] A portion of the atmospheric dust is of cosmic origin; this is
       undoubtedly proved by the fact of its containing metallic iron as
       do meteorites. Nordenskiöld found iron in the dust covering snow,
       and Tissandier in every kind of air, although naturally in very
       small quantities.

  [40] The idea of the spontaneous growth of organisms in a suitable
       medium, although still upheld by many, has since the work of
       Pasteur and his followers (and to a certain extent of his
       predecessors) been discarded, because it has been proved how,
       when, and whence (from the air, water, &c.) the germs appear;
       that fermentation as well as infectious diseases cannot take
       place without them; and chiefly because it has been shown that
       any change accompanied by the development of the organisms
       introduced may be brought about at will by the introduction of
       the germs into a suitable medium.

  [41] In further confirmation of the fact that putrefaction and
       fermentation depend on germs carried in the air, we may cite the
       circumstance that poisonous substances destroying the life of
       organisms stop or hinder the appearance of the above processes.
       Air which has been heated to redness or passed through sulphuric
       acid no longer contains the germs of organisms, and loses the
       faculty of producing fermentation and putrefaction.

  [42] Their presence in the air is naturally due to the diffusion of
       germs into the atmosphere, and owing to their microscopical
       dimensions, they, as it were, hang in the air in virtue of their
       large surfaces compared to their weight. In Paris the amount of
       dust suspended in the air equals from 6 (after rain) to 23 grams
       per 1,000 c.m. of air.

Thus we see that air contains a great variety of substances. The
nitrogen, which is found in it in the largest quantity, has the least
influence on those processes which are accomplished by the action
of air. The oxygen, which is met with in a lesser quantity than the
nitrogen, on the contrary takes a very important part in a number of
reactions; it supports combustion and respiration, it brings about
decomposition and every process of slow oxidation. The part played by
the moisture of air is well known. The carbonic anhydride, which is met
with in still smaller quantities, has an immense significance in nature,
inasmuch as it serves for the nourishment of plants. The importance of
the ammonia and nitric acid is very great, because they are the sources
of the nitrogenous substances comprising an indispensable element in all
living organisms. And, lastly, the infinitesimal quantity of germs also
have a great significance in a number of processes. Thus it is not the
quantitative but the qualitative relations of the component parts of the
atmosphere which determine its importance in nature.[43]

  [43] We see similar cases everywhere. For example, the predominating
       mass of sand and clay in the soil takes hardly any chemical part
       in the economy of the soil in respect to the nourishment of
       plants. The plants by their roots search for substances which
       are diffused in comparatively small quantities in the soil. If a
       large quantity of these nourishing substances are removed, then
       the plants will not develop in the soil, just as animals die in
       oxygen.

Air, being a mixture of various substances, may suffer considerable
_changes_ in consequence of incidental circumstances. It is particularly
necessary to remark those changes in the composition of air which take
place in dwellings and in various localities where human beings have
to remain during a lengthy period of time. The respiration of human
beings and animals alters the air.[44] A similar deterioration of air
is produced by the influence of decomposing organic substances, and
especially of substances burning in it.[45] Hence it is necessary
to have regard to the purification of the air of dwellings. The
renewal of air, the replacing of respired by fresh air, is termed
'ventilation,'[46] and the removal of foreign and injurious admixtures
from the air is called 'disinfection.'[47] The accumulation of all kinds
of impurities in the air of dwellings and cities is the reason why the
air of mountains, forests, seas, and non-marshy localities, covered with
vegetation or snow, is distinguished for its freshness, and, in all
respects, beneficial action.

  [44] A man in breathing burns about 10 grams of carbon per hour--that
       is, he produces about 880 grams, or (as 1 cub.m. of carbonic
       anhydride weighs about 2,000 grams) about 5/12 c.m. of carbonic
       anhydride. The air coming from the lungs contains 4 p.c. of
       carbonic anhydride by volume. The exhaled air acts as a direct
       poison, owing to this gas and to other impurities.

  [45] For this reason candles, lamps, and gas change the composition
       of air almost in the same way as respiration. In the burning of 1
       kilogram of stearin candles, 50 cubic metres of air are changed
       as by respiration--that is, 4 p.c. of carbonic anhydride will
       be formed in this volume of air. The respiration of animals and
       exhalations from their skins, and especially from the intestines
       and the excrements and the transformations taking place in them,
       contaminate the air to a still greater extent, because they
       introduce other volatile substances besides carbonic anhydride
       into the air. At the same time that carbonic anhydride is formed
       the amount of oxygen in the air decreases, and there is noticed
       the appearance of miasmata which occur in but small quantity,
       but which are noticeable in passing from fresh air into a
       confined space full of such adulterated air. The researches of
       Schmidt and Leblanc and others show that even with 20·6 p.c. of
       oxygen (instead of 20·9 p.c.), when the diminution is due to
       respiration, air becomes noticeably less fit for respiration,
       and that the heavy feeling experienced in such air increases
       with a lesser percentage of oxygen. It is difficult to remain
       for a few minutes in air containing 17·2 p.c. of oxygen. These
       observations were chiefly obtained by observations on the air of
       different mines, at different depths below the surface. The air
       of theatres and buildings full of people also proves to contain
       less oxygen; it was found on one occasion that at the end of a
       theatrical representation the air in the stalls contained 20·75
       p.c. of oxygen, whilst the air at the upper part of the theatre
       contained only 20·36 p.c. The amount of carbonic anhydride in
       the air may be taken as a measure of its purity (Pettenkofer).
       When it reaches 1 p.c. it is very difficult for human beings to
       remain long in such air, and it is necessary to set up a vigorous
       ventilation for the removal of the adulterated air. In order
       to keep the air in dwellings in a uniformly good state, it is
       necessary to introduce at least 10 cubic metres of fresh air per
       hour per person. We saw that a man exhales about five-twelfths
       of a cubic metre of carbonic anhydride per day. Accurate
       observations have shown that air containing one-tenth p.c. of
       exhaled carbonic anhydride (and consequently also a corresponding
       amount of the other substances evolved together with it) is not
       felt to be oppressive; and therefore the five-twelfth cubic
       metres of carbonic anhydride should be diluted with 420 cubic
       metres of fresh air if it be desired to keep not more than
       one-tenth p.c. (by volume) of carbonic anhydride in the air.
       Hence a man requires 420 cubic metres of air per day, or 18 cubic
       metres per hour. With the introduction of only 10 cubic metres of
       fresh air per person, the amount of carbonic anhydride may reach
       one-fifth p.c., and the air will not then be of the requisite
       freshness.

  [46] The _ventilation_ of inhabited buildings is most necessary, and
       is even indispensable in hospitals, schools, and similar
       buildings. In winter it is carried on by the so-called
       calorifiers or stoves heating the air before it enters. The best
       kind of calorifiers in this respect are those in which the fresh
       cold air is led through a series of pipes heated by the hot gases
       coming from a stove. In ventilation, particularly during winter,
       care is taken that the incoming air shall be moist, because
       in winter the amount of moisture in the air is very small.
       Ventilation, besides introducing fresh air into a dwelling-place,
       must also withdraw the air already spoilt by respiration and
       other causes--that is, it is necessary to construct channels for
       the escape of the bad air, besides those for the introduction of
       fresh air. In ordinary dwelling-places, where not many people
       are congregated, the ventilation is conducted by natural means,
       in the heating by fires, through crevices, windows, and various
       orifices in walls, doors, and windows. In mines, factories, and
       workrooms ventilation is of the greatest importance.

       Animal vitality may still continue for a period of several
       minutes in air containing up to 30 p.c. of carbonic anhydride, if
       the remaining 70 p.c. consist of ordinary air; but respiration
       ceases after a certain time, and death may even ensue. The
       flame of a candle is very easily extinguished in an atmosphere
       containing from 5 to 6 p.c. of carbonic anhydride, but animal
       vitality can be sustained in it for a somewhat long time,
       although the effect of such air is exceedingly painful even to
       the lower animals. There are mines in which a lighted candle
       easily goes out from the excess of carbonic anhydride, but in
       which the miners have to remain for a long time. The presence
       of 1 p.c. of carbonic oxide is deadly even to cold-blooded
       animals. The air in the galleries of a mine where blasting
       has taken place, is known to produce a state of insensibility
       resembling that produced by charcoal fumes. Deep wells and
       vaults not unfrequently contain similar substances, and their
       atmosphere often causes suffocation. The atmospheres of such
       places cannot be tested by lowering a lighted candle into it, as
       these poisonous gases would not extinguish the flame. This method
       only suffices to indicate the amount of carbonic anhydride. If a
       candle keeps alight, it signifies that there is less than 6 p.c.
       of this gas. In doubtful cases it is best to lower a dog or other
       animal into the air to be tested. If CO_{2} be very carefully
       added to air, the flame of a candle is not extinguished (although
       it becomes very much smaller) even when the gas amounts to 12
       p.c. of air. Researches made by F. Clowes (1894) show that the
       flames (in every case 3/4 in. long) of different combustible
       substances are extinguished by the gradual addition of different
       percentages of nitrogen and carbonic acid to the air; the
       percentage sufficient to extinguish the flame being as follows
       (the percentage of oxygen is given in parenthesis):

                                  p.c. CO_{2}   p.c.  N.
               Absolute alcohol    14  (18·1)    21 (16·6)
               Candle              14  (18·1)    22 (16·4)
               Hydrogen            58  ( 8·8)    70 ( 6·3)
               Coal gas            33  (14·1)    46 (11·3)
               Carbonic oxide      24  (16·0)    28 (15·1)
               Methane             10  (18·9)    17 (17·4)

       The flames of all solid and liquid substances is extinguished by
       almost the same percentage of CO_{2} or N_{2}, but the flames of
       different gases vary in this respect, and hydrogen continues to
       burn in mixtures which are far poorer in oxygen than those in
       which the flames of other combustible gases are extinguished;
       the flame of methane CH_{4} is the most easily extinguished. The
       percentage of nitrogen may be greater than that of CO_{2}. This,
       together with the fact that, under the above circumstances, the
       flame of a gas before going out becomes fainter and increases in
       size, seems to indicate that the chief reason for the extinction
       of the flame is the fall in its temperature.

  [47] Different so-called disinfectants purify the air, and prevent the
       injurious action of certain of its components by changing or
       destroying them. Disinfection is especially necessary in those
       places where a considerable amount of volatile substances
       are evolved into the air, and where organic substances are
       decomposed; for instance, in hospitals, closets, &c. The numerous
       disinfectants are of the most varied nature. They may be divided
       into oxidising, antiseptic, and absorbent substances. To the
       oxidising substances used for disinfection belong chlorine, and
       various substances evolving it, because chlorine in the presence
       of water oxidises the majority of organic substances, and this
       is why chlorine is used as a disinfectant for Siberian plagues.
       Further, to this class belong the permanganates of the alkalis
       and peroxide of hydrogen, as substances easily oxidising matters
       dissolved in water; these salts are not volatile like chlorine,
       and therefore act much more slowly, and in a much more limited
       sphere. Antiseptic substances are those which convert organic
       substances into such as are little prone to change, and prevent
       putrefaction and fermentation. They most probably kill the
       germs of organisms occurring in miasmata. The most important of
       these substances are creosote and phenol (carbolic acid), which
       occur in tar, and act in preserving smoked meat. Phenol is a
       substance little soluble in water, volatile, oily, and having the
       characteristic smell of smoked objects. Its action on animals in
       considerable quantities is injurious, but in small quantities,
       used in the form of a weak solution, it prevents the change of
       animal matter. The smell of privies, which depends on the change
       of excremental matter, may be easily removed by means of chlorine
       or phenol. Salicylic acid, thymol, common tar, and especially
       its solution in alkalis as proposed by Nensky, &c., are also
       substances having the same property. Absorbent substances are
       of no less importance, especially as preventatives, than the
       preceding two classes of disinfectants, inasmuch as they are
       innocuous. They are those substances which absorb the odoriferous
       gases and vapours emitted during putrefaction, which are chiefly
       ammonia, sulphuretted hydrogen, and other volatile compounds. To
       this class belong charcoal, certain salts of iron, gypsum, salts
       of magnesia, and similar substances, as well as peat, mould, and
       clay. Questions of disinfection and ventilation appertain to the
       most serious problems of common life and hygiene. These questions
       are so vast that we are here able only to give a short outline of
       their nature.




                               CHAPTER VI

           THE COMPOUNDS OF NITROGEN WITH HYDROGEN AND OXYGEN


[Illustration: FIG. 42.--The dry distillation of bones on a large scale.
The bones are heated in the vertical cylinders C (about 1-1/2 metre
high and 30 centimetres in diameter). The products of distillation pass
through the tubes T, into the condenser B, and receiver F. When the
distillation is completed the trap H is opened, and the burnt bones are
loaded into trucks V. The roof M is then opened, and the cylinders are
charged with a fresh quantity of bones. The ammonia water is preserved,
and goes to the preparation of ammoniacal salts, as described in the
following drawing.]

In the last chapter we saw that nitrogen does not directly combine
with hydrogen, but that a mixture of these gases in the presence of
hydrochloric acid gas, HCl, forms ammonium chloride, NH_{4}Cl, on the
passage of a series of electric sparks.[1] In ammonium chloride, HCl
is combined with NH_{3}, consequently N with H_{3} forms ammonia.[2]
Almost all the _nitrogenous substances of plants and animals_ evolve
ammonia when heated with an alkali. But even without the presence of
an alkali the majority of nitrogenous substances, when decomposed or
heated with a limited supply of air, evolve their nitrogen, if not
entirely, at all events partially, in the form of ammonia. When animal
substances such as skins, bones, flesh, hair, horns, &c., are heated
without access of air in iron retorts--they undergo what is termed
dry distillation. A portion of the resultant substances remains in the
retort and forms a carbonaceous residue, whilst the other portion, in
virtue of its volatility, escapes through the tube leading from the
retort. The vapours given off, on cooling, form a liquid which separates
into two layers; the one, which is oily, is composed of the so-called
animal oils (_oleum animale_): the other, an aqueous layer, contains
a solution of ammonia salts. If this solution be mixed with lime and
heated, the lime takes up the elements of carbonic acid from the ammonia
salts, and ammonia is evolved as a gas.[3] In ancient times ammonia
compounds were imported into Europe from Egypt, where they were prepared
from the soot obtained in the employment of camels' dung as fuel in
the locality of the temple of Jupiter Ammon (in Lybia), and therefore
the salt obtained was called 'sal-ammoniacale,' from which the name of
ammonia is derived. At the present time ammonia is obtained exclusively,
on a large scale, either from the products of the dry distillation of
animal or vegetable refuse, from urine, or from the ammoniacal liquors
collected in the destructive distillation of coal for the preparation
of coal gas. This ammoniacal liquor is placed in a retort with lime and
heated; the ammonia is then evolved together with steam.[4] In the arts,
only a small amount of ammonia is used in a free state--that is, in an
aqueous solution; the greater portion of it is converted into different
salts having technical uses, especially sal-ammoniac, NH_{4}Cl, and
ammonium sulphate, (NH_{4})_{2}SO_{4}. They are saline substances
which are formed because ammonia, NH_{3}, combines with all acids, HX,
forming ammonia salts, NH_{4}X. Sal-ammoniac, NH_{4}Cl, is a compound of
ammonia with hydrochloric acid. It is prepared by passing the vapours of
ammonia and water, evolved, as above described, from ammoniacal liquor,
into an aqueous solution of hydrochloric acid, and on evaporating the
solution sal-ammoniac is obtained in the form of soluble crystals[5]
resembling common salt in appearance and properties. Ammonia may be
very easily prepared _from_ this _sal-ammoniac_, NH_{4}Cl, as from any
other ammoniacal salt, by heating it with lime. Calcium hydroxide,
CaH_{2}O_{2}, as an alkali takes up the acid and sets free the
ammonia, forming calcium chloride, according to the equation 2NH_{4}Cl
+ CaH_{2}O_{2} = 2H_{2}O + CaCl_{2} + 2NH_{3}. In this reaction the
ammonia is evolved as a gas.[6]

[Illustration: FIG. 43.--Method of abstracting ammonia, on a large
scale, from ammonia water obtained at gas works by the dry distillation
of coal, or by the fermentation of urine, &c. This water is mixed with
lime and poured into the boiler C´´, and from thence into C´ and C
consecutively. The last boiler is heated directly over a furnace, and
hence no ammonia remains in solution after the liquid has been boiled
in it. The liquid is therefore then thrown away. The ammonia vapour and
steam pass from the boiler C, through the tube T, into the boiler C´,
and then into C´´, so that the solution in C´ becomes stronger than that
in C, and still stronger in C´´. The boilers are furnished with stirrers
A, A´, and A´´ to prevent the lime settling. From C´´ the ammonia and
steam pass through the tube T´´ into worm condensers surrounded with
cold water, thence into the Woulfe's bottle P, where the solution of
ammonia is collected, and finally the still uncondensed ammonia vapour
is led into the flat vessel R, containing acid which absorbs the last
traces of ammonia.]

  [1] The ammonia in the air, water, and soil proceeds from the
      decomposition of the nitrogenous substances of plants and animals,
      and also probably from the reduction of nitrates. Ammonia is
      always formed in the rusting of iron. Its formation in this case
      depends in all probability on the decomposition of water, and on
      the action of the hydrogen at the moment of its evolution on the
      nitric acid contained in the air (Cloez), or on the formation of
      ammonium nitrite, which takes place under many circumstances. The
      evolution of vapours of ammonia compounds is sometimes observed
      in the vicinity of volcanoes. At a red heat nitrogen combines
      directly with B Ca Mg, and with many other metals, and these
      compounds, when heated with a caustic alkali, or in the presence
      of water, give ammonia (_see_ Chapter XIV., Note 14, and Chapter
      XVII., Note 12). These are examples of the indirect combination of
      nitrogen with hydrogen.

  [2] If a silent discharge or a series of electric sparks be passed
      through ammonia gas, it is decomposed into nitrogen and
      hydrogen. This is a phenomenon of dissociation; therefore, a
      series of sparks do not totally decompose the ammonia, but
      leave a certain portion undecomposed. One volume of nitrogen
      and three volumes of hydrogen are obtained from two volumes of
      ammonia decomposed. Ramsay and Young (1884) investigated the
      decomposition of NH_{3} under the action of heat, and showed
      that at 500°, 1-1/2 p.c. is decomposed, at 600° about 18 p.c.,
      at 800° 65 p.c., but these results were hardly free from the
      influence of 'contact.' The _presence_ of free ammonia--that is,
      ammonia not combined with acids--in a gas or aqueous solution
      may be recognised by its characteristic smell. But many ammonia
      salts do not possess this smell. However, on the addition of
      an alkali (for instance, caustic lime, potash, or soda), they
      evolve ammonia gas, especially when heated. The presence of
      ammonia may be made visible by introducing a substance moistened
      with strong hydrochloric acid into its neighbourhood. A white
      cloud, or visible white vapour, then makes its appearance. This
      depends on the fact that both ammonia and hydrochloric acid are
      volatile, and on coming into contact with each other produce
      solid sal-ammoniac, NH_{4}Cl, which forms a cloud. This test is
      usually made by dipping a glass rod into hydrochloric acid, and
      holding it over the vessel from which the ammonia is evolved. With
      small amounts of ammonia this test is, however, untrustworthy,
      as the white vapour is scarcely observable. In this case it is
      best to take paper moistened with mercurous nitrate, HgNO_{3}.
      This paper turns black in the presence of ammonia, owing to the
      formation of a black compound of ammonia with mercurous oxide. The
      smallest traces of ammonia (for instance, in river water) may be
      detected by means of the so-called Nessler's reagent, containing
      a solution of mercuric chloride and potassium iodide, which forms
      a brown coloration or precipitate with the smallest quantities
      of ammonia. It will be useful here to give the thermochemical
      data (in thousands of units of heat, according to Thomsen), or
      the quantities of heat _evolved_ in the formation of ammonia and
      its compounds in quantities expressed by their formulæ. Thus, for
      instance, (N + H_{3}) 26·7 indicates that 14 grams of nitrogen
      in combining with 3 grams of hydrogen develop sufficient heat to
      raise the temperature of 26·7 kilograms of water 1°. (NH_{3} +
      nH_{2}O) 8·4 (heat of solution); (NH_{3},nH_{2}O + HCl,nH_{2}O)
      12·3; (N + H_{4} + Cl) 90·6; (NH_{3} + HCl) 41·9.

  [3] The same ammonia water is obtained, although in smaller quantities,
      in the dry distillation of plants and of coal, which consists
      of the remains of fossil plants. In all these cases the ammonia
      proceeds from the destruction of the complex nitrogenous
      substances occurring in plants and animals. The ammonia salts
      employed in the arts are prepared by this method.

  [4] The technical methods for the preparation of ammonia water, and for
      the extraction of ammonia from it, are to a certain extent
      explained in the figures accompanying the text.

  [5] Usually these crystals are sublimed by heating them in crucibles or
      pots, when the vapours of sal-ammoniac condense on the cold covers
      as a crust, in which form the salt comes into the market.

  [6] On a small scale ammonia may be prepared in a glass flask by mixing
      equal parts by weight of slaked lime and finely-powdered
      sal-ammoniac, the neck of the flask being connected with an
      arrangement for drying the gas obtained. In this instance
      neither calcium chloride nor sulphuric acid can be used for
      drying the gas, since both these substances absorb ammonia, and
      therefore solid caustic potash, which is capable of retaining
      the water, is employed. The gas-conducting tube leading from the
      desiccating apparatus is introduced into a mercury bath, if dry
      gaseous ammonia be required, because water cannot be employed
      in collecting ammonia gas. Ammonia was first obtained in this
      dry state by Priestley, and its composition was investigated by
      Berthollet at the end of the last century. Oxide of lead mixed
      with sal-ammoniac (Isambert) evolves ammonia with still greater
      ease than lime. The cause and process of the decomposition are
      almost the same, 2PbO + 2NH_{4}Cl = Pb_{2}OCl_{2} + H_{2}O +
      2NH_{3}. Lead oxychloride is (probably) formed.

It must be observed that all the complex nitrogenous substances of
plants, animals, and soils are decomposed when heated with an excess
of sulphuric acid, the whole of their nitrogen being converted into
ammonium sulphate, from which it may be liberated by treatment with
an excess of alkali. This reaction is so complete that it forms the
basis of Kjeldahl's method for estimating the amount of nitrogen in its
compounds.

Ammonia is a colourless gas, resembling those with which we are already
acquainted in its outward appearance, but clearly distinguishable from
any other gas by its very characteristic and pungent smell. It irritates
the eyes, and it is positively impossible to inhale it. Animals die in
it. Its density, referred to hydrogen, is 8·5; hence it is lighter than
air. It belongs to the class of gases which are easily liquefied.[7]
Faraday employed the following method for liquefying ammonia. Ammonia
when passed over dry silver chloride, AgCl, is absorbed by it to a
considerable extent, especially at low temperatures.[8] The solid
compound AgCl,3NH_{3} thus obtained is introduced into a bent tube
(fig. 45), whose open end c is then fused up. The compound is then
slightly heated at _a_, and the ammonia comes off, owing to the easy
dissociation of the compound. The other end of the tube is immersed in
a freezing mixture. The pressure of the gas coming off, combined with
the low temperature at one end of the tube, causes the ammonia evolved
to condense into a liquid, in which form it collects at the cold end of
the tube. If the heating be stopped, the silver chloride again absorbs
the ammonia. In this manner one tube may serve for repeated experiments.
Ammonia may also be liquefied by the ordinary methods--that is, by means
of pumping dry ammonia gas into a refrigerated space. Liquefied ammonia
is a colourless and very mobile liquid,[9] whose specific gravity at 0°
is 0·63 (E. Andréeff). At the temperature (about -70°) given by a mixture
of liquid carbonic anhydride and ether, liquid ammonia crystallises, and
in this form its odour is feeble, because at so low a temperature its
vapour tension is very inconsiderable. The boiling point (at a pressure
of 760 mm.) of liquid ammonia is about -32°. Hence this temperature may
be obtained at the ordinary pressure by the evaporation of liquefied
ammonia.

[Illustration: FIG. 45.--The liquefaction of ammonia in a thick bent
glass tube. A compound of chloride of silver and ammonia is placed in
the end _a_, and the end _c_ is then sealed up.]

  [7] [Illustration: FIG. 44.--Carré's apparatus. Described in text.]

      This is evident from the fact that its absolute boiling point lies
      at about +130° (Chapter II., Note 29). It may therefore be
      liquefied by pressure alone at the ordinary, and even at much
      higher temperatures. The latent heat of evaporation of 17 parts
      by weight of ammonia equals 4,400 units of heat, and hence
      liquid ammonia may be employed for the production of cold.
      Strong aqueous solutions of ammonia, which in parting with their
      ammonia act in a similar manner, are not unfrequently employed
      for this purpose. Suppose a saturated solution of ammonia to be
      contained in a closed vessel furnished with a receiver. If the
      ammoniacal solution be heated, the ammonia, with a small quantity
      of water, will pass off from the solution, and in accumulating
      in the apparatus will produce a considerable pressure, and will
      therefore liquefy in the cooler portions of the receiver. Hence
      liquid ammonia will be obtained in the receiver. The heating of
      the vessel containing the aqueous solution of ammonia is then
      stopped. After having been heated it contains only water, or a
      solution poor in ammonia. When once it begins to cool the ammonia
      vapours commence dissolving in it, the space becomes rarefied, and
      a rapid vaporisation of the liquefied ammonia left in the receiver
      takes place. In evaporating in the receiver it will cause the
      temperature in it to fall considerably, and will itself pass into
      the aqueous solution. In the end, the same ammoniacal solution as
      originally taken is re-obtained. Thus, in this case, on heating
      the vessel the pressure increases by itself, and on cooling it
      diminishes, so that here heat directly replaces mechanical work.
      This is the principle of the simplest forms of _Carré's ice-making
      machines_, shown in fig. 44. C is a vessel made of boiler plates
      into which the saturated solution of ammonia is poured; m is a
      tube conducting the ammonia vapour to the receiver A. All parts
      of the apparatus should be hermetically joined together, and
      should be able to withstand a pressure reaching ten atmospheres.
      The apparatus should be freed from air, which would otherwise
      hinder the liquefaction of the ammonia. The process is carried on
      as follows:--The apparatus is first so inclined that any liquid
      remaining in A may flow into C. The vessel C is then placed
      upon a stove F, and heated until the thermometer _t_ indicates
      a temperature of 130° C. During this time the ammonia has been
      expelled from C, and has liquefied in A. In order to facilitate
      the liquefaction, the receiver A should be immersed in a tank of
      water R (_see_ the left-hand drawing in fig. 44). After about
      half an hour, when it may be supposed that the ammonia has been
      expelled, the fire is removed from under C, and this is now
      immersed in the tank of water R. The apparatus is represented in
      this position in the right-hand drawing of fig. 44. The liquefied
      ammonia then evaporates, and passes over into the water in C. This
      causes the temperature of A to fall considerably. The substance
      to be refrigerated is placed in a vessel G, in the cylindrical
      space inside the receiver A. The refrigeration is also kept on for
      about half an hour, and with an apparatus of ordinary dimensions
      (containing about two litres of ammonia solution), five kilograms
      of ice are produced by the consumption of one kilogram of coal. In
      industrial works more complicated types of Carré's machines are
      employed.

  [8] Below 15° (according to Isambert), the compound AgCl,3NH_{3} is
      formed, and above 20° the compound 2AgCl,3NH_{3}. The tension of
      the ammonia evolved from the latter substance is equal to the
      atmospheric pressure at 68°, whilst for AgCl,3NH_{3} the pressures
      are equal at about 20°; consequently, at higher temperatures
      it is greater than the atmospheric pressure, whilst at lower
      temperatures the ammonia is absorbed and forms this compound.
      Consequently, all the phenomena of dissociation are here clearly
      to be observed. Joannis and Croisier (1894) investigated similar
      compounds with AgBr, AgI, AgCN and AgNO_{3}, and found that
      they all give definite compounds with NH_{3}, for instance
      AgBr,3NH_{3}, 2AgBr,3NH_{3} and AgBr,2NH_{3}; they are all
      colourless, solid substances which decompose under the atmospheric
      pressure at +3·5, +34° and +51°.

  [9] The liquefaction of ammonia may be accomplished without an increase
      of pressure, by means of refrigeration alone, in a carefully
      prepared mixture of ice and calcium chloride (because the absolute
      boiling point of NH_{3} is high, about +130°). It may even take
      place in the severe frosts of a Russian winter. The application of
      liquid ammonia as a motive power for engines forms a problem which
      has to a certain extent been solved by the French engineer Tellier.

Ammonia, containing, as it does, much hydrogen, is _capable of
combustion_; it does not, however, burn steadily, and sometimes not
at all, in ordinary atmospheric air. In pure oxygen it burns with a
greenish-yellow flame,[10] forming water, whilst the nitrogen set
free gives its oxygen compounds--that is, oxides of nitrogen. The
decomposition of ammonia into hydrogen and nitrogen not only takes
place at a red heat and under the action of electric sparks, but also
by means of many oxidising substances; for instance, by passing ammonia
through a tube containing red-hot copper oxide. The water thus formed
may be collected by substances absorbing it, and the quantity of
nitrogen may be measured in a gaseous form, and thus the composition
of ammonia determined. In this manner it is very easy to prove that
ammonia contains 3 parts by weight of hydrogen to 14 parts by weight of
nitrogen; and, by volume, 3 vols. of hydrogen and 1 vol. of nitrogen
form 2 vols. of ammonia.[11]

  [10] The combustion of ammonia in oxygen may be effected by the aid of
       platinum. A small quantity of an aqueous solution of ammonia,
       containing about 20 p.c. of the gas, is poured into a wide-necked
       beaker of about one litre capacity. A gas-conducting tube about
       10 mm. in diameter, and supplying oxygen, is immersed in the
       aqueous solution of ammonia. But before introducing the gas an
       incandescent platinum spiral is placed in the beaker; the ammonia
       in the presence of the platinum is oxidised and burns, whilst the
       platinum wire becomes still more incandescent. The solution of
       ammonia is heated, and oxygen passed through the solution. The
       oxygen, as it bubbles off from the ammonia solution, carries with
       it a part of the ammonia, and this mixture explodes on coming
       into contact with the incandescent platinum. This is followed
       by a certain cooling effect, owing to the combustion ceasing,
       but after a short interval this is renewed, so that one feeble
       explosion follows after another. During the period of oxidation
       without explosion, white vapours of ammonium nitrite and
       red-brown vapours of oxides of nitrogen make their appearance,
       while during the explosion there is complete combustion and
       consequently water and nitrogen are formed.

  [11] This may be verified by their densities. Nitrogen is 14 times
       denser than hydrogen, and ammonia is 8-1/2 times. If 3 volumes
       of hydrogen with 1 volume of nitrogen gave 4 volumes of ammonia,
       then these 4 volumes would weigh 17 times as much as 1 volume of
       hydrogen; consequently 1 volume of ammonia would be 4-1/4 times
       heavier than the same volume of hydrogen. But if these 4 volumes
       only give 2 volumes of ammonia, the latter will be 8-1/2 times as
       dense as hydrogen, which is found to be actually the case.

Ammonia is capable of combining with a number of substances, forming,
like water, substances of various degrees of stability. It is more
soluble than any of the gases yet described, both in water and in many
aqueous solutions. We have already seen, in the first chapter, that
one volume of water, at the ordinary temperature, dissolves about 700
vols. of ammonia gas. The great solubility of ammonia enables it to
be always kept ready for use in the form of an aqueous solution,[12]
which is commercially known as _spirits of hartshorn_. Ammonia water is
continually evolving ammoniacal vapour, and so has the characteristic
smell of ammonia itself. It is a very characteristic and important
fact that ammonia has an alkaline reaction, and colours litmus paper
blue, just like caustic potash or lime; it is therefore sometimes
called _caustic ammonia_ (volatile alkali). Acids may be saturated by
ammonia water or gas in exactly the same way as by any other alkali.
In this process _ammonia combines directly with acids_, and this forms
the most essential chemical reaction of this substance. If sulphuric,
nitric, acetic, or any other acid be brought into contact with ammonia
it absorbs it, and in so doing evolves a large amount of heat and forms
a compound having all the properties of a salt. Thus, for example,
sulphuric acid, H_{2}SO_{4}, in absorbing ammonia, forms (on evaporating
the solution) two salts, according to the relative quantities of ammonia
and acid. One salt is formed from NH_{3} + H_{2}SO_{4}, and consequently
has the composition NH_{5}SO_{4}, and the other is formed from 2NH_{3}
+ H_{2}SO_{4}, and its composition is therefore N_{2}H_{8}SO_{4}. The
former has an acid reaction and the latter a neutral reaction, and
they are called respectively acid ammonium sulphate (ammonium hydrogen
sulphate), and normal ammonium sulphate, or simply ammonium sulphate.
The same takes place in the action of all other acids; but certain of
them are able to form normal ammonium salts only, whilst others give
both acid and normal ammonium salts. This depends on the nature of
the acid and not on the ammonia, as we shall afterwards see. Ammonium
salts are very similar in appearance and in many of their properties
to metallic salts; for instance, sodium chloride, or table salt,
resembles sal-ammoniac, or ammonium chloride, not only in its outward
appearance but even in crystalline form, in its property of giving
precipitates with silver salts, in its solubility in water, and in
its evolving hydrochloric acid when heated with sulphuric acid--in a
word, a most perfect analogy is to be remarked in an entire series of
reactions. An analogy in composition is seen if sal-ammoniac, NH_{4}Cl,
be compared with table salt, NaCl; and the ammonium hydrogen sulphate,
NH_{4}HSO_{4}, with the sodium hydrogen sulphate, NaHSO_{4}; or ammonium
nitrate, NH_{4}NO_{3}, with sodium nitrate, NaNO_{3}.[13] It is seen,
on comparing the above compounds, that the part which sodium takes in
the sodium salts is played in ammonium salts by a group NH_{4}, which
is called _ammonium_. If table salt be called 'sodium chloride,' then
sal-ammoniac should be and is called 'ammonium chloride.'

  [12] Aqueous solutions of ammonia are lighter than water, and at 15°,
       taking water at 4° = 10,000, their specific gravity, as dependent
       on _p_, or the percentage amount (by weight) of ammonia, is given
       by the expression _s_ = 9,992-42·5_p_ + 0·21_p_^2; for instance,
       with 10 p.c. _s_ = 9,587. If _t_ represents the temperature
       between the limits of +10° and +20°, then the expression
       (15-_t_)(1·5 + 0·14_p_) must be added to the formula for the
       specific gravity. Solutions containing more than 24 p.c. have
       not been sufficiently investigated in respect to the variation
       of their specific gravity. It is, however, easy to obtain
       more concentrated solutions, and at 0° solutions approaching
       NH_{3},H_{2}O (48·6 p.c. NH_{3}) in their composition, and of
       sp. gr. 0·85, may be prepared. But such solutions give up the
       bulk of their ammonia at the ordinary temperature, so that more
       than 24 p.c. NH_{3} is rarely contained in solution. Ammoniacal
       solutions containing a considerable amount of ammonia give
       ice-like crystals which seem to contain ammonia at temperatures
       far below 0° (for instance, an 8 p.c. solution at -14°, the
       strongest solutions at -48°). The whole of the ammonia may be
       expelled from a solution by heating, even at a comparatively
       low temperature; hence on heating aqueous solutions containing
       ammonia a very strong solution of ammonia is obtained in the
       distillate. Alcohol, ether, and many other liquids are also
       capable of dissolving ammonia. Solutions of ammonia, when exposed
       to the atmosphere, give off a part of their ammonia in accordance
       with the laws of the solution of gases in liquids, which we have
       already considered. But the ammoniacal solutions at the same time
       absorb carbonic anhydride from the air, and ammonium carbonate
       remains in the solution.

       [Illustration: FIG. 46.--Apparatus for preparing solutions of
       ammonia.]

       Solutions of ammonia are required both for laboratory and factory
       operations, and have therefore to be frequently prepared. For
       this purpose the arrangement shown in fig. 46 is employed in the
       laboratory. In works the same arrangement is used, only on a
       larger scale (with earthenware or metallic vessels). The gas is
       prepared in the retort, from whence it is led into the two-necked
       globe A, and then through a series of Woulfe's bottles, B, C,
       D, E. The impurities spurting over collect in A, and the gas is
       dissolved in B, but the solution soon becomes saturated, and a
       purer (washed) ammonia passes over into the following vessels,
       in which only a pure solution is obtained. The bent funnel tube
       in the retort preserves the apparatus from the possibility both
       of the pressure of the gas evolved in it becoming too great
       (when the gas escapes through it into the air), and also from
       the pressure incidentally falling too low (for instance, owing
       to a cooling effect, or from the reaction stopping). If this
       takes place, the air passes into the retort, otherwise the liquid
       from B would be drawn into A. The safety tubes in each Woulfe's
       bottle, open at both ends, and immersed in the liquid, serve for
       the same purpose. Without them, in case of an accidental stoppage
       in the evolution of so soluble a gas as ammonia, the solution
       would be sucked from one vessel to another--for instance, from
       E into D, &c. In order to clearly see the necessity for _safety
       tubes_ in a gas apparatus, it must be remembered that the
       _gaseous pressure_ in the interior of the arrangement must exceed
       the atmospheric pressure by the height of the sum of the columns
       of liquid through which the gas has to pass.

  [13] The analogy between the ammonium and sodium salts might seem to
       be destroyed by the fact that the latter are formed from the
       alkali or oxide and an acid, with the separation of water, whilst
       the ammonium salts are directly formed from ammonia and an acid,
       without the separation of water; but the analogy is restored
       if we compare soda to ammonia water, and liken caustic soda to
       a compound of ammonia with water. Then the very preparation of
       ammonium salts from such a hydrate of ammonia will completely
       resemble the preparation of sodium salts from soda. We may cite
       as an example the action of hydrochloric acid on both substances.

           NaHO          +         HCl         =  H_{2}O  +     NaCl
       Sodium hydroxide     Hydrochloric acid      Water     Table salt

            NH_{4}HO      +         HCl         =  H_{2}0  +   NH_{4}Cl
       Ammonium hydroxide    Hydrochloric acid      Water    Sal-ammoniac

       Just as in soda the hydroxyl or aqueous radicle OH is replaced by
       chlorine, so it is in ammonia hydrate.

The hypothesis that ammoniacal salts correspond with a complex metal
ammonium bears the name of the _ammonium theory_. It was enunciated
by the famous Swedish chemist Berzelius after the proposition made by
Ampère. The analogy admitted between ammonium and metals is probable,
owing to the fact that mercury is able to form an amalgam with ammonium
similar to that which it forms with sodium or many other metals. The
only difference between ammonium amalgam and sodium amalgam consists in
the instability of the ammonium, which easily decomposes into ammonia
and hydrogen.[14] Ammonium amalgam may be prepared from sodium amalgam.
If the latter be shaken up with a strong solution of sal-ammoniac, the
mercury swells up violently and loses its mobility whilst preserving its
metallic appearance. In so doing, the mercury dissolves ammonium--that
is, the sodium in the mercury is replaced by the ammonium, and replaces
it in the sal-ammoniac, forming sodium chloride, NH_{4}Cl + HgNa =
NaCl + HgNH_{4}. Naturally, the formation of ammonium amalgam does not
entirely prove the existence of ammonium itself in a separate state; but
it shows the possibility of this substance existing, and its analogy
with the metals, because only metals dissolve in mercury.[15] Ammonium
amalgam crystallises in cubes, three times heavier than water; it is
only stable in the cold, and particularly at very low temperatures. It
begins to decompose at the ordinary temperature, evolving ammonia and
hydrogen in the proportion of two volumes of ammonia and one volume
of hydrogen, NH_{4} = NH_{3} + H. By the action of water, ammonium
amalgam gives hydrogen and ammonia water, just as sodium amalgam gives
hydrogen and sodium hydroxide; and therefore, in accordance with
the ammonium theory, ammonia water must be looked on as containing
ammonium hydroxide, NH_{4}OH,[16] just as an aqueous solution of sodium
hydroxide, contains NaOH. The ammonium hydroxide, like ammonium itself,
is an unstable substance, which easily dissociates, and can only exist
in a free state at low temperatures.[17] Ordinary solutions of ammonia
must be looked on as the products of the dissociation of this hydroxide,
inasmuch as NH_{4}OH = NH_{3} + H_{2}O.

  [14] Weyl (1864) by subjecting sodium to the action of ammonia at
       the ordinary temperature and under considerable pressures,
       obtained a liquid, which was subsequently investigated by Joannis
       (1889), who confirmed the results obtained by Weyl. At 0° and the
       atmospheric pressure the composition of this substance is Na +
       5·3NH_{3}. The removal (at 0°) of ammonia from the liquid gives
       a solid copper-red body having the composition NH_{3}Na. The
       determination of the molecular weight of this substance by the
       fall of the tension of liquid ammonia gave N_{2}H_{6}Na_{2}. It
       is, therefore, free ammonium in which one H is replaced by Na.
       The compound with potassium, obtained under the same conditions,
       proved to have an analogous composition. By the decomposition of
       NH_{3}Na at the ordinary temperature, Joannis (1891) obtained
       hydrogen and sodium-amide NH_{2}Na in small colourless crystals
       which were soluble in water. The addition of liquid ammonia to
       metallic sodium and a saturated solution of sodium chloride,
       gives NH_{2}Na_{2}Cl, and this substance is sal-ammoniac, in
       which H_{2} is replaced by Na_{2}.

       If pure oxygen be passed through a solution of these compounds
       in ammonia at a temperature of about -50°, it is seen that the
       gas is rapidly absorbed. The liquid gradually loses its dark
       red colour and becomes lighter, and when it has become quite
       colourless a gelatinous precipitate is thrown down. After the
       removal of the ammonia, this precipitate dissolves easily in
       water with a considerable evolution of heat, but without giving
       off any gaseous products. The composition of the sodium compound
       thus obtained is NH_{2}Na_{2}HO, which shows that it is a hydrate
       of bisodium-ammonium. Thus, although free ammonium has not been
       obtained, still a sodium substitution product of it is known
       which corresponds to it as a salt to a hydrate. Ammonium amalgam
       was originally obtained in exactly the same way as sodium amalgam
       (Davy); namely, a piece of sal-ammoniac was taken, and moistened
       with water (in order to render it a conductor of electricity). A
       cavity was made in it, into which mercury was poured, and it was
       laid on a sheet of platinum connected with the positive pole of a
       galvanic battery, while the negative pole was put into connection
       with the mercury. On passing a current the mercury increased
       considerably in volume, and became plastic, whilst preserving
       its metallic appearance, just as would be the case were the
       sal-ammoniac replaced by a lump of a sodium salt or of many other
       metals. In the analogous decomposition of common metallic salts,
       the metal contained in a given salt separates out at the negative
       pole, immersed in mercury, by which the metal is dissolved. A
       similar phenomenon is observed in the case of sal-ammoniac; the
       elements of ammonium, NH_{4}, in this case are also collected in
       the mercury, and are retained by it for a certain time.

  [15] We may mention, however, that under particular conditions hydrogen
       is also capable of forming an amalgam resembling the amalgam of
       ammonium. If an amalgam of zinc be shaken up with an aqueous
       solution of platinum chloride, without access of air, then a
       spongy mass is formed which easily decomposes, with the evolution
       of hydrogen.

  [16] We saw above that the solubility of ammonia in water at low
       temperatures attains to the molecular ratio NH_{3} + H_{2}O, in
       which these substances are contained in caustic ammonia, and
       perhaps it may be possible at exceedingly low temperatures to
       obtain ammonium hydroxide, NH_{4}HO, in a solid form. Regarding
       solutions as dissociated definite compounds, we should see a
       confirmation of this view in the property shown by ammonia of
       being extremely soluble in water, and in so doing of approaching
       to the limit NH_{4}HO.

  [17] In confirmation of the truth of this conclusion we may cite the
       remarkable fact that there exist, in a free state and as
       comparatively stable compounds, a series of alkaline hydroxides,
       NR_{4}HO, which are perfectly analogous to ammonium hydroxide,
       and present a striking resemblance to it and to sodium hydroxide,
       with the only difference that the hydrogen in NH_{4}HO is
       replaced by complex groups, R = CH_{3}, C_{2}H_{5}, &c., for
       instance N(CH_{3})_{4}HO. Details will be found in organic
       chemistry.

All ammoniacal salts _decompose at a red heat_ into ammonia and an acid,
which, on cooling in contact with each other, re-combine together. If
the acid be non-volatile, the ammoniacal salt, when heated, evolves the
ammonia, leaving the non-volatile acid behind; if the acid be volatile,
then, on heating, both the acid and ammonia volatilise together, and
on cooling re-combine into the salt which originally served for the
formation of their vapours.[18]

  [18] The fact that ammoniacal salts are decomposed when ignited, and not
       simply sublimed, may be proved by a direct experiment with
       sal-ammoniac, NH_{4}Cl, which in a state of vapour is decomposed
       into ammonia, NH_{3}, and hydrochloric acid, HCl, as will be
       explained in the following chapter. The readiness with which
       ammonium salts decompose is seen from the fact that a solution
       of ammonium oxalate is decomposed with the evolution of ammonia
       even at -1°. Dilute solutions of ammonium salts, when boiled give
       aqueous vapour having an alkaline reaction, owing to the presence
       of free ammonia given off from the salt.

Ammonia is not only capable of combining with acids, but also with many
salts, as was seen from its forming definite compounds, AgCl,3NH_{3}
and 2AgCl,3NH_{3}, with silver chloride. Just as ammonia is absorbed
by various oxygen salts of the metals, so also is it absorbed by the
chlorine, iodine, and bromine compounds of many metals, and in so doing
evolves heat. Certain of these compounds part with their ammonia even
when left exposed to the air, but others only do so at a red heat; many
give up their ammonia when dissolved, whilst others dissolve without
decomposition, and when evaporated separate from their solutions
unchanged. All these facts only indicate that ammoniacal, like aqueous,
compounds dissociate with greater or lesser facility.[19] Certain
metallic oxides also absorb ammonia and are dissolved in ammonia water.
Such are, for instance, the oxides of zinc, nickel, copper, and many
others; the majority of such compounds are unstable. The property of
ammonia of combining with certain oxides explains its action on certain
metals.[20] By reason of such action, copper vessels are not suitable
for holding liquids containing ammonia. Iron is not acted on by such
liquids.

  [19] Isambert studied the dissociation of ammoniacal compounds, as we
       have seen in Note 8, and showed that at low temperatures many
       salts are able to combine with a still greater amount of
       ammonia, which proves an entire analogy with hydrates; and as
       in this case it is easy to isolate the definite compounds, and
       as the least possible tension of ammonia is greater than that
       of water, therefore the ammoniacal compounds present a great
       and peculiar interest, as a means for explaining the nature of
       aqueous solutions and as a confirmation of the hypothesis of
       the formation of definite compounds in them; for these reasons
       we shall frequently refer to these compounds in the further
       exposition of this work.

  [20] Chapter V., Note 2.

The similarity between the relation of ammonia and water to salts and
other substances is more especially marked in those cases in which
the salt is capable of combining with both ammonia and water. Take,
for example, copper sulphate, CuSO_{4}. As we saw in Chapter I., it
gives with water blue crystals, CuSO_{4},5H_{2}O; but it also absorbs
ammonia in the same molecular proportion, forming a blue substance,
CuSO_{4},5NH_{3}, and therefore the ammonia combining with salts may be
termed _ammonia of crystallisation_.

Such are the _reactions of combination_ proper to ammonia. Let us
now turn our attention to the reactions of substitution proper to
this substance. If ammonia be passed through a heated tube containing
metallic sodium, hydrogen is evolved, and a compound is obtained
containing ammonia in which one atom of hydrogen is replaced by an
atom of sodium, NH_{2}Na (according to the equation NH_{3} + Na =
NH_{2}Na + H). This body is termed sodium amide. We shall afterwards
see that iodine and chlorine are also capable of directly displacing
hydrogen from ammonia, and of replacing it. In fact, the hydrogen of
ammonia may be replaced in many ways by different elements. If in this
replacement NH_{2} remains, the resultant substances NH_{2}R are called
_amides_, whilst the substitution products, NHR_{2}, in which only NH
remains, are called _imides_,[20 bis] and those in which none of the
ammoniacal hydrogen remains, NR_{3}, are known as _nitrides_. Free
amidogen, N_{2}H_{4}, is now known in a state of hydration under the
name of hydrazine;[21] it combines with acids and resembles ammonia in
this respect. In the action of different substances on ammonia it is
the _hydrogen that is substituted_, whilst the nitrogen remains in the
resultant compound, so to say, untouched. The same phenomenon is to be
observed in the action of various substances on water. In the majority
of cases the reactions of water consist in the hydrogen being evolved,
and in its being replaced by different elements. This also takes place,
as we have seen, in acids in which the hydrogen is easily displaced by
metals. This chemical mobility of hydrogen is perhaps connected with the
great lightness of the atoms of this element.

  [20 bis] Imide, NH, has not been obtained in a free state, but its
       hydrochloric acid salt, NHHCl, has apparently been obtained (1890)
       by Maumené by igniting the double bichloride of platinum and
       ammonium chloride, PtCl_{2}NH_{4}Cl = Pt + 2HCl + NHHCl. It is
       soluble in water, and crystallises from its solution in hexagonal
       rhombic prisms. It gives a double salt with FeCl_{3} of the
       composition FeCl_{3}3NHHCl. The salt NHHCl is similar (isomeric)
       with the first possible product of the metalepsis of ammonia,
       NH_{2}Cl, although it does not resemble it in any of its
       properties.

  [21] Free _amidogen_ or _hydrazine_, N_{2}H_{4}, or 2NH_{2}, was
       prepared by Curtius (1887) by means of ethyl diazoacetate,
       or triazoacetic acid. Curtius and Jay (1889) showed that
       triazoacetic acid, CHN_{2}.COOH (the formula should be tripled),
       when heated with water or a mineral acid, gives (quantitatively)
       oxalic acid and amidogen (hydrazine), CHN_{2}.COOH + 2H_{2}O =
       C_{2}O_{2}(OH)_{2} + N_{2}H_{4}--_i.e._ (empirically), the oxygen
       of the water replaces the nitrogen of the azoacetic acid. The
       amidogen is thus obtained in the form of a salt. With acids,
       amidogen forms very stable salts of the two types N_{2}H_{4}HX
       and N_[2]H_{4}H_{2}X_{2}, as, for example, with HCl, H_{2}SO_{4},
       &c. These salts are easily crystallised; in acid solutions
       they act as powerful reducing agents, evolving nitrogen; when
       ignited they are decomposed into ammoniacal salts, nitrogen,
       and hydrogen; with nitrites they evolve nitrogen. The sulphate
       N_{2}H_{4},H_{2}SO_{4} is sparingly soluble in cold water (3
       parts in 100 of water), but is very soluble in hot water; its
       specific gravity is 1·378, it fuses at 254° with decomposition.
       The hydrochloride N_{2}H_{4},2HCl crystallises in octahedra, is
       very soluble in water, but not in alcohol; it fuses at 198°,
       evolving hydrogen chloride and forming the salt N_{2}H_{4}HCl;
       when rapidly heated it decomposes with an explosion; with
       platinic chloride it immediately evolves nitrogen, forming
       platinous chloride. By the action of alkalis the salts
       N_{2}H_{4},2HX give _hydrate of amidogen_, N_{2}H_{4},H_{2}O,
       which is a fuming liquid (specific gravity 1·03), boiling at
       119°, almost without odour, and whose aqueous solution corrodes
       glass and india-rubber, has an alkaline taste and poisonous
       properties. The reducing capacities of the hydrate are clearly
       seen from the fact that it reduces the metals platinum and
       silver from their solutions. With mercuric oxide it explodes. It
       reacts directly with the aldehydes RO, forming N_{2}R_{2} and
       water; for example, with benzaldehydes it gives the very stable
       insoluble _benzalazine_ (C_{6}H_{5}CHN)_{2} of a yellow colour.
       We may add that hydrazine often forms double salts; for example,
       MgSO_{4}N_{2}H_{4}H_{2}SO_{4} or KClN_{2}H_{4}HCl, and that it is
       also formed by the action of nitrous acid upon aldehyde-ammonia.
       The products of the substitution of the hydrogen in hydrazine by
       hydrocarbon groups R (R = CH_{3}, C_{2}H_{5}, C_{6}H_{5}, &c.)
       were obtained before hydrazine itself; for example, NHRNH_{2},
       NR_{2}NH_{2}, and (NRH)_{2}.

       The heat of solution of the sulphuric acid salt (1 part in 200
       and 300 parts of water at 10°·8) is equal to -8·7 C. According
       to Berthelot and Matigon (1892), the heat of neutralisation of
       hydrazine by sulphuric acid is +5·5 C and by hydrochloric acid
       +5·2 C. Thus hydrazine is a very feeble base, for its heat of
       saturation is not only lower than that of ammonia (+12·4 C.
       for HCl), but even below that of hydroxylamine (+9·3 C.) The
       heat of formation from the elements of hydrated hydrazine -9·5 C
       was deduced from the heat of combustion, determined by burning
       N_{2}H_{4}H_{2}SO_{4} in a calorimetric bomb, +127·7 C. Thus
       hydrazine is an endothermal compound; its passage into ammonia
       by the combination of hydrogen is accompanied by the evolution
       of 51·5 C. In the presence of an acid these figures were greater
       by +14·4 C. Hence the direct converse passage from ammonia into
       hydrazine is impossible. As regards the passage of hydroxylamine
       into hydrazine, it would be accompanied by the evolution of heat
       (+21·5 C.) in an aqueous solution.

       Amidogen must be regarded as a compound which stands to ammonia
       in the same relation as hydrogen peroxide stands to water.
       Water, H(OH), gives, according to the law of substitution,
       as was clearly to be expected, (OH)(OH)--that is, peroxide
       of hydrogen is the free radicle of water (hydroxyl). So also
       ammonia, H(NH_{2}), forms hydrazine, (NH_{2})(NH_{2})--that is,
       the free radicle of ammonia, NH_{2}, or amidogen. In the case of
       phosphorus a similar substance, as we shall afterwards see, has
       long been known under the name of liquid phosphuretted hydrogen,
       P_{2}H_{4}.

In practical chemistry[21 bis] ammonia is often employed, not only
for saturating acids, but also for effecting reactions of double
decomposition with salts, and especially for separating insoluble
basic hydroxides from soluble salts. Let MHO stand for an insoluble
basic hydroxide and HX for an acid. The salt formed by them will have a
composition MHO + XH-H_{2}O = MX. If aqueous ammonia, NH_{4}OH, be added
to a solution of this salt, the ammonia will change places with the
metal M, and thus form the insoluble basic hydroxide, or, as it is said,
give a precipitate.

         MX          +    NH_{4}(OH)    =    NH_{4}X     +     MHO
  Salt of the metal.   Aqueous ammonia.   Ammonium salt.   Basic hydrate.
    In solution          In solution        In solution    As precipitate

  [21 bis] In practice, the applications of ammonia are very varied. The
       use of ammonia as a stimulant, in the forms of the so-called
       'smelling salts' or of spirits of hartshorn, in cases of
       faintness, &c., is known to everyone. The volatile carbonate of
       ammonium, or a mixture of an ammonium salt with an alkali, is also
       employed for this purpose. Ammonia also produces a well-known
       stimulating effect when rubbed on the skin, for which reason it is
       sometimes employed for external applications. Thus, for instance,
       the well-known volatile salve is prepared from any liquid oil
       shaken up with a solution of ammonia. A portion of the oil is thus
       transformed into a soapy substance. The solubility of greasy
       substances in ammonia, which proceeds from the formation both of
       emulsions and soaps, explains its use in extracting grease spots.
       It is also employed as an external application for stings from
       insects, and for bites from poisonous snakes, and in general in
       medicine. It is also remarkable that in cases of drunkenness a few
       drops of ammonia in water taken internally rapidly renders a
       person sober. A large quantity of ammonia is used in dyeing,
       either for the solution of certain dyes--for example, carmine--or
       for changing the tints of others, or else for neutralising the
       action of acids. It is also employed in the manufacture of
       artificial pearls. For this purpose the small scales of a peculiar
       small fish are mixed with ammonia, and the liquid so obtained is
       blown into small hollow glass beads shaped like pearls.

       In nature and the arts, however, ammonium salts, and not free
       ammonia, are most frequently employed. In this form a portion of
       that _nitrogen_ which is necessary for the formation of albuminous
       substances is _supplied to plants_. Owing to this, a large
       quantity of ammonium sulphate is now employed as a fertilising
       substance. But the same effect may be produced by nitre, or by
       animal refuse, which in decomposing gives ammonia. For this
       reason, an ammoniacal (hydrogen) compound may be introduced into
       the soil in the spring which will be converted into a nitrate
       (oxygen salt) in the summer.

Thus, for instance, if aqueous ammonia is added to a solution of a salt
of aluminium, then alumina hydrate is separated out as a colourless
gelatinous precipitate.[22]

  [22] As certain basic hydrates form peculiar compounds with ammonia,
       in some cases it happens that the first portions of ammonia
       added to a solution of a salt produce a precipitate, whilst
       the addition of a fresh quantity of ammonia dissolves this
       precipitate if the ammoniacal compound of the base be soluble in
       water. This, for example, takes place with the copper salts. But
       alumina does not dissolve under these circumstances.

In order to grasp the relation between ammonia and the oxygen
compounds of nitrogen it is necessary to recognise the general
_law of substitution_, applicable to all cases of substitution
between elements,[23] and therefore showing what may be the cases of
substitution between oxygen and hydrogen as component parts of water.
The law of substitution may be deduced from mechanical principles if
the molecule be conceived as a system of elementary atoms occurring
in a certain chemical and mechanical equilibrium. By likening the
molecule to a system of bodies in a state of motion--for instance,
to the sum total of the sun, planets, and satellites, existing in
conditions of mobile equilibrium--then we should expect the action
of one part, in this system, to be equal and opposite to the other,
according to Newton's third law of mechanics. Hence, given a molecule
of a compound, for instance, H_{2}O, NH_{3}, NaCl, HCl, &c., its every
two parts must in a chemical sense represent two things somewhat
alike in force and properties, and therefore _every two parts into
which a molecule of a compound may be divided are capable of replacing
each other_. In order that the application of the law should become
clear it is evident that among compounds the most stable should be
chosen. We will therefore take hydrochloric acid and water as the
most stable compounds of hydrogen.[24] According to the above law of
substitution, if the elements H and Cl are able to form a molecule, HCl,
and a stable one, they are able to replace each other. And, indeed,
we shall afterwards see (Chapter XI.) that in a number of instances a
substitution between hydrogen and chlorine can take place. Given RH,
then RCl is possible, because HCl exists and is stable. The molecule of
water, H_{2}O, may be divided in two ways, because it contains 3 atoms:
into H and (HO) on the one hand, and into H_{2} and O on the other.
Consequently, being given RH, its substitution products will be R(HO)
according to the first form, and R_{2}O according to the second; being
given RH_{2}, its corresponding substitution products will be RH(OH),
R(OH)_{2}, RO, (RH)_{2}O, &c. The group (OH) is the same hydroxyl or
aqueous radicle which we have already mentioned in the third chapter
as a component part of hydroxides and alkalis--for instance, Na(OH),
Ca(OH)_{2}, &c. It is evident, judging from H(HO) and HCl, that (OH)
can be substituted by Cl, because both are replaceable by H; and this
is of common occurrence in chemistry, because metallic chlorides--for
example, NaCl and NH_{4}Cl--correspond with hydroxides of the alkalis
Na(OH) or NH_{4}(OH). In hydrocarbons--for instance, C_{2}H_{6}--the
hydrogen is replaceable by chlorine and by hydroxyl. Thus ordinary
alcohol is C_{2}H_{6}, in which one atom of H is replaced by (OH); that
is, C_{2}H_{5}(OH). It is evident that the replacement of hydrogen by
hydroxyl essentially forms the phenomenon of oxidation, because RH gives
R(OH), or RHO. Hydrogen peroxide may in this sense be regarded as water
in which the hydrogen is replaced by hydroxyl; H(OH) gives (OH)_{2} or
H_{2}O_{2}. The other form of substitution--namely, that of O in the
place of H_{2}--is also a common chemical phenomenon. Thus alcohol,
C_{2}H_{6}O, or C_{2}H_{5}(OH), when oxidising in the air, gives acetic
acid, C_{2}H_{4}O_{2}, or C_{2}H_{3}O(OH), in which H_{2} is replaced
by O.

  [23] When the element chlorine, as we shall afterwards more fully
       learn, replaces the element hydrogen, the reaction by which
       such an exchange is accomplished proceeds as a substitution, AH
       + Cl_{2} = ACl + HCl, so that two substances, AH and chlorine,
       react on each other, and two substances, ACl and HCl, are
       formed; and further, two molecules react on each other, and
       two others are formed. The reaction proceeds very easily, but
       the substitution of one element, _A_, by another, _X_, does
       not always proceed with such ease, clearness, or simplicity.
       The substitution between oxygen and hydrogen is very rarely
       accomplished by the reaction of the free elements, but the
       substitution between these elements, one for another, forms
       the most common case of oxidation and reduction. In speaking
       of the law of substitution, I have in view the substitution of
       the elements one by another, and not the direct reaction of
       substitution. The law of substitution determines the cycle of the
       combinations of a given element, if a few of its compounds (for
       instance, the hydrogen compounds) be known. A development of the
       conceptions of the law of substitution may be found in my lecture
       given at the Royal Institution in London, 1889.

  [24] If hydrogen peroxide be taken as a starting point, then still
       higher forms of oxidation than those corresponding with water
       should be looked for. They should possess the properties of
       hydrogen peroxide, especially that of parting with their oxygen
       with extreme ease (even by contact). Such compounds are known.
       Pernitric, persulphuric, and similar acids present these
       properties, as we shall see in describing them.

In the further course of this work we shall have occasion to refer to
the law of substitution for explaining many chemical phenomena and
relations.

We will now apply these conceptions to ammonia in order to see its
relation to the oxygen compounds of nitrogen. It is evident that many
substances should be obtainable from ammonia, NH_{3}, or aqueous
ammonia, NH_{4}(OH), by substituting their hydrogen by hydroxyl,
or H_{2} by oxygen. And such is the case. The two extreme cases of
such substitution will be as follows: (1) One atom of H in NH_{3} is
substituted by (OH), and NH_{2}(OH) is produced. Such a substance, still
containing much hydrogen, should have many of the properties of ammonia.
It is known under the name of _hydroxylamine_,[25] and, in fact, is
capable, like ammonia, of giving salts with acids; for example, with
hydrochloric acid, NH_{3}(OH)Cl--which is a substance corresponding to
sal-ammoniac, in which one atom of hydrogen is replaced by hydroxyl.[25
bis] (2) The other extreme case of substitution is that given by
ammonium hydroxide, NH_{4}(OH), when the whole of the hydrogen of the
ammonium is replaced by oxygen; and, as ammonium contains 4 atoms of
hydrogen, the highest oxygen compound should be NO_{2}(OH), or NHO_{3},
as we find to be really the case, for NHO_{3} is nitric acid, exhibiting
the highest degree of oxidation of nitrogen.[26] If instead of the two
extreme aspects of substitution we take an intermediate one, we obtain
the intermediate oxygen compounds of nitrogen. For instance, N(OH)_{3}
is orthonitrous acid,[27] to which corresponds nitrous acid, NO(OH), or
NHO_{2}, equal to N(OH)_{3}-H_{2}O, and nitrous anhydride, N_{2}O_{3} =
2N(OH)_{3}-3H_{2}O. Thus nitrogen gives a series of oxygen compounds,
which we will proceed to describe. We will, however, first show by
two examples that in the first place the passage of ammonia into the
oxygen compounds of nitrogen up to nitric acid, as well as the converse
preparation of ammonia (and consequently of the intermediate compounds
also) from nitric acid, are reactions which proceed directly and easily
under many circumstances, and in the second place that the above general
principle of substitution gives the possibility of understanding many,
at first sight unexpected and complex, relations and transformations,
such as the preparation of hydronitrous acid, HN_{3}. In nature the
matter is complicated by a number of influences and circumstances, but
in the law the relations are presented in their simplest aspect.

  [25] The compound of hydroxylamine with hydrochloric acid has the
       composition NH_{2}(OH)HCl = NH_{4}ClO--that is, it is as it were
       oxidised sal-ammoniac. It was prepared by Lossen in 1865 by the
       action of tin and hydrochloric acid in the presence of water on
       a substance called ethyl nitrate, in which case the hydrogen
       liberated from the hydrochloric acid by the tin acts upon the
       elements of nitric acid--

       C_{2}H_{5}·NO_{3} + 6H + HCl = NH_{4}OCl + H_{2}O + C_{2}H_{5}·OH
         Ethyl nitrate  Hydrogen     Hydroxylamine  Water      Alcohol
                          from           + HCl
                       HCl and Sn

       Thus in this case the nitric acid is deoxidised, not directly
       into nitrogen, but into hydroxylamine. Hydroxylamine is also
       formed by passing nitric oxide, NO, into a mixture of tin and
       hydrochloric acid--that is, by the action of the hydrogen evolved
       on the nitric oxide, NO + 3H + HCl = NH_{4}OCl--and in many other
       cases. According to Lossen's method, a mixture of 30 parts of
       ethyl nitrate, 120 parts of tin, and 40 parts of a solution of
       hydrochloric acid of sp. gr. 1·06 are taken. After a certain
       time the reaction commences spontaneously. When the reaction
       has ceased the tin is separated by means of hydrogen sulphide,
       the solution is evaporated, and a large amount of sal-ammoniac
       is thus obtained (owing to the further action of hydrogen on
       the hydroxylamine compound, the hydrogen taking up oxygen from
       it and forming water); a solution ultimately remains containing
       the hydroxylamine salt; this salt is dissolved in anhydrous
       alcohol and purified by the addition of platinum chloride, which
       precipitates any ammonium salt still remaining in the solution.
       After concentrating the alcoholic solution the hydroxylamine
       hydrochloride separates in crystals. This substance melts at
       about 150°, and in so doing decomposes into nitrogen, hydrogen
       chloride, water, and sal-ammoniac. A sulphuric acid compound of
       hydroxylamine may be obtained by mixing a solution of the above
       salt with sulphuric acid. The sulphate is also soluble in water
       like the hydrochloride; this shows that hydroxylamine, like
       ammonia itself, forms a series of salts in which one acid may
       be substituted for another. It might he expected that by mixing
       a strong solution of a hydroxylamine salt with a solution of a
       caustic alkali hydroxylamine itself would be liberated, just
       as an ammonia salt under these circumstances evolves ammonia;
       but the liberated hydroxylamine is immediately decomposed with
       the formation of nitrogen and ammonia (and probably nitrous
       oxide), 3NH_{3}O = NH_{3} + 3H_{2}O + N_{2}. Dilute solutions
       give the same reaction, although very slowly, but by decomposing
       a solution of the sulphate with barium hydroxide a certain
       amount of hydroxylamine is obtained in solution (it is partly
       decomposed). Hydroxylamine in aqueous solution, like ammonia,
       precipitates basic hydrates, and it deoxidises the oxides of
       copper, silver, and other metals. Free hydroxylamine was obtained
       by Lobry de Bruyn (1891). It is a solid, colourless, crystalline
       substance, without odour, which does not melt below 27°. It has
       the property of dissolving metallic salts; for instance, sodium
       chloride. Hydroxylamine, when rapidly heated with platinum,
       decomposes with a flash and the formation of a yellow flame.
       It is almost insoluble in ordinary solvents like chloroform,
       benzine, acetic ether, and carbon bisulphide. Its aqueous
       solutions are tolerably stable, contain up to 60 per cent.
       (sp. gr. 1·15 at 20°), and may be kept for many weeks without
       undergoing any change. Lobry de Bruyn used the hydrochloric
       salt to prepare pure hydroxylamine. The salt was first treated
       with sodium methylate (CH_{3}NaO), and then methyl alcohol was
       added to the mixture. The precipitated sodium chloride was
       separated from the solution by filtration. (The methyl alcohol
       is added to prevent the precipitated chloride of sodium from
       coating the insoluble hydrochloric salt of hydroxylamine.) The
       methyl alcohol was driven off under a pressure 150-200 mm., and
       after extracting a further portion of methyl alcohol by ether
       and several fractional distillations, a solution was obtained
       containing 70 per cent. of free hydroxylamine, 8 per cent.
       water, 9·9 per cent. chloride of sodium, and 12·1 per cent. of
       the hydrochloric salt of hydroxylamine. Pure free hydroxylamine,
       NH_{3}O, is obtained by distilling under a pressure of 60 mm.;
       it then boils at 70°, and solidifies in a condenser cooled to
       0° in the form of long needles. It melts at 33°, boils at 58°
       under a pressure of 22 mm., and has a sp. gr. of about 1·235
       (Brühl). Under the action of NaHO it gives NH_{3} and NHO_{2}
       or N_{2}O, and forms nitric acid (Kolotoff, 1893) under the
       action of oxidising agents. Hydroxylamine is obtained in a great
       number of cases, for instance by the action of tin on dilute
       nitric acid, and also by the action of zinc on ethyl nitrate and
       dilute hydrochloric acid, &c. The relation between hydroxylamine,
       NH_{2}(OH), and nitrous acid, NO(OH), which is so clear in the
       sense of the law of substitutions, becomes a reality in those
       cases when reducing agents act on salts of nitrous acid. Thus
       Raschig (1888) proposed the following method for the preparation
       of the hydroxylamine sulphate. A mixture of strong solutions of
       potassium nitrite, KNO_{2}, and hydroxide, KHO, in molecular
       proportions, is prepared and cooled. An excess of sulphurous
       anhydride is then passed into the mixture, and the solution
       boiled for a long time. A mixture of the sulphates of potassium
       and hydroxylamine is thus obtained: KNO_{2} + KHO + 2SO_{2} +
       2H_{2}O = NH_{2}(OH),H_{2}SO_{4} + K_{2}SO_{4}. The salts may be
       separated from each other by crystallisation.

  [25 bis] In order to illustrate the application of the law of
       substitution to a given case, and to show the connection between
       ammonia and the oxides of nitrogen, let us consider the possible
       products of an oxygen and hydroxyl substitution in caustic
       ammonia, NH_{4}(OH). It is evident that the substitution of H by
       OH can give: (1) NH_{3}(OH)_{2}; (2) NH_{2}(OH)_{3}; (3)
       NH(OH)_{4}; and (4) N(OH)_{5}. They should all, like caustic
       ammonia itself, easily part with water and form products
       (hydroxylic) of the oxidation of ammonia. The first of them is the
       hydrate of hydroxylamine, NH_{2}(OH) + H_{2}O; the second,
       NH(OH)_{2} + H_{2}O (and also the substance NH(OH)_{4} or
       NH_{3}O_{2}), containing, as it does, both hydrogen and oxygen, is
       able to part with all its hydrogen in the form of water (which
       could not be done by the first product, since it contained too
       little oxygen), forming, as the ultimate product,
       2NH_{2}(OH)_{3}-5H_{2}O = N_{2}O--that is, it corresponds with
       nitrous oxide, or the lower degree of the oxidation of nitrogen.
       So, also, nitrous anhydride corresponds with the third of the
       above products, 2NH(OH)_{4}-5H_{2}O = N_{2}O_{3}, and nitric
       anhydride with the fourth, 2N(OH)_{5}-5H_{2}O = N_{2}O_{5}. As, in
       these three equations, two molecules of the substitution products
       (-5H_{2}O) are taken, it is also possible to combine two different
       products in one equation. For instance, the third and fourth
       products: NH(OH)_{4} + N(OH)_{5}-5H_{2}O corresponds to N_{2}O_{4}
       or 2NO_{2}, that is, to peroxide of nitrogen. Thus all the five
       (see later) oxides of nitrogen, N_{2}O, NO, N_{2}O_{3}, NO_{2},
       and N_{2}O_{5}, may be deduced from ammonia. The above may be
       expressed in a general form by the equation (it should be remarked
       that the composition of all the substitution products of caustic
       ammonia may be expressed by NH_{3}O{5-_a_}, where _a_ varies
       between 0 and 4):

         NH_{5}O_{5 - _a_} + NH_{5}O{5 - _b_} - 5H_{2}O
                                            = N_{2}O_{5 - (_a_ + _b_)},

       where _a_ + _b_ can evidently be not greater than 5; when _a_ +
       _b_ = 5 we have N_{2}--nitrogen, when = 4 we have N_{2}O nitrous
       oxide; when _a_ + _b_ = 3 we have N_{2}O_{2} or NO--nitric oxide,
       and so on to N_{2}O_{5}, when _a_ + _b_ = 0. Besides which it is
       evident that intermediate products may correspond with (and hence
       also break up into) different starting points; for instance,
       N_{2}O is obtained when _a_ + _b_ = 2, and this may occur either
       when _a_ = 0 (nitric acid), and _b_ = 2 (hydroxylamine), or when
       _a_ = _b_ = 1 (the third of the above substitution products).

  [26] Nitric acid corresponds with the anhydride N_{2}O_{5}, which will
       afterwards be described, but which must be regarded as the
       highest saline oxide of nitrogen, just as Na_{2}O (and the
       hydroxide NaHO) in the case of sodium, although sodium forms
       a peroxide possessing the property of parting with its oxygen
       with the same ease as hydrogen peroxide, if not on heating,
       at all events in reactions--for instance, with acids. So also
       nitric acid has its corresponding peroxide, which may be called
       pernitric acid. Its composition is not well known--probably
       NHO_{4}--so that its corresponding anhydride would be N_{2}O_{7}.
       It is formed by the action of a silent discharge on a mixture
       of nitrogen and oxygen, so that a portion of its oxygen is
       in a state similar to that in ozone. The instability of this
       substance (obtained by Hautefeuille, Chappuis, and Berthelot),
       which easily splits up with the formation of nitric peroxide, and
       its resemblance to persulphuric acid, which we shall afterwards
       describe, will permit our passing over the consideration of the
       little that is further known concerning it.

  [27] Phosphorus (Chapter XIX.) gives the hydride PH_{3}, corresponding
       with ammonia, NH_{3}, and forms phosphorous acid, PH_{3}O_{3},
       which is analogous to nitrous acid, just as phosphoric acid is
       to nitric acid; but phosphoric (or, better, orthophosphoric)
       acid, PH_{3}O_{4}, is able to lose water and give pyro-and
       meta-phosphoric acids. The latter is equal to the ortho-acid
       minus water = PHO_{3}, and therefore nitric acid, NHO_{3},
       is really meta-nitric acid. So also nitrous acid, HNO_{2},
       is meta-nitrous (anhydrous) acid, and thus the ortho-acid is
       NH_{3}O_{3} = N(OH)_{3}. Hence for nitric acid we should expect
       to find, besides the ordinary or meta-nitric acid, HNO_{3}
       (= 1/2N_{2}O_{3},H_{2}O), and ortho-nitric acid, H_{3}NO_{4}
       (= 1/2N_{2}O_{3},3H_{2}O), an intermediate pyro-nitric acid,
       N_{2}H_{4}O_{7}, corresponding to pyrophosphoric acid,
       P_{2}H_{4}O_{7}. We shall see (for instance, in Chapter XVI.,
       Note 21) that in nitric acid there is indeed an inclination
       of the ordinary salts (of the meta-acid), MNO_{3}, to combine
       with bases M_{2}O, and to approximate to the composition of
       ortho-compounds which are equal to meta-compound and bases
       (MNO_{3} + M_{2}O = M_{3}NO_{4}).

1. It is easy to prove the possibility of the oxidation of ammonia into
nitric acid by passing a mixture of ammonia and air over heated spongy
platinum. This causes the oxidation of the ammonia, nitric acid being
formed, which partially combines with the excess of ammonia.

The converse passage of nitric acid into ammonia is effected by the
action of hydrogen at the moment of its evolution.[28] Thus metallic
aluminium, evolving hydrogen from a solution of caustic soda, is able to
completely convert nitric acid added to the mixture (as a salt, because
the alkali gives a salt with the nitric acid) into ammonia, NHO_{3} + 8H
= NH_{3} + 3H_{2}O.

2. In 1890 Curtius in Germany obtained a gaseous substance of the
composition HN_{3} (hydrogen trinitride), having the distinctive
properties of an acid, and giving, like hydrochloric acid, salts;
for example, a sodium salt, NaN_{3}; ammonium salt, NH_{4}N_{3} =
N_{4}H_{4}; barium salt, Ba(N_{3})_{2}, &c., which he therefore named
hydronitrous acid, _HN__{3}.[28 bis] The extraordinary composition
of the compound (ammonia, NH_{3}, contains one N atom and three H
atoms; in HN_{3}, on the contrary, there are three N atoms and one H
atom), the facile decomposition of its salts with an explosion, and
above all its distinctly acid character (an aqueous solution shows a
strong acid reaction to litmus), not only indicated the importance of
this unexpected discovery, but at first gave rise to some perplexity
as to the nature of the substance obtained, for the relations in
which HN_{3} stood to other simple compounds of nitrogen which had
long been known was not at all evident, and the scientific spirit
especially requires that there should be a distinct bond between
every innovation, every fresh discovery, and that which is already
firmly established and known, for upon this basis is founded that
apparently paradoxical union in science of a conservative stability
with an irresistible and never-ceasing improvement. This missing,
connection between the newly discovered hydronitrous acid, HN_{3},
and the long known ammonia, NH_{3}, and nitric acid, HNO_{3}, may
be found in the law of substitution, starting from the well-known
properties and composition of nitric acid and ammonia, as I mentioned
in the 'Journal of the Russian Physico-Chemical Society' (1890).
The essence of the matter lies in the fact that to the hydrate of
ammonium, or caustic ammonia, NH_{4}OH, there should correspond,
according to the law of substitution, an ortho-nitric acid (_see_
Note 27), H_{3}NO_{4} = NO(OH)_{3}, which equals NH_{4}(OH) with
the substitution in it of (_a_) two atoms of hydrogen by oxygen
(O--H_{2}) and (_b_) two atoms of hydrogen by the aqueous radicle
(OH--H). Ordinary or meta-nitric acid is merely this ortho-nitric
acid minus water. To ortho-nitric acid there should correspond the
ammoniacal salts: mono-substituted, H_{2}NH_{4}NO_{4}; bi-substituted,
H(NH_{4})_{2}NO_{4}; and tri-substituted, (NH_{4})_{3}NO_{4}. These
salts, containing as they do hydrogen and oxygen, like many similar
ammoniacal salts (see, for instance, Chapter IX.--Cyanides), are able
to part with them in the form of water. Then from the first salt we
have H_{2}NH_{4}NO_{4}-4H_{2}O = N_{2}O--nitrous oxide, and from the
second H(NH_{4})_{2}NO4-4H_{2}O = HN_{3}--hydronitrous acid, and from
the third (NH_{4})_{3}NO-4H_{2}O = N_{4}H_{4}--the ammonium salt of the
same acid. The composition of HN_{3} should be thus understood, whilst
its acid properties are explained by the fact that the water (4H_{2}O)
from H(NH_{4})_{2}N_O{4} is formed at the expense of the hydrogen of the
ammonium and oxygen of the nitric acid, so that there remains the same
hydrogen as in nitric acid, or that which may be replaced by metals and
give salts. Moreover, nitrogen undoubtedly belongs to that category of
metalloids which give acids, like chlorine and carbon, and therefore,
under the influence of three of its atoms, one atom of hydrogen acquires
those properties which it has in acids, just as in HCN (hydrocyanic
acid) the hydrogen has received these properties under the influence of
the carbon and nitrogen (and HN_{3} may be regarded as HCN where C has
been replaced by N_{2}). Moreover, besides explaining the composition
and acid properties of HN_{3}, the above method gives the possibility
of foretelling the closeness of the bond between hydronitrous acid and
nitrous oxide, for N_{2}O + NH_{3} = HN_{3} + H_{2}O. This reaction,
which was foreseen from the above considerations, was accomplished
by Wislicenus (1892) by the synthesis of the sodium salt, by taking
the amide of sodium, NH_{2}Na (obtained by heating Na in a current of
NH_{3}), and acting upon it (when heated) with nitrous oxide, N_{2}O,
when 2NH_{2}Na + N_{2}O = NaN_{3} + NaHO + NH_{3}. The resultant salt,
NaN_{3}, gives hydronitrous acid when acted upon by sulphuric acid,
NaN_{3} + H_{2}SO_{4} = NaHSO_{4} +HN_{3}. The latter gives, with the
corresponding solutions of their salts, the insoluble (and easily
explosive) salts of silver, AgN_{3} (insoluble, like AgCl or AgCN), and
lead, Pb(N_{3})_{2}.

  [28] The formation of ammonia is observed in many cases of oxidation by
       means of nitric acid. This substance is even formed in the action
       of nitric acid on tin, especially if dilute acid be employed in
       the cold. A still more considerable amount of ammonia is obtained
       if, in the action of nitric acid, there are conditions directly
       tending to the evolution of hydrogen, which then reduces the acid
       to ammonia; for instance, in the action of zinc on a mixture of
       nitric and sulphuric acids.

  [28 bis] Curtius started with benzoylhydrazine, C_{6}H_{5}CONHNH_{2}
       (hydrazine, see Note 20 bis). (This substance is obtained by
       the action of hydrated hydrazine on the compound ether of benzoic
       acid). Benzoylhydrazine under the action of nitrous acid gives
       benzoylazoimide and water:

         C_{6}H_{5}CONHNH_{2} + NO_{2}H = C_{6}H_{5}CON_{3} + 2H_{2}O.

       Benzoylazoimide when treated with sodium alcoholate gives the
       sodium salt of hydronitrous acid:

         C_{6}H_{5}CON_{3} + C_{2}H_{3}ONa
                                 = C_{6}H_{5}O_{2}C_{2}H_{3} + NaN_{3}.

       The addition of ether to the resultant solution precipitates the
       NaN_{3}, and this salt when treated with sulphuric acid gives
       gaseous hydronitrous acid, HN_{3}. It has an acrid smell, and is
       easily soluble in water. The aqueous solution exhibits a strongly
       acid reaction. Metals dissolve in this solution and give the
       corresponding salts. With hydronitrous acid gaseous ammonia forms
       a white cloud, consisting of the salt of ammonium, NH_{4}N_{3}.
       This salt separates out from an alcoholic solution in the form of
       white lustrous scales. The salts of hydronitrous acid are obtained
       by a reaction of substitution with the sodium or ammonium salts.
       In this manner Curtius obtained and studied the salts of silver
       (AgN_{3}), mercury (HgN_{3}), lead (PbN_{6}), barium (BaN_{6}).
       With hydrazine, N_{2}H_{4}, hydronitrous acid forms saline
       compounds in the composition of which there are one or two
       particles of N_{3}H per one particle of hydrazine; thus N_{5}H_{5}
       and N_{8}H_{6}. The first was obtained in an almost pure form. It
       crystallises from an aqueous solution in dense, volatile, lustrous
       prisms (up to 1 in. long), which fuse at 50°, and deliquesce in
       the air; from a solution in boiling alcohol it separates out in
       bright crystalline plates. This salt, N_{5}H_{5}, has the same
       empirical composition, NH, as the ammonium salt of hydronitrous
       acid, N_{4}H_{4}, and imide; but their molecules and structure are
       different. Curtius also obtained (1893) hydronitrous acid by
       passing the vapour of N_{2}O_{5} (evolved by the action of HNO_{3}
       on As_{2}O_{3}) into a solution of hydrazine, N_{2}H_{4}.
       Similarly Angeli, by acting upon a saturated solution of silver
       nitrite with a strong solution of hydrazine, obtained the
       explosive AgN_{3} in the form of a precipitate, and this reaction,
       which is based upon the equation N_{2}H_{4} + NHO_{2} = HN_{3} +
       2H_{2}O, proceeds so easily that it forms an experiment for the
       lecture table. A thermal investigation of hydronitrous acid by
       Berthelot and Matignon gave the following figures for the heat of
       solution of the ammonium salt N_{3}HNH_{3} (1 grm. in 100 parts of
       water)-708 C., and for the heat of neutralisation by barium
       hydrate +10·0 C., and by ammonia +8·2 C. The heat of combustion of
       N_{4}H_{4} (+163·8 C. at a constant vol.) gives the heat of
       formation of the salt N_{4}H_{4} (solid) as -25·3 C. and
       (solution)-32·3 C.; this explains the explosive nature of this
       compound. In its heat of formation from the elements N_{3}H =-62·6
       C., this compound differs from all the hydrogen compounds of
       nitrogen in having a maximum absorption of heat, which explains
       its instability.

The compounds of nitrogen with oxygen present an excellent example of
the law of multiple proportions, because they contain, for 14 parts by
weight of nitrogen, 8, 16, 24, 32, and 40 parts respectively by weight
of oxygen. The composition of these compounds is as follows:--

             N_{2}O, nitrous oxide; hydrate NHO.
             N_{2}O_{2}, nitric oxide, NO.
             N_{2}O_{3}, nitrous anhydride; hydrate NHO_{2}.
             N_{2}O_{4}, peroxide of nitrogen, NO_{2}.
             N_{2}O_{5}, nitric anhydride; hydrate NHO_{3}.

Of these compounds,[29] nitrous and nitric oxides, peroxide of
nitrogen, and nitric acid, NHO_{3}, are characterised as being the most
stable. _The lower oxides, when coming into contact with the higher,
may give the intermediate forms_; for instance, NO and NO_{2} form
N_{2}O_{3}, _and the intermediate oxides may, in splitting up, give a
higher and lower oxide_. So N_{2}O_{4} gives N_{2}O_{3} and N_{2}O_{5},
or, in the presence of water, their hydrates.

  [29] According to the thermochemical determinations of Favre, Thomsen,
       and more especially of Berthelot, it follows that, in the
       formation of such quantities of the oxides of nitrogen as express
       their formulæ, if gaseous nitrogen and oxygen be taken as the
       starting points, and if the compounds formed be also gaseous, the
       following amounts of heat, expressed in thousands of heat units,
       are _absorbed_ (hence a minus sign):--

         N_{2}O   N_{2}O_{2}    N_{2}O_{3}    N_{2}O_{4}    N_{2}O_{5}
          -21        -43           -22           -5            -1
              -22           +21           +17            +4

       The difference is given in the lower line. For example, if
       N_{2}, or 28 grams of nitrogen, combine with O--that is, with 16
       grams of oxygen--then 21,000 units of heat are absorbed, that
       is, sufficient heat to raise 21,000 grams of water through 1°.
       Naturally, direct observations are impossible in this case; but
       if charcoal, phosphorus, or similar substances are burnt both in
       nitrous oxide and in oxygen, and the heat evolved is observed
       in both cases, then the difference (more heat will be evolved
       in burning in nitrous oxide) gives the figures required. If
       N_{2}O_{2}, by combining with O_{2}, gives N_{2}O_{4}, then, as
       is seen from the table, heat should be developed, namely, 38,000
       units of heat, or NO + O = 19,000 units of heat. The differences
       given in the table show that the maximum absorption of heat
       corresponds with nitric oxide, and that the higher oxides are
       formed from it with evolution of heat. If liquid nitric acid,
       NHO_{3}, were decomposed into N + O_{3} + H, then 41,000 heat
       units would be required; that is, an evolution of heat takes
       place in its formation from the gases. It should be observed
       that the formation of ammonia, NH_{3}, from the gases N + H_{3}
       evolves 12·2 thousand heat units.

We have already seen that, under certain conditions, nitrogen combines
with oxygen, and we know that ammonia may he oxidised. In these cases
various oxidation products of nitrogen are formed, but in the presence
of water and an excess of oxygen they always give nitric acid. Nitric
acid, as corresponding with the highest oxide, is able, in deoxidising,
to give the lower oxides; it is the only nitrogen acid whose salts occur
somewhat widely in nature, and it has many technical uses, for which
reason we will begin with it.

_Nitric acid_, NHO_{3}, is likewise known as aqua fortis. In a free
state it is only met with in nature in small quantities, in the air
and in rain-water after storms; but even in the atmosphere nitric
acid does not long remain free, but combines with ammonia, traces of
which are always found in air. On falling on the soil and into running
water, &c., the nitric acid everywhere comes into contact with bases
(or their carbonates), which easily act on it, and therefore it is
converted into the nitrates of these bases. Hence nitric acid is always
met with in the form of salts in nature. The soluble salts of nitric
acid are called _nitres_. This name is derived from the Latin _sal
nitri_. The potassium salt, KNO_{3}, is common nitre, and the sodium
salt, NaNO_{3}, Chili saltpetre, or cubic nitre. Nitres are formed
in the soil when a nitrogenous substance is slowly oxidised in the
presence of an alkali by means of the oxygen of the atmosphere. In
nature there are very frequent instances of such oxidation. For this
reason certain soils and rubbish heaps--for instance, lime rubbish (in
the presence of a base)--lime contain a more or less considerable amount
of nitre. One of these nitres--sodium nitrate--is extracted from the
earth in large quantities in Chili, where it was probably formed by the
oxidation of animal refuse. This kind of nitre is employed in practice
for the manufacture of nitric acid and the other oxygen compounds of
nitrogen. Nitric acid is obtained from _Chili saltpetre_ by heating it
with _sulphuric acid_. The hydrogen of the sulphuric acid replaces the
sodium in the nitre. The sulphuric acid then forms either an acid salt,
NaHSO_{4}, or a normal salt, Na_{2}SO_{4}, whilst nitric acid is formed
from the nitre and is volatilised. The decomposition is expressed by
the equations: (1) NaNO_{3} + H_{2}SO_{4} = HNO_{3} + NaHSO_{4}, if the
acid salt be formed, and (2) 2NaNO_{3} + H_{2}SO_{4} = Na_{2}SO_{4}
+ 2HNO_{3}, if the normal sodium sulphate is formed. With an excess
of sulphuric acid, at a moderate heat, and at the commencement of the
reaction, the decomposition proceeds according to the first equation;
and on further heating with a sufficient amount of nitre according to
the second, because the acid salt NaHSO_{4} itself acts like an acid
(its hydrogen being replaceable as in acids), according to the equation
NaNO_{3} + NaHSO_{4} = Na_{2}SO_{4} + HNO_{3}.

[Illustration: FIG. 47.--Method of preparing nitric acid on a large
scale. A cast-iron retort, C, is fixed into the furnace, and heated
by the fire, B. The flame and products of combustion are at first led
along the flue, M (in order to heat the receivers), and afterwards into
L. The retort is charged with Chili saltpetre and sulphuric acid, and
the cover is luted on with clay and gypsum. A clay tube, _a_, is fixed
into the neck of the retort (in order to prevent the nitric acid from
corroding the cast iron), and a bent glass tube, D, is luted on to it.
This tube carries the vapours into a series of earthenware receivers, E.
Nitric acid mixed with sulphuric acid collects in the first. The purest
nitric acid is procured from the second, whilst that which condenses in
the third receiver contains hydrochloric acid, and that in the fourth
nitrous oxide. Water is poured into the last receiver in order to
condense the residual vapours.]

The sulphuric acid, as it is said, here displaces the nitric acid from
its compound with the base.[29 bis] Thus, in the reaction of sulphuric
acid on nitre there is formed a non-volatile salt of sulphuric acid,
which remains, together with an excess of this acid, in the distilling
apparatus, and nitric acid, which is converted into vapour, and may be
condensed, because it is a liquid and volatile substance. On a small
scale, this reaction may be carried on in a glass retort with a glass
condenser. On a large scale, in chemical works, the process is exactly
similar, only iron retorts are employed for holding the mixture of
nitre and sulphuric acid, and earthenware three-necked bottles are used
instead of a condenser,[30] as shown in fig. 47.

  [29 bis] This often gives rise to the supposition that sulphuric acid
        possesses a considerable degree of affinity or energy compared
       with nitric acid, but we shall afterwards see that the idea of the
       relative degree of affinity of acids and bases is, in many cases,
       exceedingly unbiassed; it need not be accepted so long as it is
       possible to explain the observed phenomena without admitting any
       supposition whatever of the degree of the force of affinity,
       because the latter cannot be measured. The action of sulphuric
       acid upon nitre may be explained by the fact alone that the
       resultant nitric acid is volatile. The nitric acid is the only one
       of all the substances partaking in the reaction which is able to
       pass into vapour; it alone is volatile, while the remainder are
       non-volatile, or, more strictly speaking, exceedingly difficultly
       volatile substances. Let us imagine that the sulphuric acid is
       only able to set free a small quantity of nitric acid from its
       salt, and this will suffice to explain the decomposition of the
       whole of the nitre by the sulphuric acid, because once the nitric
       acid is separated it passes into vapour when heated, and passes
       away from the sphere of action of the remaining substances; then
       the free sulphuric acid will set free a fresh small quantity of
       nitric acid, and so on until it drives off the entire quantity. It
       is evident that, in this explanation, it is essential that the
       sulphuric acid should be in excess (although not greatly)
       throughout the reaction; according to the equation expressing the
       reaction, 98 parts of sulphuric acid are required per 85 parts of
       Chili nitre; but if this proportion be maintained in practice the
       nitric acid is not all disengaged by the sulphuric acid; an excess
       of the latter must be taken, and generally 80 parts of Chili nitre
       are taken per 98 parts of acid, so that a portion of the sulphuric
       acid remains free to the very end of the reaction.

  [30] It must be observed that sulphuric acid, at least when undiluted
       (60° Baumé), corrodes cast iron with difficulty, so that the
       acid may be heated in cast-iron retorts. Nevertheless, both
       sulphuric and nitric acids have a certain action on cast iron,
       and therefore the acid obtained will contain traces of iron. In
       practice sodium nitrate (Chili saltpetre) is usually employed
       because it is cheaper, but in the laboratory it is best to take
       potassium nitrate, because it is purer and does not froth up so
       much as sodium nitrate when heated with sulphuric acid. In the
       action of an excess of sulphuric acid on nitre and nitric acid
       a portion of the latter is decomposed, forming lower oxides of
       nitrogen, which are dissolved in the nitric acid. A portion
       of the sulphuric acid itself is also carried over as spray by
       the vapours of the nitric acid. Hence sulphuric acid occurs
       as an impurity in commercial nitric acid. A certain amount of
       hydrochloric acid will also be found to be present in it, because
       sodium chloride is generally found as an impurity in nitre, and
       under the action of sulphuric acid it forms hydrochloric acid.
       Commercial acid further contains a considerable excess of water
       above that necessary for the formation of the hydrate, because
       water is first poured into the earthenware vessels employed for
       condensing the nitric acid in order to facilitate its cooling and
       condensation. Further, the acid of composition HNO_{3} decomposes
       with great ease, with the evolution of oxides of nitrogen. Thus
       the commercial acid contains a great number of impurities, and
       is frequently purified in the following manner:--Lead nitrate is
       first added to the acid because it forms non-volatile and almost
       insoluble (precipitated) substances with the free sulphuric
       and hydrochloric acids, and liberates nitric acid in so doing,
       according to the equations Pb(NO_{3})_{2} + 2HCl = PbCl_{2} +
       2NHO_{3} and Pb(NO_{3})_{2} + H_{2}SO_{4} = PbSO_{4} + 2NHO_{3}.
       Potassium chromate is then added to the impure nitric acid,
       by which means oxygen is liberated from the chromic acid, and
       this oxygen, at the moment of its evolution, oxidises the lower
       oxides of nitrogen and converts them into nitric acid. A pure
       nitric acid, containing no impurities other than water, may be
       then obtained by carefully distilling the acid, treated as above
       described, and particularly if only the middle portions of the
       distillate are collected. Such acid should give no precipitate,
       either with a solution of barium chloride (a precipitate shows
       the presence of sulphuric acid) or with a solution of silver
       nitrate (a precipitate shows the presence of hydrochloric acid),
       nor should it, after being diluted with water, give a coloration
       with starch containing potassium iodide (a coloration shows the
       admixture of other oxides of nitrogen). The oxides of nitrogen
       may be most easily removed from impure nitric acid by heating
       for a certain time with a small quantity of pure charcoal. By
       the action of nitric acid on the charcoal carbonic anhydride
       is evolved, which carries off the lower oxides of nitrogen. On
       redistilling, pure acid is obtained. The oxides of nitrogen
       occurring in solution may also be removed by passing air through
       the nitric acid.

Nitric acid so obtained always contains water. It is extremely difficult
to deprive it of all the admixed water without destroying a portion of
the acid itself and partially converting it into lower oxides, because
without the presence of an excess of water it is very unstable. When
rapidly distilled a portion is decomposed, and there are obtained free
oxygen and lower oxides of nitrogen, which, together with the water,
remain in solution with the nitric acid. Therefore it is necessary to
work with great care in order to obtain a pure hydrate of nitric acid,
HNO_{3}, and especially to mix the nitric acid obtained from nitre, as
above described, with sulphuric acid, which takes up the water, and
to distil it at the lowest possible temperature--that is, by placing
the retort holding the mixture in a water or oil bath and carefully
heating it. The first portion of the nitric acid thus distilled boils
at 86°, has a specific gravity at 15° of 1·526, and solidifies at -50°;
it is very unstable at higher temperatures. This is the normal hydrate,
HNO_{3}, which corresponds with the salts, NMO_{3}, of nitric acid.
When diluted with water nitric acid presents a higher boiling point,
not only as compared with that of the nitric acid itself, but also with
that of water; so that, if very dilute nitric acid be distilled, the
first portions passing over will consist of almost pure water, until
the boiling point in the vapours reaches 121°. At this temperature a
compound of nitric acid with water, containing about 70 p.c. of nitric
acid,[31] distils over; its specific gravity at 15° = 1·421. If the
solution contain less than 25 p.c. of water, then, the specific gravity
of the solution being above 1·44, HNO_{3} evaporates off and fumes in
the air, forming the above hydrate, whose vapour tension is less than
that of water. Such solutions form _fuming nitric acid_. On distilling
it gives monohydrated acid,[32] HNO_{3}; it is a hydrate boiling at
121°, so that it is obtained from both weak and strong solutions. Fuming
nitric acid, under the action not only of organic substances, but
even of heat, loses a portion of its oxygen, forming lower oxides of
nitrogen, which impart a _red-brown colour_ to it;[33] the pure acid is
colourless.

  [31] Dalton, Smith, Bineau, and others considered that the hydrate of
       constant boiling point (see Chapter I., Note 60) for nitric
       acid was the compound 2HNO_{3},3H_{2}O, but Roscoe showed that
       its composition changes with a variation of the pressure and
       temperature under which the distillation proceeds. Thus, at
       a pressure of 1 atmosphere the solution of constant boiling
       point contains 68·6 p.c., and at one-tenth of an atmosphere
       66·8 p.c. Judging from what has been said concerning solutions
       of hydrochloric acid, and from the variation of specific
       gravity, I think that the comparatively large decrease in the
       tensions of the vapours depends on the formation of a hydrate,
       NHO_{3},2H_{2}O (= 63·6 p.c.). Such a hydrate may be expressed by
       N(HO)_{5}, that is, as NH_{4}(HO), in which all the equivalents
       of hydrogen are replaced by hydroxyl. The constant boiling point
       will then be the temperature of the decomposition of this hydrate.

       The variation of the specific gravity at 15° from water (_p_
       = 0) to the hydrate NHO_{3},5H_{2}O (41·2 p.c. HNO_{3}) is
       expressed by _s_ = 9992 + 57·4_p_ + 0·16_p^2_, if water = 10,000
       at 4°. For example, when _p_ = 30 p.c., _s_ = 11,860. For more
       concentrated solutions, at least, the above-mentioned hydrate,
       HNO_{3},2H_{2}O, must be taken, up to which the specific gravity
       _s_ = 9570 + 84·18_p_-0·240_p_^2; but perhaps (since the results
       of observations of the specific gravity of the solutions are not
       in sufficient agreement to arrive at a conclusion) the hydrate
       HNO_{3},3H_{2}O should be recognised, as is indicated by many
       nitrates (Al, Mg, Co, &c.), which crystallise with this amount
       of water of crystallisation. From HNO_{3},2H_{2}O to HNO_{3}
       the specific gravity of the solutions (at 15°) _s_ = 10,652 +
       62·08_p_-0·160_p_^2. The hydrate HNO_{3},2H_{2}O is recognised by
       Berthelot on the basis of the thermochemical data for solutions
       of nitric acid, because on approaching to this composition there
       is a rapid change in the amount of heat evolved by mixing nitric
       acid with water. Pickering (1892) by refrigeration obtained
       the crystalline hydrates: HNO_{3},H_{2}O, melting at -37° and
       HNO_{3},3H_{2}O, melting at -18°. A more detailed study of the
       reactions of hydrated nitric acid would no doubt show the
       existence of change in the process and rapidity of reaction in
       approaching these hydrates.

  [32] The normal hydrate HNO_{3}, corresponding with the ordinary salts,
       may be termed the monohydrated acid, because the anhydride
       N_{2}O_{5} with water forms this normal nitric acid. In this
       sense the hydrate HNO_{3},2H_{2}O is the pentahydrated acid.

  [33] For technical and laboratory purposes recourse is frequently had
       to _red fuming nitric acid_--that is, the normal nitric acid,
       HNO_{3}, containing lower oxides of nitrogen in solution. This
       acid is prepared by decomposing nitre with half its weight
       of strong sulphuric acid, or by distilling nitric acid with
       an excess of sulphuric acid. The normal nitric acid is first
       obtained, but it partially decomposes, and gives the lower
       oxidation products of nitrogen, which are dissolved by the nitric
       acid, to which they impart its usual pale-brown or reddish
       colour. This acid fumes in the air, from which it attracts
       moisture, forming a less volatile hydrate. If carbonic anhydride
       be passed through the red-brown fuming nitric acid for a long
       period of time, especially if, assisted by a moderate heat, it
       expels all the lower oxides, and leaves a colourless acid free
       from these oxides. It is necessary, in the preparation of the red
       acid, that the receivers should be kept quite cool, because it
       is only when cold that nitric acid is able to dissolve a large
       proportion of the oxides of nitrogen. The strong red fuming acid
       has a specific gravity 1·56 at 20°, and has a suffocating smell
       of the oxides of nitrogen. When the red acid is mixed with water
       it turns green, and then of a bluish colour, and with an excess
       of water ultimately becomes colourless. This is owing to the fact
       that the oxides of nitrogen in the presence of water and nitric
       acid are changed, and give  solutions.

       Markleffsky (1892) showed that the green solutions contain
       (besides HNO_{3}) HNO_{2} and N_{2}O_{4}, whilst the blue
       solutions only contain HNO_{2} (_see_ Note 48).

       The action of red fuming nitric acid (or a mixture with sulphuric
       acid) is in many cases very powerful and rapid, and it sometimes
       acts differently from pure nitric acid. Thus iron becomes
       covered with a coating of oxides, and insoluble in acids; it
       becomes, as is said, passive. Thus chromic acid (and potassium
       dichromate) gives oxide of chromium in this red acid--that is, it
       is deoxidised. This is owing to the presence of the lower oxides
       of nitrogen, which are capable of being oxidised--that is, of
       passing into nitric acid like the higher oxides. But, generally,
       the action of fuming nitric acid, both red and colourless, is
       powerfully oxidising.

Nitric acid, as an _acid hydrate_, enters into reactions of double
decomposition with bases, basic hydrates (alkalis), and with salts.
In all these cases a salt of nitric acid is obtained. An alkali and
nitric acid give water and a salt; so, also, a basic oxide with nitric
acid gives a salt and water; for instance, lime, CaO + 2HNO_{3} =
Ca(NO_{3})_{2} + H_{2}O. Many of these salts are termed nitres.[34]
The composition of the ordinary salts of nitric acid may be expressed
by the general formula M(NO_{3})_{_n_}, where M indicates a metal
replacing the hydrogen in one or several (_n_) equivalents of nitric
acid. We shall find afterwards that the atoms M of metals are equivalent
to one (K, Na, Ag) atom of hydrogen, or two (Ca, Mg, Ba), or three
(Al, In), or, in general, _n_ atoms of hydrogen. _The salts of nitric
acid_ are especially characterised by being all _soluble in water_.[35]
From the property common to all these salts of entering into double
decompositions, and owing to the volatility of nitric acid, they evolve
nitric acid when heated with sulphuric acid. They all, like the acid
itself, are capable of evolving oxygen when heated, and consequently of
acting as oxidising substances; they therefore, for instance, deflagrate
with ignited carbon, the carbon burning at the expense of the oxygen of
the salt and forming gaseous products of combustion.[36]

  [34] Hydrogen is not evolved in the action of nitric acid (especially
       strong) on metals, even with those metals which evolve hydrogen
       under the action of other acids. This is because the hydrogen
       at the moment of its separation reduces the nitric acid, with
       formation of the lower oxides of nitrogen, as we shall afterwards
       see.

  [35] Certain basic salts of nitric acid, however (for example, the
       basic salt of bismuth), are insoluble in water; whilst, on the
       other hand, all the normal salts are soluble, and this forms an
       exceptional phenomenon among acids, because all the ordinary
       acids form insoluble salts with one or another base. Thus, for
       sulphuric acid the salts of barium, lead, &c., for hydrochloric
       acid the salts of silver, &c., are insoluble in water. However,
       the normal salts of acetic and certain other acids are all
       soluble.

  [36] _Ammonium nitrate_, NH_{4}NO_{3}, is easily obtained by adding a
       solution of ammonia or of ammonium carbonate to nitric acid until
       it becomes neutral. On evaporating this solution, crystals of
       the salt are formed which contain no water of crystallisation.
       It crystallises in prisms like those formed by common nitre, and
       has a refreshing taste; 100 parts of water at _t_° dissolve 54 +
       0·61_t_ parts by weight of the salt. It is soluble in alcohol,
       melts at 160°, and is decomposed at about 180°, forming water
       and nitrous oxide, NH_{4}NO_{3} = 2H_{2}O + N_{2}O. If ammonium
       nitrate be mixed with sulphuric acid, and the mixture be heated
       to about the boiling point of water, then nitric acid is evolved,
       and ammonium hydrogen sulphate remains in solution; but if the
       mixture be heated rapidly to 16O°, then nitrous oxide is evolved.
       In the first case the sulphuric acid takes up ammonia, and in the
       second place water. Ammonium nitrate is employed in practice for
       the artificial production of cold, because in dissolving in water
       it lowers the temperature very considerably. For this purpose it
       is best to take equal parts by weight of the salt and water. The
       salt must first be reduced to a powder and then rapidly stirred
       up in the water, when the temperature will fall from +15° to -10°,
       so that the water freezes.

       Ammonium nitrate absorbs ammonia, with which it forms
       unstable compounds resembling compounds containing water
       of crystallisation. (Divers 1872, Raoult 1873.) At -10°
       NH_{4}NO_{3},2NH_{3} is formed: it is a liquid of sp. gr. 1·15,
       which loses all its ammonia under the influence of heat. At +28°
       NH_{4}NO_{3},NH_{3} is formed: it is a solid which easily parts
       with its ammonia when heated, especially in solution.

       Troost (1882) investigated the tension of the dissociation of
       the compounds formed, and came to the conclusion that a definite
       compound corresponding to the formula 2NH_{4}NO_{3},3NH_{3} is
       formed, because the tension of dissociation remains constant
       in the decomposition of such a compound at 0°. Y. Kouriloff
       (1893), however, considers that the constancy of the tension of
       the ammonia evolved is due to the decomposition of a saturated
       solution, and not of a definite compound. During decomposition
       the system is composed of a liquid and a solid; the tension
       only becomes constant from the moment the solid falls down. The
       composition 2NH_{4}NO_{3},3NH_{3} corresponds to a saturated
       solution at 0°, and the solubility of NH_{4}NO_{3} in NH_{3}
       increases with a rise of temperature.

Nitric acid also enters into double decompositions with a number of
hydrocarbons not in any way possessing alkaline characters and not
reacting with other acids. Under these circumstances the nitric acid
gives water and a new substance termed a _nitro-compound_. The chemical
character of the nitro-compound is the same as that of the original
substance; for example, if an indifferent substance be taken, then the
nitro compound obtained from it will also be indifferent; if an acid
be taken, then an acid is obtained also.[36 bis] Benzene, C_{6}H_{6},
for instance, acts according to the equation C_{6}H_{6} + HNO_{3} =
H_{2}O + C_{6}H_{5}NO_{2}. Nitrobenzene is produced. The substance
taken, C_{6}H_{6}, is a liquid hydrocarbon having a faint tarry smell,
boiling at 80°, and lighter than water; by the action of nitric acid
nitrobenzene is obtained, which is a substance boiling at about 210°,
heavier than water, and having an almond-like odour: it is employed in
large quantities for the preparation of aniline and aniline dyes.[37]
As the nitro-compounds contain both combustible elements (hydrogen
and carbon), as well as oxygen in unstable combination with nitrogen,
in the form of the radicle NO_{2} of nitric acid, they decompose
with an explosion when ignited or even struck, owing to the pressure
of the vapours and gases formed--free nitrogen, carbonic anhydride,
CO_{2}, carbonic oxide, CO, and aqueous vapour. In the explosion of
nitro compounds[37 bis] much heat is evolved, as in the combustion of
gunpowder or detonating gas, and in this case the force of explosion in
a closed space is great, because from a solid or liquid nitro-compound
occupying a small space there proceed vapours and gases whose elasticity
is great not only from the small space in which they are formed, but
owing to the high temperature corresponding to the combustion of the
nitro-compound.[38]

  [36 bis] This is explained by saying that in true nitro-compounds
       the residue of nitric acid NO_{2} takes the place of the hydrogen
       in the hydrocarbon group. For example, if C_{6}H_{5}OH be given,
       then C_{6}H_{4}(NO_{2})OH will be a true nitro-compound having the
       radical properties of C_{6}H_{5}OH. If, on the other hand, the
       NO_{2} replace the hydrogen of the aqueous radicle
       (C_{6}H_{5}ONO_{2}), then the chemical character varies, as in the
       passage of KOH into KONO_{2} (nitre) (_see_ Note 37 and Organic
       Chemistry).

  [37] The compound ethers of nitric acid in which the hydrogen of the
       aqueous radicle (OH) is replaced by the residue of nitric acid
       (NO_{2}) are frequently called nitro-compounds. But in their
       chemical character they differ from true nitro-compounds (for
       details _see_ Organic Chemistry) and do not burn like them.

       The action of nitric acid on cellulose, C_{6}H_{10}O_{5}, is an
       example. This substance, which forms the outer coating of all
       plant cells, occurs in an almost pure state in cotton, in common
       writing-paper, and in flax, &c.; under the action of nitric
       acid it forms water and nitrocellulose (like water and KNO_{3}
       from KHO), which, although it has the same appearance as the
       cotton originally taken, differs from it entirely in properties.
       It explodes when struck, bursts into flame very easily under
       the action of sparks, and acts like gunpowder, whence its name
       of pyroxylin, or gun-cotton. The composition of gun-cotton is
       C_{6}H_{7}N_{3}O_{11} = C_{6}H_{10}O_{5} + 3NHO_{3}-3H_{2}O.
       The proportion of the group NO_{2} in nitrocellulose may be
       decreased by limiting the action of the nitric acid and compounds
       obtained with different properties; for instance, the (impure)
       well-known _collodion cotton_, containing from 11 to 12 per
       cent. of nitrogen, and _pyro-collodion_ (Mendeléeff, 1890),
       containing 12·4 per cent. of nitrogen. Both these products are
       soluble in a mixture of alcohol and ether (in collodion a portion
       of the substance is soluble in alcohol), and the solution when
       evaporated gives a transparent film, which is insoluble in water.
       A solution of collodion is employed in medicine for covering
       wounds, and in wet-plate photography for giving on glass an even
       coating of a substance into which the various reagents employed
       in the process are introduced. Extremely fine threads (obtained
       by forcing a gelatinous mixture of collodion, ether, and alcohol
       through capillary tubes in water) of collodion form artificial
       silk.

  [37 bis] The property possessed by nitroglycerin (occurring in
       dynamite), nitrocellulose, and the other nitro-compounds, of
       burning with an explosion, and their employment for smokeless
       powder and as explosives in general, depends on the reasons in
       virtue of which a mixture of nitre and charcoal deflagrates and
       explodes; in both cases the elements of the nitric acid occurring
       in the compound are decomposed, the oxygen in burning unites with
       the carbon, and the nitrogen is set free; thus a very large volume
       of gaseous substances (nitrogen and oxides of carbon) is rapidly
       formed from the solid substances originally taken. These gases
       occupy an incomparably larger volume than the original substance,
       and therefore produce a powerful pressure and explosion. It is
       evident that in exploding with the development of heat (that is,
       in decomposing, not with the absorption of energy, as is generally
       the case, but with the evolution of energy) the nitro-compounds
       form stores of energy which are easily set free, and that
       consequently their elements occur in a state of particularly
       energetic motion, which is especially strong in the group NO_{2}:
       this group is common to all nitro-compounds, and all the oxygen
       compounds of nitrogen are unstable, easily decomposable, and (Note
       29) absorb heat in their formation. On the other hand, the
       nitro-compounds are instructive as an example and proof of the
       fact that the elements and groups forming compounds are united in
       definite order in the molecules of a compound. A blow, concussion,
       or rise of temperature is necessary to bring the combustible
       elements C and H into the most intimate contact with NO_{2}, and
       to distribute the elements in a new order in new compounds.

       As regards the composition of the nitro-compounds, it will be seen
       that the hydrogen of a given substance is replaced by the complex
       group NO_{2} of the nitric acid. The same is observed in the
       passage of alkalis into nitrates, so that the reactions of
       substitution of nitric acid--that is, the formation of salts and
       nitro-compounds--may be expressed in the following manner. In
       these cases the hydrogen is replaced by the so-called _radicle of
       nitric acid_ NO_{2}, as is evident from the following table:--

                {Caustic potash   KHO.
                {Nitre            K(NO_{2})O.

                {Hydrate of lime  CaH_{2}O_{2}.
                {Calcium nitrate  Ca(NO_{2})_{2}O_{2}.

                {Glycerin         C_{3}H_{5}H_{3}O_{3}.
                {Nitroglycerin    C_{3}H_{5}(NO_{2})_{3}O_{3}.

                {Phenol           C_{6}H_{5}OH.
                {Picric acid      C_{6}H_{2}(NO_{2})_{3}OH, &c.

       The difference between the salts formed by nitric acid and the
       nitro-compounds consists in the fact that nitric acid is very
       easily separated from the salts of nitric acid by means of
       sulphuric acid (that is, by a method of double saline
       decomposition), whilst nitric acid is not displaced by sulphuric
       acid from true nitro-compounds; for instance, nitrobenzene,
       C_{6}H_{5}·NO_{2}. As nitro-compounds are formed exclusively from
       hydrocarbons, they are described with them in organic chemistry.

       The group NO_{2} of nitro-compounds in many cases (like all the
       oxidised compounds of nitrogen) passes into the ammonia group or
       into the ammonia radicle NH_{2}. This requires the action of
       reducing substances evolving hydrogen: RNO_{2} + 6H = RNH_{3} +
       2H_{2}O. Thus Zinin converted nitrobenzene, C_{6}H_{5}·NO_{2},
       into aniline, C_{6}H_{5}·NH_{2}, by the action of hydrogen
       sulphide.

       Admitting the existence of the group NO_{2}, as replacing hydrogen
       in various compounds, then nitric acid may be considered as water
       in which half the hydrogen is replaced by the radical of nitric
       acid. In this sense nitric acid is nitro-water, NO_{2}OH, and its
       anhydride dinitro-water, (NO_{2})_{2}O. In nitric acid the radical
       of nitric acid is combined with hydroxyl, just as in nitrobenzene
       it is combined with the radical of benzene.

       It should here be remarked that the group NO_{3} may be recognised
       in the salts of nitric acid, because the salts have the
       composition M(NO_{3})_{n}, just as the metallic chlorides have the
       composition MCl_{n}. But the group NO_{3} does not form any other
       compounds beyond the salts, and therefore it should he considered
       as hydroxyl, HO, in which H is replaced by NO_{2}.

  [38] The nitro-compounds play a very important part in mining and
       artillery. Detailed accounts of them must be looked for in
       special works, among which the works of A. R. Shuliachenke
       and T. M. Chelletsoff occupy an important place in the
       Russian literature on this subject, although historically the
       scientific works of Abel in England and Berthelot in France
       stand pre-eminent. The latter elucidated much in connection
       with explosive compounds by a series of both experimental and
       theoretical researches. Among explosives a particularly important
       place from a practical point of view is occupied by ordinary or
       black gunpowder (Chapter XIII., Note 16), fulminating mercury
       (Chapter XVI., Note 26), the different forms of gun-cotton
       (Chapter VI., Note 37), and nitro-glycerine (Chapter VIII.,
       Note 45, and Chapter XII., Note 33). The latter when mixed
       with solid pulverulent substances, like magnesia, tripoli,
       &c., forms dynamite, which is so largely used in quarries and
       mines in driving tunnels, &c. We may add that the simplest true
       nitro-compound, or marsh gas, CH_{4}, in which all the hydrogens
       are replaced by NO_{2} groups has been obtained by L. N.
       Shishkoff, C(NO_{2})_{4}, as well as nitroform, CH(NO_{2})_{3}.

[Illustration: FIG. 48.--The method of decomposition of nitrous
anhydride, also applicable to the other oxides of nitrogen, and to
their analysis. NO_{2} is generated from nitrate of lead in the retort
A. Nitric acid and other less volatile products are condensed in B.
The tube C C contains copper, and is heated from below. Undecomposed
volatile products (if any are formed) are condensed in D, which is
cooled. If the decomposition be incomplete, brown fumes make their
appearance in this receiver. The gaseous nitrogen is collected in the
cylinder E.]

If the vapour of nitric acid is passed through an even moderately
heated glass tube, the formation of dark-brown fumes of the lower
oxides of nitrogen and the separation of free oxygen may be observed,
2NHO_{3} = H_{2}O + 2NO_{2} + O. The decomposition is complete at a
white heat--that is, nitrogen is formed, 2NHO_{3} = H_{2}O + N_{2} +
O_{5}. Hence it is easily understood that nitric acid may part with its
oxygen to a number of substances capable of being oxidised.[39] It is
consequently an _oxidising agent_. Charcoal, as we have already seen,
burns in nitric acid; phosphorus, sulphur, iodine, and the majority of
metals also decompose nitric acid, some on heating and others even at
the ordinary temperature: the substances taken are oxidised and the
nitric acid is deoxidised, yielding compounds containing less oxygen.
Only a few metals, such as gold and platinum, do not act on nitric acid,
but the majority decompose it; in so doing, an oxide of the metal is
formed, which, if it has the character of a base, acts on the remaining
nitric acid; hence, with the majority of metals the result of the
reaction is usually not an oxide of the metal, but the corresponding
salt of nitric acid, and, at the same time, one of the lower oxides of
nitrogen. The resulting salts of the metals are soluble, and hence it
is said that nitric acid _dissolves_ nearly all metals.[40] This case
is termed the solution of metals by acids, although it is not a case
of simple solution, but a complex chemical change of the substances
taken. When treated with this acid, those metals whose oxides do not
combine with nitric acid yield the oxide itself, and not a salt; for
example, tin acts in this manner on nitric acid, forming a hydrated
oxide, SnH_{2}O_{3}, which is obtained in the form of a white powder,
Sn + 4NHO_{3} = H_{2}SnO_{3} + 4NO_{2} + H_{2}O. Silver is able to
take up still more oxygen, and to convert a large portion of nitric
acid into nitrous anhydride, 4Ag + 6HNO_{3} = 4AgNO_{3} + N_{2}O_{3} +
3H_{2}O. Copper takes up still more oxygen from nitric acid, converting
it into nitric oxide, and, by the action of zinc, nitric acid is able
to give up a still further quantity of nitrogen, forming nitrous oxide,
4Zn + 10NHO_{3} = 4Zn(NO_{3})_{2} + N_{2}O + 5H_{2}O.[41] Sometimes,
and especially with dilute solutions of nitric acid, the deoxidation
proceeds as far as the formation of hydroxylamine and ammonia, and
sometimes it leads to the formation of nitrogen itself. The formation of
one or other nitrogenous substance from nitric acid is determined, not
only by the nature of the reacting substances, but also by the relative
mass of water and nitric acid, and also by the temperature and pressure,
or the sum total of the conditions of reaction; and as in a given
mixture even these conditions vary (the temperature and the relative
mass vary), it not unfrequently happens that a mixture of different
products of the deoxidation of nitric acid is formed.

  [39] [Illustration: FIG. 49.--Decomposition of nitrous oxide by
       sodium.]

       Nitric acid may be entirely decomposed by passing its vapour
       over highly incandescent copper, because the oxides of nitrogen
       first formed give up their oxygen to the red-hot metallic copper,
       so that water and nitrogen gas alone are obtained. This forms a
       means for determining the composition both of nitric acid and
       of all the other compounds of nitrogen with oxygen, because
       by collecting the gaseous nitrogen formed it is possible to
       calculate, from its volume, its weight and consequently its
       amount in a given quantity of a nitrogenous substance, and by
       weighing the copper before and after the decomposition it is
       possible to determine the amount of oxygen by the increase
       in weight. The complete decomposition of nitric acid is also
       accomplished by passing a mixture of hydrogen and nitric acid
       vapours through a red-hot tube. Sodium also decomposes the oxides
       of nitrogen at a red-heat, taking up all the oxygen. This method
       is sometimes used for determining the composition of the oxides
       of nitrogen.

  [40] The application of this acid for etching copper or steel in
       engraving is based on this fact. The copper is covered with a
       coating of wax, resin, &c. (etching ground), on which nitric acid
       does not act, and then the ground is removed in certain parts
       with a needle, and the whole is washed in nitric acid. The parts
       coated remain untouched, whilst the uncovered portions are eaten
       into by the acid. Copper plates for etchings, aquatints, &c., are
       prepared in this manner.

  [41] The formation of such complex equations as the above often
       presents some difficulty to the beginner. It should be observed
       that if the reacting and resultant substances be known, it is
       easy to form an equation for the reaction. Thus, if we wish
       to form an equation expressing the reaction that nitric acid
       acting on zinc gives nitrous oxide, N_{2}O, and zinc nitrate,
       Zn(NO_{3})_{2}, we must reason as follows:--Nitric acid contains
       hydrogen, whilst the salt and nitrous oxide do not; hence water
       is formed, and therefore it is as though anhydrous nitric acid,
       N_{2}O_{5}, were acting. For its conversion into nitrous oxide
       it parts with four equivalents of oxygen, and hence it is able
       to oxidise four equivalents of zinc and to convert it into zinc
       oxide, ZnO. These four equivalents of zinc oxide require for
       their conversion into the salt four more equivalents of nitric
       anhydride; consequently five equivalents in all of the latter
       are required, or ten equivalents of nitric acid. Thus ten
       equivalents of nitric acid are necessary for four equivalents
       of zinc in order to express the reaction in whole equivalents.
       It must not be forgotten, however, that there are very few such
       reactions which can be entirely expressed by simple equations.
       The majority of equations of reactions only express the chief
       and ultimate products of reaction, and thus none of the three
       preceding equations express all that in reality occurs in the
       action of metals on nitric acid. In no one of them is only
       one oxide of nitrogen formed, but always several together or
       consecutively--one after the other, according to the temperature
       and strength of the acid. And this is easily intelligible. The
       resulting oxide is itself capable of acting on metals and of
       being deoxidised, and in the presence of the nitric acid it may
       change the acid and be itself changed. The equations given must
       be looked on as a systematic expression of the main features of
       reactions, or as a limit towards which they tend, but to which
       they only attain in the absence of disturbing influences.

Thus the action of nitric acid on metals consists in their being
oxidised, whilst the acid itself is converted, according to the
temperature, concentration in which it is taken, and the nature of the
metal, &c., into lower oxides, ammonia, or even into nitrogen.[42] Many
compounds are oxidised by nitric acid like metals and other elements;
for instance, lower oxides are converted into higher oxides. Thus,
arsenious acid is converted into arsenic acid, suboxide of iron into
oxide, sulphurous acid into sulphuric acid, the sulphides of the metals,
M_{2}S, into sulphates, M_{2}SO_{4}, &c.; in a word, nitric acid brings
about oxidation, its oxygen is taken up and transferred to many other
substances. Certain substances are oxidised by strong nitric acid so
rapidly and with so great an evolution of heat that they deflagrate
and burst into flame. Thus turpentine, C_{10}H_{16}, bursts into
flame when poured into fuming nitric acid. In virtue of its oxidising
property, nitric acid _removes the hydrogen_ from many substances. Thus
it decomposes hydriodic acid, separating the iodine and forming water;
and if fuming nitric acid be poured into a flask containing gaseous
hydriodic acid, then a rapid reaction takes place, accompanied by flame
and the separation of violet vapours of iodine and brown fumes of oxides
of nitrogen.[43]

  [42] Montemartini endeavours to show that the products evolved in the
       action of nitric acid upon metals (and their amount) is in direct
       connection with both the concentration of the acid and the
       capacity of the metals to decompose water. Those metals which
       only decompose water at a high temperature give, under the action
       of nitric acid, NO_{2}, N_{2}O_{4}, and NO; whilst those metals
       which decompose water at a lower temperature give, besides the
       above products, N_{2}O, N, and NH_{3}; and, lastly, the metals
       which decompose water at the ordinary temperature also evolve
       hydrogen. It is observed that concentrated nitric acid oxidises
       many metals with much greater difficulty than when diluted
       with water; iron, copper, and tin are very easily oxidised by
       dilute nitric acid, but remain unaltered under the influence of
       monohydrated nitric acid or of the pure hydrate NHO_{3}. Nitric
       acid diluted with a large quantity of water does not oxidise
       copper, but it oxidises tin; dilute nitric acid also does not
       oxidise either silver or mercury; but, on the addition of nitrous
       acid, even dilute acid acts on the above metals. This naturally
       depends on the smaller stability of nitrous acid, and on the fact
       that after the commencement of the action the nitric acid is
       itself converted into nitrous acid, which continues to act on the
       silver and mercury. Veley (Oxford 1891) made detailed researches
       on the action of nitric acid upon Cu, Hg, and Bi, and showed that
       nitric acid of 30 p.c. strength does not act upon these metals at
       the ordinary temperature if nitrous acid (traces are destroyed by
       urea) and oxidising agents such as H_{2}O_{2}, KClO_{3}, &c. be
       entirely absent; but in the presence of even a small amount of
       nitrous acid the metals form nitrites, which, with HNO_{3}, form
       nitrates and the oxides of nitrogen, which re-form the nitrous
       acid necessary for starting the reaction, because the reaction
       2NO + HNO_{3} + H_{2}O = 3HNO_{2} is reversible. The above metals
       are quickly dissolved in a 1 p.c. solution of nitrous acid.
       Moreover, Veley observed that nitric acid is partially converted
       into nitrous acid by gaseous hydrogen in the presence of the
       nitrates of Cu and Pb.

  [43] When nitric acid acts on many organic substances it often happens
       that not only is hydrogen removed, but also oxygen is combined;
       thus, for example, nitric acid converts toluene, C_{7}H_{8},
       into benzoic acid, C_{7}H_{6}O_{2}. In certain cases, also, a
       portion of the carbon contained in an organic substance burns at
       the expense of the oxygen of the nitric acid. So, for instance,
       phthalic acid, C_{8}H_{6}O_{4}, is obtained from naphthalene,
       _{10}H_{8}. Thus the action of nitric acid on the hydrocarbons
       is often most complex; not only does nitrification take place,
       but also separation of carbon, displacement of hydrogen, and
       combination of oxygen. There are few organic substances which can
       withstand the action of nitric acid, and it causes fundamental
       changes in a number of them. It leaves a yellow stain on the
       skin, and in a large quantity causes a wound and entirely eats
       away the membranes of the body. The membranes of plants are eaten
       into with the greatest ease by strong nitric acid in just the
       same manner. One of the most durable blue vegetable dyes employed
       in dyeing tissues is _indigo_; yet it is easily _converted into a
       yellow substance_ by the action of nitric acid, and small traces
       of free nitric acid may be recognised by this means.

As nitric acid is very easily decomposed with the separation of oxygen,
it was for a long time supposed that it was not capable of forming the
corresponding _nitric anhydride_, N_{2}O_{5}; but Deville first and
subsequently Weber and others, discovered the methods of its formation.
Deville obtained nitric anhydride by decomposing silver nitrate by
chlorine under the influence of a moderate heat. Chlorine acts on
the above salt at a temperature of 95° (2AgNO_{3} + Cl_{2} = 2AgCl +
N_{2}O_{5} + O), and when once the reaction is started, it continues
by itself without further heating. Brown fumes are given off, which
are condensed in a tube surrounded by a freezing-mixture. A portion
condenses in this tube and a portion remains in a gaseous state. The
latter contains free oxygen. A crystalline mass and a liquid substance
are obtained in the tube; the liquid is poured off, and a current of dry
carbonic acid gas is passed through the apparatus in order to remove
all traces of volatile substances (liquid oxides of nitrogen) adhering
to the crystals of nitric anhydride. These form a voluminous mass of
rhombic crystals (density 1·64), which sometimes are of rather large
size; they melt at about 30° and distil at about 47°. In distilling,
a portion of the substance is decomposed. With water these crystals
give nitric acid. Nitric anhydride is also obtained by the action of
phosphoric anhydride, P_{2}O_{5}, on cold pure nitric acid (below
0°). During the very careful distillation of equal parts by weight
of these two substances a portion of the acid decomposes, giving a
liquid compound, H_{2}O,2N_{2}O_{5} = N_{2}O_{5},2HNO_{3}, whilst the
greater part of the nitric acid gives the anhydride according to the
equation 2NHO_{3} + P_{2}O_{5} = 2PHO_{3} + N_{2}O_{5}. On heating,
nitric anhydride decomposes with an explosion, or gradually, into nitric
peroxide and oxygen, N_{2}O_{5} = N_{2}O_{4} + O.

_Nitrogen peroxide_, N_{2}O_{4}, and _nitrogen dioxide_, NO_{2}, express
one and the same composition, but they should be distinguished like
ordinary oxygen and ozone, although in this case their mutual conversion
is more easily effected and takes place on vaporisation; also, O_{3}
loses heat in passing into O_{2}, whilst N_{2}O_{4} absorbs heat in
forming NO_{2}.

Nitric acid in acting on tin and on many organic substances (for
example, starch) gives brown vapours, consisting of a mixture of
N_{2}O_{3} and NO_{2}. A purer product is obtained by the decomposition
of lead nitrate by heat, Pb(NO_{3})_{2} = 2NO_{2} + O + PbO, when
non-volatile lead oxide, oxygen gas, and nitrogen peroxide are formed.
The latter condenses, in a well-cooled vessel, to a brown liquid, which
boils at about 22°. The purest peroxide of nitrogen, solidifying at -9°,
is obtained by mixing dry oxygen in a freezing-mixture with twice its
volume of dry nitric oxide, NO, when transparent prisms of nitrogen
peroxide are formed in the receiver: they melt into a colourless liquid
at about -10°. When the temperature of the receiver is above -9°, the
crystals melt,[44] and at 0° give a reddish yellow liquid, like that
obtained in the decomposition of lead nitrate. The vapours of nitrogen
peroxide have a characteristic odour, and at the ordinary temperature
are of a dark-brown colour, but at lower temperatures the colour of
the vapour is much fainter. When heated, especially above 50°, the
colour becomes a very dark brown, so that the vapours almost lose their
transparency.

  [44] According to certain investigations, if a brown liquid is formed
       from the melted crystals by beating above -9°, then they no longer
       solidify at -10°, probably because a certain amount of N_{2}O_{3}
       (and oxygen) is formed, and this substance remains liquid at
       -30°, or it may be that the passage from 2NO_{2} into N_{2}O_{4}
       is not so easily accomplished as the passage from N_{2}O_{4} into
       2NO_{2}.

       Liquid nitrogen peroxide (that is, a mixture of NO_{2} and
       N_{2}O_{4}) is employed in admixture with hydrocarbons as an
       explosive.

The causes of these peculiarities of nitrogen peroxide were not
clearly understood until Deville and Troost determined the density and
dissociation of the vapour of this substance at different temperatures,
and showed that the density varies. If the density be referred to that
of hydrogen at the same temperature and pressure, then it is found to
vary from 38 at the boiling point, or about 27°, to 23 at 135°, after
which the density remains constant up to those high temperatures at
which the oxides of nitrogen are decomposed. As on the basis of the
laws enunciated in the following chapter, the density 23 corresponds
with the compound NO_{2} (because the weight corresponding with this
molecular formula = 46, and the density referred to hydrogen as unity
is equal to half the molecular weight); therefore at temperatures above
135° the existence of nitrogen dioxide only must be recognised. It is
this gas which is of a brown colour. At a lower temperature it forms
nitrogen peroxide, N_{2}O_{4}, whose molecular weight, and therefore
density, is twice that of the dioxide. This substance, which is isomeric
with nitrogen dioxide, as ozone is isomeric with oxygen, and has twice
as great a vapour density (46 referred to hydrogen), is formed in
greater quantity the lower the temperature, and crystallises at -10°. The
reasons both of the variation of the colour of the gas (N_{2}O_{4} gives
colourless and transparent vapours, whilst those of NO_{2} are brown
and opaque) and the variation of the vapour density with the variation
of temperature are thus made quite clear; and as at the boiling point
a density 38 was obtained, therefore at that temperature the vapours
consist of a mixture of 79 parts by weight of N_{2}O_{4} with 21 parts
by weight of NO_{2}.[45] It is evident that a decomposition here takes
place the peculiarity of which consists in the fact that the product of
decomposition, NO_{2}, is polymerised (_i.e._ becomes denser, combines
with itself) at a lower temperature; that is, the reaction

                      N_{2}O_{4} = NO_{2} + NO_{2}

is a reversible reaction, and consequently the whole phenomenon
represents a _dissociation_ in a homogeneous gaseous medium, where the
original substance, N_{2}O_{4}, and the resultant, NO_{2}, are both
gases. The _measure of dissociation_ will be expressed if we find the
proportion of the quantity of the substance decomposed to the whole
amount of the substance. At the boiling point, therefore, the measure
of the decomposition of nitrogen peroxide will be 21 p.c.; at 135° it =
1, and at 10° it = 0; that is, the N_{2}O_{4} is not then decomposable.
Consequently the limits of dissociation here are -10° and 135° at the
atmospheric pressure.[46] Within the limits of these temperatures the
vapours of nitrogen peroxide have not a constant density, but, on the
other hand, above and below these limits definite substances exist. Thus
above 135° N_{2}O_{4} has ceased to exist and NO_{2} alone remains.
It is evident that at the ordinary temperature there is a partially
dissociated system or mixture of nitrogen peroxide, N_{2}O_{4}, and
nitrogen dioxide, NO_{2}. In the brown liquid boiling at 22° probably
a portion of the N_{2}O_{4} has already passed into NO_{2}, and it is
only the colourless liquid and crystalline substance at -10° that can be
considered as pure nitrogen peroxide.[47]

  [45] Because if _x_ equal the amount by weight of N_{2}O_{4}, its
       volume will = _x_/46, and the amount of NO_{2} will = 100-_x_,
       and consequently its volume will = (100-_x_)/23. But the mixture,
       having a density 38, will weigh 100; consequently its volume will
       = 100/38. Hence _x_/46 + (100-_x_)/23 = 100/38, or _x_ = 79·O.

  [46] The phenomena and laws of dissociation, which we shall consider
       only in particular instances, are discussed in detail in works
       on theoretical chemistry. Nevertheless, in respect to nitrogen
       peroxide, as an historically important example of dissociation
       in a homogeneous gaseous medium, we will cite the results of the
       careful investigations (1885-1880) of E. and L. Natanson, who
       determined the densities under various conditions of temperature
       and pressure. The degree of dissociation, expressed as above (it
       may also he expressed otherwise--for example, by the ratio of
       the quantity of substance decomposed to that unaltered), proves
       to increase at all temperatures as the pressure diminishes,
       which would he expected for a homogeneous gaseous medium, as a
       decreasing pressure aids the formation of the lightest product of
       dissociation (that having the least density or largest volume).
       Thus, in the Natansons' experiments the degree of dissociation
       at 0° increases from 10 p.c. to 30 p.c., with a decrease of
       pressure of from 251 to 38 mm.; at 49°·7 it increases from 49
       p.c. to 93 p.c., with a fall of pressure of from 498 to 27 mm.,
       and at 100° it increases from 89·2 p.c. to 99·7 p.c., with a
       fall of pressure of from 732·5 to 11·7 mm. At 130° and 150° the
       decomposition is complete--that is, only NO_{2} remains at the
       low pressures (less than the atmospheric) at which the Natansons
       made their determinations; but it is probable that at higher
       pressures (of several atmospheres) molecules of N_{2}O_{4} would
       still be formed, and it would be exceedingly interesting to trace
       the phenomena under the conditions of both very considerable
       pressures and of relatively large volumes.

  [47] Liquid nitrogen peroxide is said by Geuther to boil at 22°-26°,
       and to have a sp. gr. at 0° = 1·494 and at 15° = 1·474. It
       is evident that, in the liquid as in the gaseous state, the
       variation of density with the temperature depends, not only
       on physical, but also on chemical changes, as the amount of
       N_{2}O_{4} decreases and the amount of NO_{2} increases with
       the temperature, and they (as polymeric substances) should
       have different densities, as we find, for instance, in the
       hydrocarbons C_{5}H_{10} and C_{10}H_{20}.

       It may not be superfluous to mention here that the measurement
       of the specific heat of a mixture of the vapours of N_{2}O_{4}
       and NO_{2} enabled Berthelot to determine that the transformation
       of 2NO_{2} into N_{2}O_{4} is accompanied by the evolution of
       about 13,000 units of heat, and as the reaction proceeds with
       equal facility in either direction, it will be exothermal in the
       one direction and endothermal in the other; and this clearly
       demonstrates the possibility of reactions taking place in either
       direction, although, as a rule, reactions evolving heat proceed
       with greater ease.

The above explains the action of nitrogen peroxide on water at low
temperatures. N_{2}O_{4} then acts on water like a mixture of the
anhydrides of nitrous and nitric acids. The first, N_{2}O_{3}, may
be looked on as water in which each of the two atoms of hydrogen is
replaced by the radicle NO, while in the second each hydrogen is
replaced by the radicle NO_{2}, proper to nitric acid; and in nitrogen
peroxide one atom of the hydrogen of water is replaced by NO and the
other by NO_{2}, as is seen from the formulæ--

               H}       NO}         NO    }        NO_{2}}
               H} O;    NO} O;      NO_{2}} O;     NO_{2}} O;

  or           H_{2}O;  N_{2}O{3};  N_{2}O_{4};  N_{2}O_{5}.

In fact, nitrogen peroxide at low temperatures gives with water (ice)
both nitric, HNO_{3}, and nitrous, HNO_{2}, acids. The latter, as
we shall afterwards see, splits up into water and the anhydride,
N_{2}O_{3}. If, however, warm water act on nitrogen peroxide, only
nitric acid and monoxide of nitrogen are formed: 3NO_{2} + H_{2}O = NO +
2NHO_{3}.

Although NO_{2} is not decomposed into N and O even at 500°, still
in many cases it acts as an oxidising agent. Thus, for instance, it
oxidises mercury, converting it into mercurous nitrate, 2NO_{2} + Hg
= HgNO_{3} + NO, being itself deoxidised into nitric oxide, into which
the dioxide in many other instances passes, and from which it is easily
formed.[48]

  [48] Nitric acid of sp. gr. 1·51 in dissolving nitrogen peroxide becomes
       brown, whilst nitric acid of sp. gr. 1·32 is  greenish
       blue, and acid of sp. gr. below 1·15 remains colourless after
       absorbing nitrogen peroxide (Note 33).

_Nitrous anhydride_, N_{2}O_{3}, corresponds[49] to nitrous acid,
NHO_{2}, which forms a series of salts, the nitrites--for example,
the sodium salt NaNO_{2}, the potassium salt KNO_{2}, the ammonium
salt (NH_{4})NO_{2},[50] the silver salt AgNO_{2},[51] &c. Neither
the anhydride nor the hydrate of the acid is known in a perfectly
pure state. The anhydride has only been obtained as a very unstable
substance, and has not yet been fully investigated; and on attempting
to obtain the acid NHO_{2} from its salts, it always gives water and
the anhydride, whilst the latter, as an intermediate oxide, partially
or wholly splits up into NO + NO_{2}. But the salts of nitrous acid are
distinguished for their great stability. Potassium nitrate, KNO_{3}, may
be converted into potassium nitrite by depriving it of a portion of
its oxygen; for instance, by fusing it (at not too high a temperature)
with metals, such as lead, KNO_{3} + Pb = KNO_{2} + PbO.[51 bis]
The resultant salt is soluble in water, whilst the oxide of lead is
insoluble. With sulphuric and other acids the solution of potassium
nitrite[52] immediately evolves a brown gas, nitrous anhydride:
2KNO_{2} + H_{2}SO_{4} = K_{2}SO_{4} + N_{2}O_{3} + H_{2}O. The same
gas (N_{2}O_{3}) is obtained by passing nitric oxide at 0° through
liquid peroxide of nitrogen,[53] or by heating starch with nitric acid
of sp. gr. 1·3. At a very low temperature it condenses into a blue
liquid boiling below 0°,[54] but then partially decomposing into NO +
NO_{2}. Nitrous anhydride possesses a remarkable capacity for oxidising.
Ignited bodies burn in it, nitric acid absorbs it, and then acquires
the property of acting on silver and other metals, even when diluted.
_Potassium iodide_ is oxidised by this gas just as it is by ozone (and
by peroxide of hydrogen, chromic and other acids, but not by dilute
nitric acid nor by sulphuric acid), with the _separation of iodine_.
This iodine may he recognised (_see_ Ozone, Chapter IV.) by its turning
starch blue. Very small traces of nitrites may be easily detected by
this method. If, for example, starch and potassium iodide are added
to a solution of potassium nitrite (at first there will be no change,
there being no free nitrous acid), and then sulphuric acid be added, the
nitrous acid (or its anhydride) immediately set free liberates iodine,
which produces a blue colour with the starch. Nitric acid does not act
in this manner, but in the presence of zinc the coloration takes place,
which proves the formation of nitrous acid in the deoxidation of nitric
acid.[55] Nitrous acid acts directly on ammonia, forming nitrogen and
water, HNO_{2} + NH_{3} = N_{2} + 2H_{2}O.[56]

  [49] Nitrogen peroxide as a mixed substance has no corresponding
       independent salts, but Sabatier and Senderens (1892) showed that
       under certain conditions NO_{2} combines directly with some
       metals--for instance, copper and cobalt--forming Cu_{2}NO_{2} and
       CoNO_{2} as dark brown powders, which do not, however, exhibit
       the reactions of salts. Thus by passing gaseous nitrogen dioxide
       over freshly reduced (from the oxides by heating with hydrogen)
       copper at 25°-30°, Cu_{2}NO_{2} is directly formed. With water
       it partly gives off NO_{2} and partly forms nitrite of copper,
       leaving metallic copper and its suboxide. The nature of these
       compounds has not yet been sufficiently investigated.

  [50] Ammonium nitrite may be easily obtained in solution by a similar
       method of double decomposition (for instance, of the barium salt
       with ammonium sulphate) to the other salts of nitrous acid, but
       it decomposes with great ease when evaporated, with evolution
       of gaseous nitrogen, as already mentioned (Chapter V.) If the
       solution, however, be evaporated at the ordinary temperature
       under the receiver of an air-pump, a solid saline mass is
       obtained, which is easily decomposed when heated. The dry salt
       even decomposes with an explosion when struck, or when heated to
       about 70°--NH_{4}NO_{2} = 2H_{2}O + N_{2}. It is also formed by
       the action of aqueous ammonia on a mixture of nitric oxide and
       oxygen, or by the action of ozone on ammonia, and in many other
       instances. Zörensen (1894) prepared NH_{4}NO_{2} by the action
       of a mixture of N_{2}O_{3} and other oxides of nitrogen on lumps
       of ammonium carbonate, extracting the nitrite of ammonium formed
       with absolute alcohol, and precipitating it from this solution
       by ether. This salt is crystalline, dissolves in water with
       absorption of heat, and attracts moisture from the air. The solid
       salt and its concentrated solutions decompose with an explosion
       when heated to 50°-80°, especially in the presence of traces
       of foreign acids. Decomposition also proceeds at the ordinary
       temperature, but more slowly; and in order to preserve the salt
       it should be covered with a layer of pure dry ether.

  [51] Silver nitrite, AgNO_{2}, is obtained as a very slightly soluble
       substance, as a precipitate, on mixing solutions of silver
       nitrate, AgNO_{3}, and potassium nitrite, KNO_{2}. It is soluble
       in a large volume of water, and this is taken advantage of
       to free it from silver oxide, which is also present in the
       precipitate, owing to the fact that potassium nitrite always
       contains a certain amount of oxide, which with water gives the
       hydroxide, forming oxide of silver with silver nitrate. The
       solution of silver nitrite gives, by double decomposition with
       metallic chlorides (for instance, barium chloride), insoluble
       silver chloride and the nitrite of the metal taken (in this case,
       barium nitrite, Ba(NO_{2})_{2}).

  [51 bis] Leroy (1889) obtained KNO_{2} by mixing powdered KNO_{3} with
       BaS, igniting the mixture in a crucible and washing the fused
       salts; BaSO_{4} is then left as an insoluble residue, and KNO_{2}
       passes into solution: 4KNO_{3} + BaS = 4KNO_{2} + BaSO_{4}.

  [52] Probably potassium nitrite, KNO_{2}, when strongly heated,
       especially with metallic oxides, evolves N and O, and gives
       potassium oxide, K_{2}O, because nitre is liable to such a
       decomposition; but it has, as yet, been but little investigated.

  [53] There are many researches which lead to the conclusion that the
       reaction N_{2}O_{3} = NO_{2}-NO is reversible, _i.e._ resembles
       the conversion of N_{2}O_{4} into NO_{2}. The brown colour of the
       fumes of N_{2}O_{3} is due to the formation of NO_{2}.

       If nitrogen peroxide be cooled to -20°, and half its weight
       of water be added to it drop by drop, then the peroxide is
       decomposed, as we have already said, into nitrous and nitric
       acids; the former does not then remain as a hydrate, but
       straightway passes into the anhydride, and, hence, if the
       resultant liquid be slightly warmed vapours of nitrous anhydride,
       N_{2}O_{3}, are evolved, and condense into a blue liquid, as
       Fritzsche showed. This method of preparing nitrous anhydride
       apparently gives the purest product, but it easily dissociates,
       forming NO and NO_{2} (and therefore also nitric acid in the
       presence of water).

  [54] According to Thorpe, N_{2}O_{3} boils at +18°. According to
       Geuther, at +3°·5, and its sp. gr. at 0° = 1·449.

  [55] In its oxidising action nitrous anhydride gives nitric oxide,
       N_{2}O_{3} = 2NO + O. Thus its analogy to ozone becomes still
       more marked, because in ozone it is only one-third of the oxygen
       that acts in oxidising; from O_{3} there is obtained O, which
       acts as an oxidiser, and common oxygen O_{2}. In a physical
       aspect the relation between N_{2}O_{3} and O_{3} is revealed in
       the fact that both substances are of a blue colour when in the
       liquid state.

  [56] This reaction is taken advantage of for converting the amides,
       NH_{2}R (where R is an element or a complex group) into
       hydroxides, RHO. In this case NH_{2}R + NHO_{2} forms 2N
       + H_{2}O + RHO; NH_{2}, is replaced by HO, the radicle of
       ammonia by the radicle of water. This reaction is employed
       for transforming many nitrogenous organic substances having
       the properties of amides into their corresponding hydroxides.
       Thus aniline, C_{6}H_{5}·NH_{2}, which is obtained from
       nitrobenzene, C_{6}H_{5}·NO_{2} (Note 37), is converted by
       nitrous anhydride into phenol, C_{6}H_{5}·OH, which occurs in the
       creosote extracted from coal tar. Thus the H of the benzene is
       successively replaced by NO_{2}, NH_{2}, and HO; a method which
       is suitable for other cases also.

As nitrous anhydride easily splits up into NO_{2} + NO, so, like NO_{2},
with warm water it gives nitric acid and nitric oxide, according to the
equation 3N_{2}O_{3} + H_{2}O = 4NO + 2NHO_{3}.

Being in a lower degree of oxidation than nitric acid, nitrous
acid and its anhydride are oxidised in solutions by many oxidising
substances--for example, by potassium permanganate--into nitric acid.[57]

  [57] The action of a solution of potassium permanganate, KMnO_{4},
       on nitrous acid in the presence of sulphuric acid is determined
       by the fact that the higher oxide of manganese, Mn_{2}O_{7},
       contained in the permanganate is converted into the lower oxide,
       MnO, which as a base forms manganese sulphate, MnSO_{4}, and
       the oxygen serves for the oxidation of the N_{2}O_{3} into
       N_{2}O_{5}, or its hydrate. As the solution of the permanganate
       is of a red colour, whilst that of manganese sulphate is almost
       colourless, this reaction is clearly seen, and may be employed
       for the detection and determination of nitrous acid and its salts.

_Nitric oxide_, NO.--This permanent gas[58] (that is, unliquefiable
by pressure without the aid of cold) may be obtained from all the
above-described compounds of nitrogen with oxygen. The deoxidation of
nitric acid by metals is the usual method employed for its preparation.
Dilute nitric acid (sp. gr. 1·18, but not stronger, as then N_{2}O_{3}
and NO_{2} are produced) is poured into a flask containing metallic
copper.[59] The reaction commences at the ordinary temperature. Mercury
and silver also give nitric oxide with nitric acid. In these reactions
with metals one portion of the nitric acid is employed in the oxidation
of the metal, whilst the other, and by far the greater, portion combines
with the metallic oxide so obtained, with formation of the nitrate
corresponding with the metal taken. The first action of the copper on
the nitric acid is thus expressed by the equation

                  2NHO_{3} + 3Cu = H_{2}O + 3CuO + 2NO.

The second reaction consists in the formation of copper nitrate--

              6NHO_{3} + 3CuO = 3H_{2}O + 3Cu(NO_{3})_{2}.

  [58] The absolute boiling point = -93° (_see_ Chapter II., Note 29).

  [59] Kammerer proposed preparing nitric oxide, NO, by pouring a solution
       of sodium nitrate over copper shavings, and adding sulphuric
       acid drop by drop. The oxidation of ferrous salts by nitric acid
       also gives NO. One part of strong hydrochloric acid is taken and
       iron is dissolved in it (FeCl_{2}), and then an equal quantity
       of hydrochloric acid and nitre is added to the solution. On
       heating, nitric oxide is evolved. In the presence of an excess
       of sulphuric acid and mercury the conversion of nitric acid into
       nitric oxide is complete (that is, the reaction proceeds to the
       end and the nitric oxide is obtained without other products), and
       upon this is founded one of the methods for determining nitric
       acid (in nitrometers of various kinds, described in text-books
       of analytical chemistry), as the amount of NO can be easily
       and accurately measured volumetrically. The amount of nitrogen
       in gun-cotton, for instance, is determined by dissolving it in
       sulphuric acid. Nitrous acid acts in the same manner. Upon this
       property Emich (1892) founds his method for preparing pure NO.
       He pours mercury into a flask, and then covers it with sulphuric
       acid, in which a certain amount of NaNO_{2} or other substance
       corresponding to HNO_{2} or HNO_{3} has been dissolved. The
       evolution of NO proceeds at the ordinary temperature, being more
       rapid as the surface of the mercury is increased (if shaken, the
       reaction proceeds very rapidly). If the gas be passed over KHO,
       it is obtained quite pure, because KHO does not act upon NO at
       the ordinary temperature (if heated, KNO_{2} and N_{2}O or N_{2},
       are formed).

Nitric oxide is a colourless gas which is only slightly soluble in
water (1/20 of a volume at the ordinary temperature). Reactions of
double decomposition in which nitric oxide readily takes part are not
known--that is to say, it is an indifferent, not a saline, oxide.
Like the other oxides of nitrogen, it is decomposed into its elements
at a red heat (starting from 900°, at 1,200° 60 per cent. give N_{2}
and 2N_{2}O_{3}, but complete decomposition into N_{2} and O_{2} only
takes place at the melting point of platinum, Emich 1892). The most
characteristic property of nitric oxide is its capacity for directly
and easily combining with oxygen (owing to the evolution of heat in
the combination). With oxygen it forms nitrous anhydride and nitrogen
peroxide, 2NO + O = N_{2}O_{3}, 2NO + O_{2} = 2NO_{2}. If nitric oxide
is mixed with oxygen and immediately shaken up with caustic potash,
it is almost entirely converted into potassium nitrite; whilst after
a certain time, when the formation of nitric peroxide has already
commenced, a mixture of potassium nitrite and nitrate is obtained.
If oxygen is passed into a bell jar filled with nitric oxide, brown
fumes of nitrous anhydride and nitric peroxide are formed, even in the
absence of moisture; these in the presence of water give, as we already
know, nitric acid and nitric oxide, so that in the presence of an
excess of water and oxygen the whole of the nitric oxide is easily and
directly converted into nitric acid. This reaction of the re-formation
of nitric acid from nitric oxide, air, and water, 2NO + H_{2}O + O_{3}
= 2HNO_{3}, is frequently made use of in practice. The experiment
showing the conversion of nitric oxide into nitric acid is very striking
and instructive. As the intermixture of the oxygen with the oxide of
nitrogen proceeds, the nitric acid formed dissolves in water, and if an
excess of oxygen has not been added the whole of the gas (nitric oxide),
being converted into HNO_{3}, is absorbed, and the water entirely fills
the bell jar previously containing the gas.[60] It is evident that
nitric oxide[61] in combining with oxygen has a strong tendency to give
the higher types of nitrogen compounds, which we see in nitric acid,
HNO_{3} or NO_{2}(OH), in nitric anhydride, N_{2}O_{5} or (NO_{2})_{2}O,
and in ammonium chloride, NH_{4}Cl. If X stand for an atom of hydrogen,
or its equivalents, chlorine, hydroxyl, &c., and if O, which is,
according to the law of substitution, equivalent to H_{2}, be indicated
by X_{2}, then the three compounds of nitrogen above named should be
considered as compounds of the type or form NX_{5}. For example, in
nitric acid X_{5} = O_{2} + (OH), where O_{2} = X_{4}, and OH = X;
whilst nitric oxide is a compound of the form NX_{2}. Hence this lower
form, like lower forms in general, strives by combination to attain to
the higher forms proper to the compounds of a given element. NX_{2}
passes consecutively into NX_{3}--namely, into N_{2}O_{3} and NHO_{2},
NX_{4} (for instance NO_{2}) and NX_{5}.

  [60] This transformation of the permanent gases nitric oxide and oxygen
       into liquid nitric acid in the presence of water, and with
       the evolution of heat, presents a most striking instance
       of liquefaction produced by the action of chemical forces.
       They perform with ease the work which physical (cooling) and
       mechanical (pressure) forces effect with difficulty. In this the
       motion, which is so distinctively the property of the gaseous
       molecules, is apparently destroyed. In other cases of chemical
       action it is apparently created, arising, no doubt, from latent
       energy--that is, from the internal motion of the atoms in the
       molecules.

  [61] Nitric oxide is capable of entering into many characteristic
       combinations; it is absorbed by the solutions of many acids, for
       instance, tartaric, acetic, phosphoric, sulphuric, and metallic
       chlorides (for example, SbCl_{5}, BiCl_{3}, &c., with which it
       forms definite compounds; Besson 1889), and also by the solutions
       of many salts, especially those formed by suboxide of iron (for
       instance, ferrous sulphate). In this case a brown compound is
       formed which is exceedingly unstable, like all the analogous
       compounds of nitric oxide. The amount of nitric oxide combined
       in this manner is in atomic proportion with the amount of the
       substance taken; thus ferrous sulphate, FeSO_{4}, absorbs it in
       the proportion of NO to 2FeSO_{4}. Ammonia is obtained by the
       action of a caustic alkali on the resultant compound, because
       the oxygen of the nitric oxide and water are transferred to the
       ferrous oxide, forming ferric oxide, whilst the nitrogen combines
       with the hydrogen of the water. According to the investigations
       of Gay (1885), the compound is formed with the evolution of
       a large quantity of heat, and is easily dissociated, like a
       solution of ammonia in water. It is evident that oxidising
       substances (for example, potassium permanganate, KMnO_{4}, Note
       57) are able to convert it into nitric acid. If the presence
       of a radicle NO_{2}, composed like nitrogen peroxide, must be
       recognised in the compounds of nitric acid, then a radicle NO,
       having the composition of nitric oxide, may be admitted in the
       compounds of nitrous acid. The compounds in which the radicle NO
       is recognised are called _nitroso-compounds_. These substances
       are described in Prof. Bunge's work (Kief, 1868).

As the decomposition of nitric oxide begins at temperatures above 900°,
many substances burn in it; thus, ignited phosphorus continues to burn
in nitric oxide, but sulphur and charcoal are extinguished in it. This
is due to the fact that the heat evolved in the combustion of these two
substances is insufficient for the decomposition of the nitric oxide,
whilst the heat developed by burning phosphorus suffices to produce this
decomposition. That nitric oxide really supports combustion, owing to
its being decomposed by the action of heat, is proved by the fact that
strongly ignited charcoal continues to burn in the same nitric oxide[62]
in which a feebly incandescent piece of charcoal is extinguished.

  [62] A mixture of nitric oxide and hydrogen is inflammable. If a
       mixture of the two gases be passed over spongy platinum the
       nitrogen and hydrogen even combine, forming ammonia. A mixture
       of nitric oxide with many combustible vapours and gases is very
       inflammable. A very characteristic flame is obtained in burning a
       mixture of nitric oxide and the vapour of the combustible carbon
       bisulphide, CS_{2}. The latter substance is very volatile, so
       that it is sufficient to pass the nitric oxide through a layer
       of the carbon bisulphide (for instance, in a Woulfe's bottle)
       in order that the gas escaping should contain a considerable
       amount of the vapours of this substance. This mixture continues
       to burn when ignited, and the flame emits a large quantity of
       the so-called ultra-violet rays, which are capable of inducing
       chemical combinations and decompositions, and therefore the flame
       may be employed in photography in the absence of sufficient
       daylight (magnesium light and electric light have the same
       property). There are many gases (for instance, ammonia) which
       when mixed with nitric oxide explode in a eudiometer.

The compounds of nitrogen with oxygen which we have so far considered
may all be prepared from nitric oxide, and may themselves be converted
into it. Thus nitric oxide stands in intimate connection with them.[63]
The passage of nitric oxide into the higher degrees of oxidation and the
converse reaction is employed in practice as a means for _transferring_
the oxygen of the air to substances capable of being oxidised. Starting
with nitric oxide, it may easily be converted, with the aid of the
oxygen of the atmosphere and water, into nitric acid, nitrous anhydride,
and nitric peroxide, and by their means employed to oxidise other
substances. In this oxidising action nitric oxide is again formed, and
it may again be converted into nitric acid, and so on continuously,
if only oxygen and water be present. Hence the fact, which at first
appears to be a paradox, that by means of a small quantity of nitric
oxide in the presence of oxygen and water it is possible to oxidise
an indefinitely large quantity of substances which cannot be directly
oxidised either by the action of the atmospheric oxygen or by the
action of nitric oxide itself. The sulphurous anhydride, SO_{2}, which
is obtained in the combustion of sulphur and in roasting many metallic
sulphides in the air is an example of this kind. In practice this gas is
obtained by burning sulphur or iron pyrites, the latter being thereby
converted into oxide of iron and sulphurous anhydride. In contact with
the oxygen of the atmosphere this gas does not pass into the higher
degree of oxidation, sulphuric anhydride, SO_{3}, and if it does form
sulphuric acid with water and the oxygen of the atmosphere, SO_{2} +
H_{2}O + O = H_{2}SO_{4}, it does so very slowly. With nitric acid (and
especially with nitrous acid, but not with nitrogen peroxide) and water,
sulphurous anhydride, on the contrary, very easily forms sulphuric acid,
and especially so when slightly heated (about 40°), the nitric acid (or,
better still, nitrous acid) being converted into nitric oxide--

           3SO_{2} + 2NHO_{3} + 2H_{2}O = 2H_{2}SO_{4} + 2NO.

  [63] The oxides of nitrogen naturally do not proceed directly from
       oxygen and nitrogen by contact alone, because their formation
       is accompanied by the absorption of a large quantity of heat,
       for (_see_ Note 29) about 21,500 heat units are absorbed when 16
       parts of oxygen and 14 parts of nitrogen combine; consequently
       the decomposition of nitric oxide into oxygen and nitrogen
       is accompanied by the evolution of this amount of heat; and
       therefore with nitric oxide, as with all explosive substances and
       mixtures, the reaction once started is able to proceed by itself.
       In fact, Berthelot remarked the decomposition of nitric oxide
       in the explosion of fulminate of mercury. This decomposition
       does not take place spontaneously; substances even burn with
       difficulty in nitric oxide, probably because a certain portion of
       the nitric oxide in decomposing gives oxygen, which combines with
       another portion of nitric oxide, and forms nitric peroxide, a
       somewhat more stable compound of nitrogen and oxygen. The further
       combinations of nitric oxide with oxygen all proceed with the
       evolution of heat, and take place spontaneously by contact with
       air alone. It is evident from these examples that the application
       of thermochemical data is limited.

The presence of water is absolutely indispensable here, otherwise
sulphuric anhydride is formed, which combines with the oxides of
nitrogen (nitrous anhydride), forming a crystalline substance containing
oxides of nitrogen (_chamber crystals_, which will be described in
Chapter XX.) Water destroys this compound, forming sulphuric acid
and separating the oxides of nitrogen. The water must be taken in a
greater quantity than that required for the formation of the hydrate
H_{2}SO_{4}, because the latter absorbs oxides of nitrogen. With an
excess of water, however, solution does not take place. If, in the above
reaction, only water, sulphurous anhydride, and nitric or nitrous acid
be taken in a definite quantity, then a definite quantity of sulphuric
acid and nitric oxide will be formed, according to the preceding
equation; but there the reaction ends and the excess of sulphurous
anhydride, if there be any, will remain unchanged. But if we add air and
water, then the nitric oxide will unite with the oxygen to form nitrogen
peroxide, and the latter with water to form nitric and nitrous acids,
which again give sulphuric acid from a fresh quantity of sulphurous
anhydride. Nitric oxide is again formed, which is able to start the
oxidation afresh if there be sufficient air. Thus it is possible
with a definite quantity of nitric oxide to convert an indefinitely
large quantity of sulphurous anhydride into sulphuric acid, water and
oxygen only being required.[64] This may be easily demonstrated by an
experiment on a small scale, if a certain quantity of nitric oxide be
first introduced into a flask, and sulphurous anhydride, steam, and
oxygen be then continually passed in. Thus the above-described reaction
may be expressed in the following manner:--

  _n_SO_{2} + _n_O + (_n_ + _m_)H_{2}O + NO
                                          = _n_H_{2}SO_{4},_m_H_{2}O + NO

if we consider only the original substances and those finally formed.
In this way a definite quantity of nitric oxide may serve for the
conversion of an indefinite quantity of sulphurous anhydride, oxygen,
and water into sulphuric acid. In reality, however, there is a limit to
this, because air, and not pure oxygen, is employed for the oxidation,
so that it is necessary to remove the nitrogen of the air and to
introduce a fresh quantity of air. A certain quantity of nitric oxide
will pass away with this nitrogen, and will in this way be lost.[65]

  [64] The instance of the action of a small quantity of NO in inducing
       a definite chemical reaction between large masses (SO_{2} +
       O + H_{2}O = H_{2}SO_{4}) is very instructive, because the
       particulars relating to it have been studied, and show that
       intermediate forms of reaction may be discovered in the so-called
       contact or catalytic phenomena. The essence of the matter here is
       that A (= SO_{2}) reacts upon B (= O and H_{2}O) in the presence
       of C, because it gives BC, a substance which forms AB with A, and
       again liberates C. Consequently C is a medium, a transferring
       substance, without which the reaction does not proceed. Many
       similar phenomena may be found in other departments of life. Thus
       the merchant is an indispensable medium between the producer
       and the consumer; experiment is a medium between the phenomena
       of nature and the cognisant faculties, and language, customs,
       and laws are media which are as necessary for the exchanges of
       social intercourse as nitric oxide for those between sulphurous
       anhydride and oxygen and water.

  [65] If the sulphurous anhydride be prepared by roasting iron pyrites,
       FeS_{2}, then each equivalent of pyrites (equivalent of iron,
       56, of sulphur 32, of pyrites 120) requires six equivalents of
       oxygen (that is 96 parts) for the conversion of its sulphur
       into sulphuric acid (for forming 2H_{2}SO_{4} with water),
       besides 1-1/2 equivalents (24 parts) for converting the iron
       into oxide, Fe_{2}O_{3}; hence the combustion of the pyrites for
       the formation of sulphuric acid and ferric oxide requires the
       introduction of an equal weight of oxygen (120 parts of oxygen to
       120 parts of pyrites), or five times its weight of air, whilst
       four parts by weight of nitrogen will remain inactive, and in the
       removal of the exhausted air will carry off the remaining nitric
       oxide. If not all, at least a large portion of the nitric oxide
       may be collected by passing the escaping air, still containing
       some oxygen, through substances which absorb oxides of nitrogen.
       Sulphuric acid itself may be employed for this purpose if it be
       used in the form of the hydrate H_{2}SO_{4}, or containing only
       a small amount of water, because such sulphuric acid dissolves
       the oxides of nitrogen. They may be easily expelled from this
       solution by heating or by dilution with water, as they are only
       slightly soluble in aqueous sulphuric acid. Besides which,
       sulphurous anhydride acts on such sulphuric acid, being oxidised
       at the expense of the nitrous anhydride, and forming nitric
       oxide from it, which again enters into the cycle of action. For
       this reason the sulphuric acid which has absorbed the oxides of
       nitrogen escaping from the chambers in the tower K (_see_ fig.
       50) is led back into the first chamber, where it comes into
       contact with sulphurous anhydride, by which means the oxides of
       nitrogen are reintroduced into the reaction which proceeds in
       the chambers. This is the use of the towers (Gay-Lussac's and
       Glover's) which are erected at either end of the chambers.

The preceding series of changes serve as the basis of the _manufacture
of sulphuric acid_ or so-called _chamber acid_. This acid is prepared
on a very large scale in chemical works because it is the cheapest acid
whose action can be applied in a great number of cases. It is therefore
used in immense quantities.

[Illustration: FIG. 50.--Section of sulphuric acid chambers, the first
and last chambers only being represented. The tower to the left is
called the Glover's tower, and that on the right the Gay-Lussac's tower.
Less than 1/10th of the natural size.]

The process is carried on in a series of chambers (or in one divided
by partitions as in fig. 50, which shows the beginning and end of a
chamber) constructed of sheet lead. These chambers are placed one
against the other, and communicate by tubes or special orifices so
placed that the inlet tubes are in the upper portion of the chamber, and
the outlet in the lower and opposite end. The current of steam and gases
necessary for the preparation of the sulphuric acid passes through these
chambers and tubes. The acid as it is formed falls to the bottom of the
chambers or runs down their walls, and flows from chamber to chamber
(from the last towards the first), to permit of which the partitions
do not reach to the bottom. The floor and walls of the chambers should
therefore be made of a material on which the sulphuric acid will not
act. Among the ordinary metals lead is the only one suitable.[65 bis]

  [65 bis] Other metals, iron, copper, zinc, are corroded by it;
       glass and china are not acted upon, but they crack from the
       variations of temperature taking place in the chambers, and
       besides they are more difficult to join properly than lead; wood,
       &c., becomes charred.

For the formation of the sulphuric acid it is necessary to introduce
sulphurous anhydride, steam, air, and nitric acid, or some oxide of
nitrogen, into the chambers. The sulphurous anhydride is produced by
burning sulphur or iron pyrites. This is carried on in the furnace with
four hearths to the left of the drawing. Air is led into the chambers
and furnace through orifices in the furnace doors. The current of air
and oxygen is regulated by opening or closing these orifices to a
greater or less extent. The ingoing draught in the chambers is brought
about by the fact that heated gases and vapours pass into the chambers,
whose temperature is further raised by the reaction itself, and also
by the remaining nitrogen being continually withdrawn from the outlet
(above the tower K) by a tall chimney situated near the chambers. Nitric
acid is prepared from a mixture of sulphuric acid and Chili saltpetre,
in the same furnaces in which the sulphurous anhydride is evolved (or in
special furnaces). Not more than 8 parts of nitre are taken to 100 parts
of sulphur burnt. On leaving the furnace the vapours of nitric acid and
oxides of nitrogen mixed with air and sulphurous anhydride first pass
along the horizontal tubes T into the receiver B B, which is partially
cooled by water flowing in on the right-hand side and running out on the
left by _o_, in order to reduce the temperature of the gases entering
the chamber. The gases then pass up a tower filled with coke, and shown
to the left of the drawing. In this tower are placed lumps of coke (the
residue from the dry distillation of coal), over which sulphuric acid
trickles from the reservoir M. This acid has absorbed in the end tower K
the oxides of nitrogen escaping from the chamber. This end tower is also
filled with coke, over which a stream of strong sulphuric acid trickles
from the reservoir M. The acid spreads over the coke, and, owing to the
large surface offered by the latter, absorbs the greater part of the
oxides of nitrogen escaping from the chambers. The sulphuric acid in
passing down the tower becomes saturated with the oxides of nitrogen,
and flows out at _h_ into a special receiver (in the drawing situated
by the side of the furnaces), from which it is forced up the tubes _h´
h´_ by steam pressure into the reservoir M, situated above the first
tower. The gases passing through this tower (hot) from the furnace
on coming into contact with the sulphuric acid take up the oxides of
nitrogen contained in it, and these are thus returned to the chamber and
again participate in the reaction. The sulphuric acid left after their
extraction flows into the chambers. Thus, on leaving the first coke
tower the sulphurous anhydride, air, and vapours of nitric acid and of
the oxides of nitrogen pass through the upper tube _m_ into the chamber.
Here they come into contact with steam introduced by lead tubes into
various parts of the chamber. The reaction takes place in the presence
of water, the sulphuric acid falls to the bottom of the chamber, and the
same process takes place in the following chambers until the whole of
the sulphurous anhydride is consumed. A somewhat greater proportion of
air than is strictly necessary is passed in, in order that no sulphurous
anhydride should be left unaltered for want of sufficient oxygen. The
presence of an excess of oxygen is shown by the colour of the gases
escaping from the last chamber. If they be of a pale colour it indicates
an insufficiency of air (and the presence of sulphurous anhydride), as
otherwise peroxide of nitrogen would be formed. A very dark colour shows
an excess of air, which is also disadvantageous, because it increases
the inevitable loss of nitric oxide by increasing the mass of escaping
gases.[66]

  [66] By this means as much as 2,500,000 kilograms of chamber acid,
       containing about 60 per cent. of the hydrate H_{2}SO_{4} and
       about 40 per cent. of water, may be manufactured per year in one
       plant of 5,000 cubic metres capacity (without stoppages). This
       process has been brought to such a degree of perfection that
       as much as 300 parts of the hydrate H_{2}SO_{4} are obtained
       from 100 parts of sulphur, whilst the theoretical amount is
       not greater than 306 parts. The acid parts with its excess of
       water on heating. For this purpose it is heated in lead vessels.
       However, the acid containing about 75 per cent. of the hydrate
       (60° Baumé) already begins to act on the lead when heated,
       and therefore the further removal of water is conducted by
       evaporating in glass or platinum vessels, as will he described
       in Chapter XX. The aqueous acid (50° Baumé) obtained in the
       chambers is termed chamber acid. The acid concentrated to 60°
       Baumé is more generally employed, and sometimes the hydrate (66°
       Baumé) termed vitriol acid is also used. In England alone more
       than 1,000 million kilograms of chamber acid are produced by
       this method. The formation of sulphuric acid by the action of
       nitric acid was discovered by Drebbel, and the first lead chamber
       was erected by Roebuck, in Scotland, in the middle of the last
       century. The essence of the process was only brought to light
       at the beginning of this century, when many improvements were
       introduced into practice.

_Nitrous oxide_, N_{2}O,[67] is similar to water in its volumetric
composition. Two volumes of nitrous oxide are formed from two volumes
of nitrogen and one volume of oxygen, which may be shown by the
ordinary method for the analysis of the oxides of nitrogen (by passing
them over red-hot copper or sodium). In contradistinction to the other
oxides of nitrogen, it is not directly oxidised by oxygen, but it
may be obtained from the higher oxides of nitrogen by the action of
certain deoxidising substances; thus, for example, a mixture of two
volumes of nitric oxide and one volume of sulphurous anhydride if left
in contact with water and spongy platinum is converted into sulphuric
acid and nitrous oxide, 2NO + SO_{2} + H_{2}O = H_{2}SO_{4} + N_{2}O.
Nitric acid, also, under the action of certain metals--for instance, of
zinc[68]--gives nitrous oxide, although in this case mixed with nitric
oxide. The usual method of preparing nitrous oxide consists in the
decomposition of ammonium nitrate by the aid of heat, because in this
case only water and nitrous oxide are formed, NH_{4}NO_{3} = 2H_{2}O
+ N_{2}O (a mixture of NH_{4}Cl and KNO_{3} is sometimes taken). The
decomposition[69] proceeds very easily in an apparatus like that used
for the preparation of ammonia or oxygen--that is, in a retort or flask
with a gas-conducting tube. The decomposition must, however, be carried
on carefully, as otherwise nitrogen is formed from the decomposition of
the nitrous oxide.[70]

  [67] If the hydrate HNO_{3} corresponds to N_{2}O_{4}, the hydrate HNO,
       _hyponitrous acid_, corresponds to N_{2}O, and in this
       sense N_{2}O is _hyponitrous anhydride_. Hyponitrous acid,
       corresponding with nitrous oxide (as its anhydride), is not
       known in a pure state, but its salts (Divers) are known. They
       are prepared by the reduction of nitrous (and consequently of
       nitric) salts by sodium amalgam. If this amalgam he added to a
       cold solution of an alkaline nitrite until the evolution of gas
       ceases, and the excess of alkali saturated with acetic acid,
       an insoluble yellow precipitate of silver hyponitrite, NAgO,
       will he obtained on adding a solution of silver nitrate. This
       hyponitrite is insoluble in cold acetic acid, and decomposes when
       heated, with the evolution of nitrous oxide. If rapidly heated it
       decomposes with an explosion. It is dissolved unchanged by weak
       mineral acids, whilst the stronger acids (for example, sulphuric
       and hydrochloric acids) decompose it, with the evolution of
       nitrogen, nitric and nitrous acids remaining in solution. Among
       the other salts of hyponitrous acid, HNO, the salts of lead,
       copper, and mercury are insoluble in water. Judging by the bond
       between hyponitrous acid and the other compounds of nitrogen,
       there is reason for thinking that its formula should he doubled,
       N_{2}H_{2}O_{2}. For instance, Thoune (1893) on gradually
       oxidising hydroxylamine, NH_{2}(OH), into nitrous acid, NO(OH)
       (Note 25), by means of an alkaline solution of KMnO_{4}, first
       obtained hyponitrous acid, N_{2}H_{2}O_{2}, and then a peculiar
       intermediate acid, N_{2}H_{2}O_{3}, which, by further oxidation,
       gave nitrous acid. On the other hand, Wislicenus (1893) showed
       that in the action of the sulphuric acid salt of hydroxylamine
       upon nitrite of sodium, there is formed, besides, nitrous
       oxide (according to V. Meyer, NH_{3}O,H_{2}SO_{4} + NaNO_{2} =
       NaHSO_{4} + 2H_{2}O + N_{2}O), a small amount of hyponitrous
       acid which may be precipitated in the form of the silver salt;
       and this reaction is most simply expressed by taking the doubled
       formula of hyponitrous acid, NH_{2}(OH) + NO(OH) = H_{2}O +
       N_{2}H_{2}O_{2}. The best argument in favour of the doubled
       formula is the property possessed by hyponitrous acid of forming
       acid salts, HNaN_{2}O_{2} (Zorn).

       According to Thoune, the following are the properties of
       hyponitrous acid. When liberated from the dry silver salt by
       the action of dry sulphuretted hydrogen, hyponitrous acid is
       unstable, and easily explodes even at low temperatures. But
       when dissolved in water (having been formed by the action of
       hydrochloric acid upon the silver salt), it is stable even
       when boiled with dilute acids and alkalis. The solution is
       colourless and has a strongly acid reaction. In the course of
       time, however, the aqueous solution also decomposes into nitrous
       oxide and water. The complete oxidation by permanganate of potash
       proceeds according to the following equation: 5H_{2}N_{2}O_{2} +
       8KMnO_{4} + 12H_{2}SO_{4} = 10HNO_{3} + 4K_{2}SO_{4} + 8MnSO_{4}
       + 12H_{2}O. In an alkaline solution, KMnO_{4} only oxidises
       hyponitrous acid into nitrous and not into nitric acid. Nitrous
       acid has a decomposing action upon hyponitrous acid, and if
       the aqueous solutions of the two acids be mixed together they
       immediately give off oxides of nitrogen. Hyponitrous acid does
       not liberate CO_{2} from its salts, but on the other hand it is
       not displaced by CO_{2}.

  [68] It is remarkable that electro-deposited copper powder gives
       nitrous oxide with a 10 p.c. solution of nitric acid, whilst
       ordinary copper gives nitric oxide. It is here evident that
       the physical and mechanical structure of the substance affects
       the course of the reaction--that is to say, it is a case of
       contact-action.

  [69] This decomposition is accompanied by the evolution of about 25,000
       calories per molecular quantity, NH_{4}NO_{3}, and therefore
       takes place with ease, and sometimes with an explosion.

  [70] In order to remove any nitric oxide that might be present, the
       gas obtained is passed through a solution of ferrous sulphate. As
       nitrous oxide is very soluble in cold water (at 0°, 100 volumes
       of water dissolve 130 volumes of N_{2}O; at 20°, 67 volumes),
       it must be collected over warm water. The nitrous oxide is much
       more soluble than nitric oxide, which is in agreement with the
       fact that nitrous oxide is much more easily liquefied than nitric
       oxide. Villard obtained a crystallohydrate, N_{2}O,6H_{2}O, which
       was tolerably stable at 0°.

[Illustration: FIG. 51.--Natterer's apparatus for the preparation of
liquid nitrous oxide and carbonic anhydride. The gas first passes though
the vessel V, for drying, and then into the pump (a section of the upper
part of the apparatus is given on the left). The piston _t_ of the
force pump is moved by the crank E and fly-wheel turned by hand. The
gas is pumped into the iron chamber A, where it is liquefied. The valve
S allows the gas to enter A, but not to escape from it. The chamber
and pump are cooled by the jacket B, filled with ice. When the gas is
liquefied the vessel A is unscrewed from the pump, and the liquid may be
poured from it by inverting it and unscrewing the valve _v_, when the
liquid runs out of the tube _x_.]

Nitrous oxide is not a permanent gas (absolute boiling point +36°); it
is easily liquefied by the action of cold under a high pressure; at
15° it may be liquefied by a pressure of about 40 atmospheres. This
gas is usually liquefied by means of the force pump[71] shown in fig.
51. As it is liquefied with comparative ease, and as the cold produced
by its vaporisation is very considerable,[72] it (as also liquid
carbonic anhydride) is often employed in investigations requiring a low
temperature. Nitrous oxide forms a very mobile, colourless liquid, which
acts on the skin, and is incapable in a cold state of oxidising either
metallic potassium, phosphorus, or carbon; its specific gravity is
slightly less than that of water (0° = 0·910, 10° = 0·856, 35° = 0·60,
39° = 0·45, Villard, 1894). When evaporated under the receiver of an
air-pump, the temperature falls to -100°, and the liquid solidifies into
a snow-like mass, and partially forms transparent crystals. Both these
substances are solid nitrous oxide. Mercury is immediately solidified in
contact with evaporating liquid nitrous oxide.[73]

  [71] Faraday obtained liquid nitrous oxide by the same method as liquid
       ammonia, by beating dry ammonium nitrate in a closed bent
       tube, one arm of which was immersed in a freezing mixture. In
       this case two layers of liquid are obtained at the cooled end,
       a lower layer of water and an upper layer of nitrous oxide.
       This experiment should be conducted with great care, as the
       pressure of the nitrous oxide in a liquid state is considerable,
       namely (according to Regnault), at +10° = 45 atmospheres, at
       0° = 36 atmospheres, at -10° = 29 atmospheres, and at -20° = 23
       atmospheres. It boils at -92°, and the pressure is then therefore
       = 1 atmosphere (_see_ Chapter II., Note 27).

  [72] Liquid nitrous oxide, in vaporising at the same pressure as liquid
       carbonic anhydride, gives rise to almost equal or even slightly
       lower temperatures. Thus at a pressure of 25 mm. carbonic
       anhydride gives a temperature as low as -115°, and nitrous oxide
       of -125° (Dewar). The similarity of these properties and even of
       the absolute boiling point (CO_{2} + 32°, N_{2}O +36°) is all
       the more remarkable because these gases have the same molecular
       weight = 44 (Chapter VII.)

  [73] A very characteristic experiment of simultaneous combustion and
       intense cold may be performed by means of liquid nitrous oxide;
       if liquid nitrous oxide be poured into a test tube containing
       some mercury the mercury will solidify, and if a piece of red-hot
       charcoal be thrown upon the surface of the nitrous oxide it
       will continue to burn very brilliantly, giving rise to a high
       temperature.

When introduced into the respiratory organs (and consequently into the
blood also) nitrous oxide produces a peculiar kind of intoxication
accompanied by spasmodic movements, and hence this gas, discovered by
Priestley in 1776, received the name of 'laughing gas.' On a prolonged
respiration it produces a state of insensibility (it is an anæsthetic
like chloroform), and is therefore employed in dental and surgical
operations.

Nitrous oxide is easily decomposed into nitrogen and oxygen by the
action of heat, or a series of electric sparks; and this explains why a
number of substances which cannot burn in nitric oxide do so with great
ease in nitrous oxide. In fact, when nitric oxide gives some oxygen on
decomposition, this oxygen immediately unites with a fresh portion of
the gas to form nitric peroxide, whilst nitrous oxide does not possess
this capacity for further combination with oxygen.[74] A mixture of
nitrous oxide with hydrogen explodes like detonating gas, gaseous
nitrogen being formed, N_{2}O + H_{2} = H_{2}O + N_{2}. The volume
of the remaining nitrogen is equal to the original volume of nitrous
oxide, and is equal to the volume of hydrogen entering into combination
with the oxygen; hence in this reaction equal volumes of nitrogen
and hydrogen replace each other. Nitrous oxide is also very easily
decomposed by red-hot metals; and sulphur, phosphorus, and charcoal
burn in it, although not so brilliantly as in oxygen. A substance in
burning in nitrous oxide evolves more heat than an equal quantity
burning in oxygen; which most clearly shows that in the formation of
nitrous oxide by the combination of nitrogen with oxygen there was not
an evolution but an absorption of heat, there being no other source for
the excess of heat in the combustion of substances in nitrous oxide
(_see_ Note 29). If a given volume of nitrous oxide be decomposed by
a metal--for instance, sodium--then there remains, after cooling and
total decomposition, a volume of nitrogen, exactly equal to that of the
nitrous oxide taken; consequently the oxygen is, so to say, distributed
between the atoms of nitrogen without producing an increase in the
volume of the nitrogen.

  [74] In the following chapter we shall consider the volumetric
       composition of the oxides of nitrogen. It explains the difference
       between nitric and nitrous oxide. Nitrous oxide is formed with
       a diminution of volumes (contraction), nitric oxide without
       contraction, its volume being equal to the sum of the volumes of
       the nitrogen and oxygen of which it is composed. By oxidation, if
       it could be directly accomplished, two volumes of nitrous oxide
       and one volume of oxygen would not give three but four volumes
       of nitric oxide. These facts must be taken into consideration in
       comparing the calorific equivalents of formation, the capacity
       for supporting combustion, and other properties of nitrous and
       nitric oxides, N_{2}O and NO.




                               CHAPTER VII

    MOLECULES AND ATOMS. THE LAWS OF GAY-LUSSAC AND AVOGADRO-GERHARDT


Hydrogen combines with oxygen in the proportion of two volumes to one.
The composition by volume of nitrous oxide is exactly similar--it
is composed of two volumes of nitrogen and one volume of oxygen. By
decomposing ammonia by the action of an electric spark it is easy to
prove that it contains one volume of nitrogen to three volumes of
hydrogen. So, similarly, it is found, whenever a compound is decomposed
and the volumes of the gases proceeding from it are measured, that
the volumes of the gases or vapours entering into combination are in
a very simple proportion to one another. With water, nitrous oxide,
&c., this may be proved by direct observation; but in the majority of
cases, and especially with substances which, although volatile--that is,
capable of passing into a gaseous (or vaporous) state--are liquid at
the ordinary temperature, such a direct method of observation presents
many difficulties. But, then, if the densities of the vapours and gases
be known, the same simplicity in their ratio is shown by calculation.
The volume of a substance is proportional to its weight, and inversely
proportional to its density, and therefore by dividing the amount by
weight of each substance entering into the composition of a compound by
its density in the gaseous or vaporous state we shall obtain factors
which will be in the same proportion as the volumes of the substances
entering into the composition of the compound.[1] So, for example,
water contains eight parts by weight of oxygen to one part by weight
of hydrogen, and their densities are 16 and 1, consequently their
volumes (or the above-mentioned factors) are 1 and 1/2, and therefore
it is seen without direct experiment that water contains two volumes of
hydrogen for every one volume of oxygen. So also, knowing that nitric
oxide contains fourteen parts of nitrogen and sixteen parts of oxygen,
and knowing that the specific gravities of these last two gases are
fourteen and sixteen, we find that the volumes in which nitrogen and
oxygen combine for the formation of nitric oxide are in the proportion
of 1 : 1. We will cite another example. In the last chapter we saw that
the density of NO_{2} only becomes constant and equal to twenty-three
(referred to hydrogen) above 135°, and as a matter of fact a method
of direct observation of the volumetric composition of this substance
would be very difficult at so high a temperature. But it may be easily
calculated. NO_{2}, as is seen from its formula and analysis, contains
thirty-two parts by weight of oxygen to fourteen parts by weight of
nitrogen, forming forty-six parts by weight of NO_{2}, and knowing the
densities of these gases we find that one volume of nitrogen with two
volumes of oxygen gives two volumes of nitrogen peroxide. Therefore,
knowing the amounts by weight of the substances participating in a
reaction or forming a given substance, and knowing the density of the
gas or vapour,[2] the volumetric relations of the substances acting in
a reaction or entering into the composition of a compound, may be also
determined.

  [1] If the weight be indicated by P, the density by D, and the volume
      by V, then

                               P/D = _K_V

      where _K_ is a coefficient depending on the system of the
      expressions P, D, and V. If D be the weight of a cubic measure
      of a substance referred to the weight of the same measure of
      water--if, as in the metrical system (Chapter I., Note 9), the
      cubic measure of one part by weight of water be taken as a unit
      of volume--then _K_ = 1. But, whatever it be, it is cancelled in
      dealing with the comparison of volumes, because comparative and
      not absolute measures of volumes are taken. In this chapter, as
      throughout the book, the weight P is given in grams in dealing
      with absolute weights; and if comparative, as in the expression
      of chemical composition, then the weight of an atom is taken as
      unity. The density of gases, D, is also taken in reference to the
      density of hydrogen, and the volume V in metrical units (cubic
      centimetres), if it be a matter of absolute magnitudes of volumes,
      and if it be a matter of chemical transformations--that is, of
      relative volumes--then the volume of an atom of hydrogen, or of
      one part by weight of hydrogen, is taken as unity, and all volumes
      are expressed according to these units.

  [2] As the volumetric relations of vapours and gases, next to the
      relations of substances by weight, form the most important
      province of chemistry, and a most important means for the
      attainment of chemical conclusions, and inasmuch as these
      volumetric relations are determined by the densities of gases and
      vapours, necessarily the methods of determining the densities of
      vapours (and also of gases) are important factors in chemical
      research. These methods are described in detail in works on
      physics and physical and analytical chemistry, and therefore we
      here only touch on the general principles of the subject.

      If we know the weight _p_ and volume _v_, occupied by the vapour
      of a given substance at a temperature _t_ and pressure _h_, then
      its density may be directly obtained by dividing _p_ by the weight
      of a volume _v_ of hydrogen (if the density be expressed according
      to hydrogen, _see_ Chapter II., Note 23) at _t_ and _h_. Hence,
      the methods of determining the density of vapours and gases are
      based on the determination of _p_, _v_, _t_, and _h_. The two
      last data (the temperature _t_ and pressure _h_) are given by the
      thermometer and barometer and the heights of mercury or other
      liquid confining the gas, and therefore do not require further
      explanation. It need only be remarked that: (1) In the case of
      easily volatile liquids there is no difficulty in procuring a
      bath with a constant temperature, but that it is nevertheless
      best (especially considering the inaccuracy of thermometers) to
      have a medium of absolutely constant temperature, and therefore
      to take either a bath in which some substance is melting--such
      as melting ice at 0° or crystals of sodium acetate, melting at
      +56°--or, as is more generally practised, to place the vessel
      containing the substance to be experimented with in the vapour
      of a liquid boiling at a definite temperature, and knowing the
      pressure under which it is boiling, to determine the temperature
      of the vapour. For this purpose the boiling points of water at
      different pressures are given in Chapter I., Note 11, and the
      boiling points of certain easily procurable liquids at various
      pressures are given in Chapter II., Note 27. (2) With respect to
      temperatures above 300° (below which mercurial thermometers may
      be conveniently employed), they are most simply obtained constant
      (to give time for the weight and volume of a substance being
      observed in a given space, and to allow that space to attain the
      calculated temperature _t_) by means of substances boiling at a
      high temperature. Thus, for instance, at the ordinary atmospheric
      pressure the temperature _t_ of the vapour of sulphur is about
      445°, of phosphorus pentasulphide 518°, of tin chloride 606°,
      of cadmium 770°, of zinc 930° (according to Violle and others),
      or 1040° (according to Deville), &c. (3) The indications of the
      hydrogen thermometer must be considered as the most exact (but
      as hydrogen diffuses through incandescent platinum, nitrogen is
      usually employed). (4) The temperature of the vapours used as
      the bath should in every case be several degrees higher than the
      boiling point of the liquid whose density is to be determined,
      in order that no portion should remain in a liquid state. But
      even in this case, as is seen from the example of nitric peroxide
      (Chapter VI.), the vapour density does not always remain constant
      with a change of _t_, as it should were the law of the expansion
      of gases and vapours absolutely exact (Chapter II., Note 26). If
      variations of a chemical and physical nature similar to that which
      we saw in nitric peroxide take place in the vapours, the main
      interest is centred in _constant_ densities, which do not vary
      with _t_, and therefore the possible effect of _t_ on the density
      must always be kept in mind in having recourse to this means
      of investigation. (5) Usually, for the sake of convenience of
      observation, the vapour density is determined at the atmospheric
      pressure which is read on the barometer; but in the case of
      substances which are volatilised with difficulty, and also of
      substances which decompose, or, in general, vary at temperatures
      near their boiling points, it is best or even indispensable to
      conduct the determination at low pressures, whilst for substances
      which decompose at low pressures the observations have to be
      conducted under a more or less considerably increased pressure.
      (6) In many cases it is convenient to determine the vapour density
      of a substance in admixture with other gases, and consequently
      under the partial pressure, which may be calculated from the
      volume of the mixture and that of the intermixed gas (_see_
      Chapter I., Note 1). This method is especially important for
      substances which are easily decomposable, because, as shown by the
      phenomena of dissociation, a substance is able to remain unchanged
      in the atmosphere of one of its products of decomposition. Thus,
      Wurtz determined the density of phosphoric chloride, PCl_{5},
      in admixture with the vapour of phosphorous chloride, PCl_{3}.
      (7) It is evident, from the example of nitric peroxide, that a
      change of pressure may alter the density and aid decomposition,
      and therefore identical results are sometimes obtained (if the
      density be variable) by raising _t_ and lowering _h_; but if the
      density does not vary under these variable conditions (at least,
      to an extent appreciably exceeding the limits of experimental
      error), then this _constant_ density indicates the _gaseous_
      and _invariable_ state of a substance. The laws hereafter laid
      down refer only to such vapour densities. But the majority of
      volatile substances show such a constant density at a certain
      degree above their boiling points up to the starting point of
      decomposition. Thus, the density of aqueous vapour does not vary
      for _t_ between the ordinary temperature and 1000° (there are no
      trustworthy determinations beyond this) and for pressures varying
      from fractions of an atmosphere up to several atmospheres. If,
      however, the density does vary considerably with a variation of
      _h_ and _t_, the fact may serve as a guide for the investigation
      of the chemical changes which are undergone by the substance in a
      state of vapour, or at least as an indication of a deviation from
      the laws of Boyle, Mariotte, and Gay-Lussac (for the expansion of
      gases with _t_). In certain cases the separation of one form of
      deviation from the other may be explained by special hypotheses.

      With respect to the means of determining _p_ and _v_, with a
      view to finding the vapour density, we may distinguish three
      chief methods: (_a_) by weight, by ascertaining the weight of
      a definite volume of vapour; (_b_) by volume, by measuring the
      volume occupied by the vapour of a definite weight of a substance;
      and (_c_) by displacement. The last-mentioned is essentially
      volumetric, because a known weight of a substance is taken, and
      the volume of the air displaced by the vapour at a given _t_ and
      _h_ is determined.

      [Illustration: FIG. 52.--Apparatus for determining the vapour
      density by Dumas' method. A small quantity of the liquid whose
      vapour density is to be determined is placed in the glass
      globe, and heated in a water or oil bath to a temperature above
      the boiling point of the liquid. When all the liquid has been
      converted into vapour and has displaced all the air from the
      globe, the latter is sealed up and weighed. The capacity of the
      globe is then measured, and in this manner the volume occupied by
      a known weight of vapour at a known temperature is determined.]

      [Illustration: FIG. 53.--Deville and Troost's apparatus for
      determining the vapour densities, according to Dumas' method, of
      substances which boil at high temperatures. A porcelain globe
      containing the substance whose vapour density is to be determined
      is heated in the vapour of mercury (350°), sulphur (410°), cadmium
      (850°), or zinc (1,040°). The globe is sealed up in an oxyhydrogen
      flame.]

      The method by weight (_a_) is the most trustworthy and
      historically important. _Dumas' method_ is typical. An ordinary
      spherical glass or porcelain vessel, like those shown respectively
      in figs. 52 and 53, is taken, and an excess of the substance to be
      experimented upon is introduced into it. The vessel is heated to a
      temperature _t_ higher than the boiling point of the liquid: this
      gives a vapour which displaces the air, and fills the spherical
      space. When the air and vapour cease escaping from the sphere, it
      is fused up or closed by some means; and when cool, the weight
      of the vapour remaining in the sphere is determined (either by
      direct weighing of the vessel with the vapour and introducing the
      necessary corrections for the weight of the air and of the vapour
      itself, or the weight of the volatilised substance is determined
      by chemical methods), and the volume of the vapour at _t_ and the
      barometric pressure _h_ are then calculated.

      _The volumetric method_ (_b_) originally employed by Gay-Lussac
      and then modified by Hofmann and others is based on the principle
      that a weighed quantity of the liquid to be experimented with
      (placed in a small closed vessel, which is sometimes fused up
      before weighing, and, if quite full of the liquid, breaks when
      heated in a vacuum) is introduced into a graduated cylinder heated
      to _t_, or simply into a Torricellian vacuum, as shown in fig. 54,
      and the number of volumes occupied by the vapour noted when the
      space holding it is heated to the desired temperature _t_.

      [Illustration: FIG. 54.--Hofmann's apparatus for determining
      vapour densities. The internal tube, about one metre long, which
      is calibrated and graduated, is filled with mercury and inverted
      in a mercury bath. A small bottle (depicted in its natural size
      on the left) containing a weighed quantity of the liquid whose
      vapour density is to be determined, is introduced into the
      Torricellian vacuum. Steam, or the vapour of amyl alcohol, &c., is
      passed through the outer tube, and heats the internal tube to the
      temperature _t_, at which the volume of vapour is measured.]

      [Illustration: FIG. 55.--Victor Meyer's apparatus for determining
      vapour densities. The tube _b_ is heated in the vapour of a liquid
      of constant boiling point. A glass tube, containing the liquid
      to be experimented upon, is caused to fall from _d_. The air
      displaced is collected in the cylinder _e_, in the trough _f_.]

      _The method of displacement_ (_c_) proposed by Victor Meyer
      is based on the fact that a space _b_ is heated to a constant
      temperature _t_ (by the surrounding vapours of a liquid of
      constant boiling point), and the air (or other gas enclosed in
      this space) is allowed to attain this temperature, and when it
      has done so a glass bulb containing a weighed quantity of the
      substance to be experimented with is dropped into the space. The
      substance is immediately converted into vapour, and displaces
      the air into the graduated cylinder _e_. The amount of this air
      is calculated from its volume, and hence the volume at _t_, and
      therefore also the volume occupied by the vapour, is found. The
      general arrangement of the apparatus is given in fig. 55.

Such an investigation (either direct, or by calculation from the
densities and composition) of every chemical reaction, resulting in the
formation of definite chemical compounds, shows that the volumes of
the reacting substances in a gaseous or vaporous state are either equal
or are in simple multiple proportion.[3] This forms the _first law_
of those discovered by _Gay-Lussac_. It may be formulated as follows:
_The amounts of substances entering into chemical reaction occupy under
similar physical conditions, in a gaseous or vaporous state, equal or
simple multiple volumes._ This law refers not only to elements, but
also to compounds entering into mutual chemical combination; thus, for
example, one volume of ammonia gas combines with one volume of hydrogen
chloride. For in the formation of sal-ammoniac, NH_{4}Cl, there enter
into reaction 17 parts by weight of ammonia, NH_{3}, which is 8·5 times
denser than hydrogen, and 36·5 parts by weight of hydrogen chloride,
whose vapour density is 18·25 times that of hydrogen, as has been
proved by direct experiment. By dividing the weights by the respective
densities we find that the volume of ammonia, NH_{3}, is equal to two,
and so also the volume of hydrogen chloride. Hence the volumes of the
compounds which here combine together are equal to each other. Taking
into consideration that the law of Gay-Lussac holds good, not only for
elements, but also for compounds, it should be expressed as follows:
_Substances interact with one another in commensurable volumes of their
vapours._[4]

  [3] Vapours and gases, as already explained in the second chapter, are
      subject to the same laws, which are, however, only approximate. It
      is evident that for the deduction of the laws which will presently
      be enunciated it is only possible to take into consideration
      a perfect gaseous state (far removed from the liquid state)
      and chemical invariability in which the _vapour density is
      constant_--that is, the volume of a given gas or vapour varies
      like a volume of hydrogen, air, or other gas, with the pressure
      and temperature.

      It is necessary to make this statement in order that it may be
      clearly seen that the laws of gaseous volumes, which we shall
      describe presently, are in the most intimate connection with the
      laws of the variations of volumes with pressure and temperature.
      And as these latter laws (Chapter II.) are not infallible, but
      only approximately exact, the same, therefore, applies to the laws
      about to be described. And as it is possible to find more exact
      laws (a second approximation) for the variation of _v_ with _p_
      and _t_ (for example, van der Waals' formula, Chapter II., Note
      33), so also a more exact expression of the relation between the
      composition and the density of vapours and gases is also possible.
      But to prevent any doubt arising at the very beginning as to the
      breadth and general application of the laws of volumes, it will be
      sufficient to mention that the density of such gases as oxygen,
      nitrogen, and carbonic anhydride is already known to _remain
      constant_ (within the limits of experimental error) between the
      ordinary temperature and a white heat; whilst, judging from what
      is said in my work on the 'Tension of Gases' (vol. i. p. 9),
      it may be said that, as regards pressure, the relative density
      remains very constant, even when the deviations from Mariotte's
      law are very considerable. However, in this respect the number of
      data is as yet too small to arrive at an exact conclusion.

  [4] We must recollect that this law is only approximate, like Boyle
      and Mariotte's law, and that, therefore, like the latter, a more
      exact expression may be found for the exceptions.

The law of combining volumes and the law of multiple proportion were
discovered independently of each other--the one in France by Gay-Lussac,
the other in England by Dalton--almost simultaneously. In the language
of the atomic hypothesis it may be said that atomic quantities of
elements occupy equal or multiple volumes.

The first law of Gay-Lussac expresses the relation between the volumes
of the component parts of a compound. Let us now consider the relation
existing between the volumes of the component parts and of the compounds
which proceed from them. This may sometimes be determined by direct
observation. Thus the volume occupied by water, formed by two volumes
of hydrogen and one volume of oxygen, may be determined by the aid of
the apparatus shown in fig. 56. The long glass tube is closed at the
top and open at the bottom, which is immersed in a cylinder containing
mercury. The closed end is furnished with wires like a eudiometer. The
tube is filled with mercury, and then a certain volume of detonating gas
is introduced. This gas is obtained from the decomposition of water,
and therefore in every three volumes contains two volumes of hydrogen
and one volume of oxygen. The tube is surrounded by a second and wider
glass tube, and the vapour of a substance boiling above 100°--that is,
whose boiling point is higher than that of water--is passed through
the annular space between them. Amyl alcohol, whose boiling point is
132°, may be taken for this purpose. The amyl alcohol is boiled in the
vessel to the right hand and its vapour passed between the walls of the
two tubes. In the case of amyl alcohol the outer glass tube should be
connected with a condenser to prevent the escape into the air of the
unpleasant-smelling vapour. The detonating gas is thus heated up to a
temperature of 132°. When its volume becomes constant it is measured,
the height of the column of mercury in the tube above the level of the
mercury in the cylinder being noted. Let this volume equal _v_; it will
therefore contain 1/3 _v_ of oxygen and 2/3 _v_ of hydrogen. The current
of vapour is then stopped, and the gas exploded; water is formed, which
condenses into a liquid. The volume occupied by the vapour of the water
formed has now to be determined. For this purpose the vapour of the amyl
alcohol is again passed between the tubes, and thus the whole of the
water formed is converted into vapour at the same temperature as that at
which the detonating gas was measured; and the cylinder of mercury being
raised until the column of mercury in the tube stands at the same height
above the surface of the mercury in the cylinder as it did before the
explosion, it is found that the volume of the water formed is equal to
2/3 _v_--that is, it is equal to the volume of the hydrogen contained
in it. Consequently the volumetric composition of water is expressed in
the following terms: Two volumes of hydrogen combine with one volume of
oxygen to form two volumes of aqueous vapour. For substances which are
gaseous at the ordinary temperature, this direct method of observation
is sometimes very easily conducted; for instance, with ammonia, nitric
and nitrous oxides. Thus to determine the composition by volume of
nitrous oxide, the above-described apparatus may be employed. Nitrous
oxide is introduced into the tube, and after measuring its volume
electric sparks are passed through the gas; it is then found that two
volumes of nitrous oxide have given three volumes of gases--namely,
two volumes of nitrogen and one volume of oxygen. Consequently the
composition of nitrous oxide is similar to that of water; two volumes of
nitrogen and one volume of oxygen give two volumes of nitrous oxide. By
decomposing ammonia it is found to be composed in such a manner that two
volumes give one volume of nitrogen and three volumes of hydrogen; also
two volumes of nitric oxide are formed by the union of one volume of
oxygen with one volume of nitrogen. The same relations may be proved by
calculation from the vapour densities, as was described above.

[Illustration: FIG. 56.--Apparatus for demonstrating the volume occupied
by the steam formed from the explosion of detonating gas.]

Comparisons of various results made by the aid of direct observations or
calculation, an example of which has just been cited, led Gay-Lussac to
the conclusion that _the volume of a compound in a gaseous or vaporous
state is always in simple multiple proportion to the volume of each of
the component parts of which it is formed_ (and consequently to the
sum of the volumes of the elements of which it is formed). This is the
_second law of Gay-Lussac_; it extends the simplicity of the volumetric
relations to compounds, and is of the same nature as that presented
by the elements entering into mutual combination. Hence not only the
substances forming a given compound, but also the substances formed,
exhibit a simple relation of volume when measured as vapour or gas.[5]

  [5] This second law of volumes may be considered as a consequence of the
      first law. The first law requires simple ratios between the
      volumes of the combining substances _A_ and _B_. A substance _AB_
      is produced by their combination. It may, according to the law
      of multiple proportion, combine, not only with substances _C_,
      _D_, &c., but also with _A_ and with _B_. In this new combination
      the volume of _AB_, combining with the volume of _A_, should be
      in simple multiple proportion with the volume of _A_; hence the
      volume of the compound _AB_ is in simple proportion to the volume
      of its component parts. Therefore only one law of volumes need be
      accepted. We shall afterwards see that there is a third law of
      volumes embracing also the two first laws.

When a compound is formed from two or more components, there may
or may not be a contraction; the volume of the reacting substances
is in this case either equal to or greater than the volume of the
resultant compound. The reverse is naturally observed in the case of
decompositions, when from one substance there are produced several of
simpler nature. Therefore in the future we shall term _combination_ a
reaction in which a contraction is observed--that is, a diminution in
the volume of the component bodies in a state of vapour or gas; and we
shall term _decomposition_ a reaction in which an expansion is produced;
while those reactions in which the volumes in a gaseous or vaporous
state remain constant (the volumes being naturally compared at the same
temperature and pressure) we shall term reactions of _substitution_ or
of double decomposition. Thus the transition of oxygen into ozone is a
reaction of combination, the formation of nitrous oxide from oxygen and
nitrogen will also be a combination, the formation of nitric oxide from
the same will be a reaction of substitution, the action of oxygen on
nitric oxide a combination, and so on.

The degree of contraction produced in the formation of chemical
compounds not unfrequently leads to the possibility of distinguishing
the degree of change which takes place in the chemical character of the
components when combined. In those cases in which a contraction occurs,
the properties of the resultant compound are very different from the
properties of the substances of which it is composed. Thus ammonia bears
no resemblance in its physical or chemical properties to the elements
from which it is derived; a contraction takes place in a state of
vapour, indicating a proximation of the elements--the distance between
the atoms is diminished, and from gaseous substances there is formed a
liquid substance, or at any rate one which is easily liquefied. For this
reason nitrous oxide formed by the condensation of two permanent gases
is a substance which is somewhat easily converted into a liquid; again,
nitric acid, which is formed from elements which are permanent gases,
is a liquid, whilst, on the contrary, nitric oxide, which is formed
without contraction and is decomposed without expansion, remains a gas
which is as difficult to liquefy as nitrogen and oxygen. In order to
obtain a still more complete idea of the dependence of the properties
of a compound on the properties of the component substances, it is
further necessary to know the quantity of heat which is developed in the
formation of the compound. If this quantity be large--as, for example,
in the formation of water--then the amount of energy in the resultant
compound will be considerably less than the energy of the elements
entering into its composition; whilst, on the contrary, if the amount of
heat evolved in the formation of a compound be small, or if there even
be an absorption of heat, as in the formation of nitrous oxide, then the
energy of the elements is not destroyed, or is only altered to a slight
extent; hence, notwithstanding the contraction (compression) involved in
its formation, nitrous oxide supports combustion.

The preceding laws were deduced from purely experimental and empirical
data and as such evoke further consequences, as the law of multiple
proportions gave rise to the atomic theory and the law of equivalents
(Chapter IV.) In view of the atomic conception of the constitution of
substances, the question naturally arises as to what, then, are the
relative volumes proper to those physically indivisible molecules which
chemically react on each other and consist of the atoms of elements.
The simplest possible hypothesis in this respect would be that the
volumes of the molecules of substances are equal; or, what is the same
thing, to suppose that equal volumes of vapours and gases contain an
equal number of molecules. This proposition was first enunciated by the
Italian savant _Avogadro_ in 1810. It was also admitted by the French
physico-mathematician _Ampère_ (1815) for the sake of simplifying all
kinds of physico-mathematical conceptions respecting gases. But Avogadro
and Ampère's propositions were not generally received in science
until Gerhardt in the forties had applied them to the generalisation
of chemical reactions, and had demonstrated, by aid of a series of
phenomena, that the reactions of substances actually take place with
the greatest simplicity, and more especially that such reactions take
place between those quantities of substances which occupy equal volumes,
and until he had stated the hypothesis in an exact manner and deduced
the consequences that necessarily follow from it. Following Gerhardt,
Clausius, in the fifties, placed this hypothesis of the equality of
the number of molecules in equal volumes of gases and vapours on the
basis of the kinetic theory of gases. At the present day the hypothesis
of Avogadro and Gerhardt lies at the basis of contemporary physical,
mechanical, and chemical conceptions; the consequences arising from it
have often been subject to doubt, but in the end have been verified by
the most diverse methods; and now, when all efforts to refute those
consequences have proved fruitless, the hypothesis must be considered
as verified,[6] and the _law of Avogadro-Gerhardt_ must be spoken of as
fundamental, and as of great importance for the comprehension of the
phenomena of nature. The law may now be formulated from two points of
view. In the first place, from a physical aspect: _equal volumes of
gases_ (or vapours) at equal temperatures and pressures _contain the
same number of molecules_--or of particles of matter which are neither
mechanically nor physically divisible--previous to chemical change. In
the second place, from a chemical aspect, the same law may be expressed
thus: _the quantities of substances entering into chemical reactions
occupy, in a state of vapour, equal volumes_. For our purpose the
chemical aspect is the most important, and therefore, before developing
the law and its consequences, we will consider the chemical phenomena
from which the law is deduced or which it serves to explain.

  [6] It must not be forgotten that Newton's law of gravity was first a
      hypothesis, but it became a trustworthy, perfect theory, and
      acquired the qualities of a fundamental law owing to the concord
      between its deductions and actual facts. All laws, all theories,
      of natural phenomena, are at first hypotheses. Some are rapidly
      established by their consequences exactly agreeing with facts;
      others only take root by slow degrees; and there are many which
      are destined to be refuted owing to their consequences being found
      to be at variance with facts.

When two isolated substances interact with each other directly and
easily--as, for instance, an alkali and an acid--then it is found that
the reaction is accomplished between quantities which in a gaseous
state occupy equal volumes. Thus ammonia, NH_{3}, reacts directly with
hydrochloric acid, HCl, forming sal-ammoniac, NH_{4}Cl, and in this
case the 17 parts by weight of ammonia occupy the same volume as the
36·5 parts by weight of hydrochloric acid.[7] Ethylene, C_{2}H_{4},
combines with chlorine, Cl_{2}, in only one proportion, forming
ethylene dichloride, C_{2}H_{4}Cl_{2}, and this combination proceeds
directly and with great facility, the reacting quantities occupying
equal volumes. Chlorine reacts with hydrogen in only one proportion,
forming hydrochloric acid, HCl, and in this case equal volumes interact
with each other. If an equality of volumes is observed in cases of
combination, it should be even more frequently encountered in cases
of decomposition, taking place in substances which split up into two
others. Indeed, acetic acid breaks up into marsh gas, CH_{4}, and
carbonic anhydride, CO_{2}, and in the proportions in which they
are formed from acetic acid they occupy equal volumes. Also from
phthalic acid, C_{8}H_{6}O_{4}, there may be obtained benzoic acid,
C_{7}H_{6}O_{2}, and carbonic anhydride, CO_{2}, and as all the elements
of phthalic acid enter into the composition of these substances, it
follows that, although they cannot re-form it by their direct action
on each other (the reaction is not reversible), still they form the
direct products of its decomposition, and they occupy equal volumes.
But benzoic acid, C_{7}H_{6}O_{2}, is itself composed of benzene,
C_{6}H_{6}, and carbonic anhydride, CO_{2}, which also occupy equal
volumes.[8] There is an immense number of similar examples among those
organic substances to whose study Gerhardt consecrated his whole life
and work, and he did not allow such facts as these to escape his
attention. Still more frequently in the phenomena of substitution,
when two substances react on one another, and two are produced without
a change of volume, it is found that the two substances acting on
each other occupy equal volumes as well as each of the two resultant
substances. Thus, in general, reactions of substitution take place
between volatile acids, HX, and volatile alcohols, R(OH), with the
formation of ethereal salts, RX, and water, H(OH), and the volume of the
vapour of the reacting quantities, HX, R(OH), and RX, is the same as
that of water H(OH), whose weight, corresponding with the formula, 18,
occupies 2 volumes, if 1 part by weight of hydrogen occupy 1 volume and
the density of aqueous vapour referred to hydrogen is 9. Such general
examples, of which there are many,[9] show that the reaction of equal
volumes forms a chemical phenomenon of frequent occurrence, indicating
the necessity for acknowledging the law of Avogadro-Gerhardt.

  [7] This is not only seen from the above calculations, but may be proved
      by experiment. A glass tube, divided in the middle by a stopcock,
      is taken and one portion filled with _dry_ hydrogen chloride
      (the dryness of the gases is very necessary, because ammonia
      and hydrogen chloride are both very soluble in water, so that a
      small trace of water may contain a large amount of these gases in
      solution) and the other with dry ammonia, under the atmospheric
      pressure. One orifice (for instance, of that portion which
      contains the ammonia) is firmly closed, and the other is immersed
      under mercury, and the cock is then opened. Solid sal-ammoniac is
      formed, but if the volume of one gas be greater than that of the
      other, some of the first gas will remain. By immersing the tube in
      the mercury in order that the internal pressure shall equal the
      atmospheric pressure, it may easily be shown that the volume of
      the remaining gas is equal to the difference between the volumes
      of the two portions of the tube, and that this remaining gas is
      part of that whose volume was the greater.

  [8] Let us demonstrate this by figures. From 122 grams of benzoic acid
      there are obtained (_a_) 78 grams of benzene, whose density
      referred to hydrogen = 39, hence the relative volume = 2; and
      (_b_) 44 grams of carbonic anhydride, whose density = 22, and
      hence the volume = 2. It is the same in other cases.

  [9] A large number of such generalised reactions, showing reaction
      by equal volumes, occur in the case of the hydrocarbon
      derivatives, because many of these compounds are volatile. The
      reactions of alkalis on acids, or anhydrides on water, &c.,
      which are so frequent between mineral substances, present but
      few such examples, because many of these substances are not
      volatile and their vapour densities are unknown. But essentially
      the same is seen in these cases also; for instance, sulphuric
      acid, H_{2}SO_{4}, breaks up into the anhydride, SO_{3}, and
      water, H_{2}O, which exhibit an equality of volumes. Let us take
      another example where three substances combine in equal volumes:
      carbonic anhydride, CO_{2}, ammonia, NH_{3}, and water, H_{2}O
      (the volumes of all are equal to 2), form acid ammonium carbonate,
      (NH_{4})HCO_{3}.

But the question arises, What is the relation of volumes if the reaction
of two substances takes place in more than one proportion, according to
the law of multiple proportions? A definite answer can only be given in
cases which have been very thoroughly studied. Thus chlorine, in acting
on marsh gas, CH_{4}, forms four compounds, CH_{3}Cl, CH_{2}Cl_{2},
CHCl_{3}, and CCl_{4}, and it may be established by direct experiment
that the substance CH_{3}Cl (methylic chloride) precedes the remainder,
and that the latter proceed from it by the further action of chlorine.
And this substance, CH_{3}Cl, is formed by the reaction of equal volumes
of marsh gas, CH_{4}, and chlorine, Cl_{2}, according to the equation
CH_{4} + Cl_{2} = CH_{3}Cl + HCl. A great number of similar cases are
met with amongst organic--that is, carbon--compounds. Gerhardt was led
to the discovery of his law by investigating many such reactions, and by
observing that in them the reaction of equal volumes precedes all others.

But if nitrogen or hydrogen give several compounds with oxygen, the
question proposed above cannot be answered with complete clearness,
because the successive formations of the different combinations cannot
be so strictly defined. It may be supposed, but neither definitely
affirmed nor experimentally confirmed, that nitrogen and oxygen
first give nitric oxide, NO, and only subsequently the brown vapours
N_{2}O_{3} and NO_{2}. Such a sequence in the combination of nitrogen
with oxygen can only be supposed on the basis of the fact that NO forms
N_{2}O_{3} and NO_{2} directly with oxygen. If it be admitted that NO
(and not N_{2}O or NO_{2}) be first formed, then this instance would
also confirm the law of Avogadro-Gerhardt, because nitric oxide contains
equal volumes of nitrogen and oxygen. So, also, it may be admitted that,
in the combination of hydrogen with oxygen, hydrogen peroxide is first
formed (equal volumes of hydrogen and oxygen), which is decomposed by
the heat evolved into water and oxygen. This explains the presence of
traces of hydrogen peroxide (Chapter IV.) in almost all cases of the
combustion or oxidation of hydrogenous substances; for it cannot be
supposed that water is first formed and then the peroxide of hydrogen,
because up to now such a reaction has not been observed, whilst the
formation of H_{2}O from H_{2}O_{2} is very easily reproduced.[10]

  [10] This opinion which I have always held (since the first editions
       of this work), as to the primary origin of hydrogen peroxide
       and of the formation of water by means of its decomposition,
       has in latter days become more generally accepted, thanks more
       especially to the work of Traube. Probably it explains most
       simply the necessity for the presence of traces of water in
       many reactions, as, for instance, in the explosion of carbonic
       oxide with oxygen, and perhaps the theory of the explosion of
       detonating gas itself and of the combustion of hydrogen will
       gain in clearness and truth if we take into consideration the
       preliminary formation of hydrogen peroxide and its decomposition.
       We may here point out the fact that Ettingen (at Dorpat, 1888)
       observed the existence of currents and waves in the explosion of
       detonating gas by taking photographs, which showed the periods of
       combustion and the waves of explosion, which should be taken into
       consideration in the theory of this subject. As the formation of
       H_{2}O_{2} from O_{2} and H_{2} corresponds with a less amount of
       heat than the formation of water from H_{2} and O, it may be that
       the temperature of the flame of detonating gas depends on the
       pre-formation of hydrogen peroxide.

Thus a whole series of phenomena show that the chemical reaction of
substances actually takes place, as a rule, between equal volumes,
but this does not preclude the possibility of the frequent reaction
of unequal volumes, although, in this case, it is often possible to
discover a preceding reaction between equal volumes.[11]

  [11] The possibility of reactions between unequal volumes,
       notwithstanding the general application of the law of
       Avogadro-Gerhardt, may, in addition to what has been said
       above, depend on the fact that the participating substances,
       at the moment of reaction, undergo a preliminary modification,
       decomposition, isomeric (polymeric) transformation, &c. Thus,
       if NO_{2}, seems to proceed from N_{2}O_{4}, if O_{2} is formed
       from O_{3}, and the converse, then it cannot be denied that
       the production of molecules containing only one atom is also
       possible--for instance, of oxygen--as also of higher polymeric
       forms--as the molecule N from N_{2}, or H_{3} from H_{2}. In
       this manner it is obviously possible, by means of a series of
       hypotheses, to explain the cases of the formation of ammonia,
       NH_{3}, from 3 vols. of hydrogen and 1 vol. of nitrogen. But
       it must be observed that perhaps our information in similar
       instances is, as yet, far from being complete. If hydrazine or
       diamide N_{2}H_{4} (Chapter VI. Note 20 bis) is formed and the
       imide N_{2}H_{2} in which 2 vols. of hydrogen are combined with
       2 vols. of nitrogen, then the reaction here perhaps first takes
       place between equal volumes. If it be shown that diamide gives
       nitrogen and ammonia (3N_{2}H_{4} = N_{2} + 4NH_{3}) under the
       action of sparks, heat, or the silent discharge, &c., then it
       will be possible to admit that it is formed before ammonia. And
       perhaps the still less stable imide N_{2}H_{2}, which may also
       decompose with the formation of ammonia, is produced before the
       amide N_{2}H_{4}.

       I mention this to show that the fact of apparent exceptions
       existing to the law of reactions between equal volumes does not
       prove the impossibility of their being included under the law
       on further study of the subject. Having put forward a certain
       law or hypothesis, consequences must be deduced from it, and
       if by their means clearness and consistency are attained--and
       especially, if by their means that which could not otherwise
       be known can be predicted--then the consequences verify the
       hypothesis. This was the case with the law now under discussion.
       The mere simplicity of the deduction of the weights proper to
       the atoms of the elements, or the mere fact that having admitted
       the law it follows (as will afterwards be shown) that the _vis
       viva_ of the molecules of all gases is a constant quantity, is
       quite sufficient reason for retaining the hypothesis, if not
       for believing in it as a fact beyond doubt. And such is the
       whole doctrine of atoms. And since by the acceptance of the law
       it became possible to foretell even the properties and atomic
       weights of elements which had not yet been discovered, and these
       predictions afterwards proved to be in agreement with the actual
       facts, it is evident that the law of Avogadro-Gerhardt penetrates
       deeply into the nature of the chemical relation of substances.
       This being granted, it is possible at the present time to exhibit
       and deduce the truth under consideration in many ways, and in
       every case, like all that is highest in science (for example,
       the laws of the indestructibility of matter, of the conservation
       of energy, of gravity, &c.), it proves to be not an empirical
       conclusion from direct observation and experiment, not a direct
       result of analysis, but a creation, or instinctive penetration,
       of the inquiring mind, guided and directed by experiment and
       observation--a synthesis of which the exact sciences are
       capable equally with the highest forms of art. Without such a
       synthetical process of reasoning, science would only be a mass
       of disconnected results of arduous labour, and would not be
       distinguished by that vitality with which it is really endowed
       when once it succeeds in attaining a synthesis, or concordance
       of outward form with the inner nature of things, without losing
       sight of the diversities of individual parts; in short, when
       it discovers by means of outward phenomena, which are apparent
       to the sense of touch, to observation, and to the common mind,
       the internal signification of things--discovering simplicity in
       complexity and uniformity in diversity. And this is the highest
       problem of science.

The law of Avogadro-Gerhardt may also be easily expressed in an
algebraical form. If the weight of a molecule, or of that quantity
of a substance which enters into chemical reaction and occupies in a
state of vapour, according to the law, a volume equal to that occupied
by the molecules of other bodies, be indicated by the letters M_{1},
M_{2} ... or, in general, M, and if the letters D_{1}, D_{2}, ... or,
in general, D, stand for the density or weight of a given volume of the
gases or vapours of the corresponding substances under certain definite
conditions of temperature and pressure, then the law requires that

                 M_{1}/D_{1} = M_{2}/D_{2} ... = M/D = C

where C is a certain constant. This expression shows directly that the
volumes corresponding with the weights M_{1}, M_{2} ... M, are equal to
a certain constant, because the volume is proportional to the weight and
inversely proportional to the density. The magnitude of C is naturally
conditioned by and dependent on the units taken for the expression of
the weights of the molecules and the densities. The weight of a molecule
(equal to the sum of the atomic weights of the elements forming it) is
usually expressed by taking the weight of an atom of hydrogen as unity,
and hydrogen is now also chosen as the unit for the expression of the
densities of gases and vapours; it is therefore only necessary to find
the magnitude of the constant for any one compound, as it will be the
same for all others. Let us take water. Its reacting mass is expressed
(conditionally and relatively) by the formula or molecule H_{2}O,
for which M = 18, if H = 1, as we already know from the composition
of water. Its vapour density, or D, compared to hydrogen = 9, and
consequently for water C = 2, and therefore and in general for the
molecules of all substances M/D = 2.

Consequently the weight of a molecule is equal to twice its vapour
density expressed in relation to hydrogen, and conversely _the density
of a gas is equal to half the molecular weight referred to hydrogen_.

The truth of this may be seen from a very large number of observed
vapour densities by comparing them with the results obtained by
calculation. As an illustration, we may point out that for ammonia,
NH_{3}, the weight of the molecule or quantity of the reacting
substance, as well as the composition and weight corresponding with
the formula, is expressed by the figures 14 + 3 = 17. Consequently M
= 17. Hence, according to the law, D = 8·5. And this result is also
obtained by experiment. The density, according to both formula and
experiment, of nitrous oxide, N_{2}O, is 22, of nitric acid 15, and of
nitric peroxide 23. In the case of nitrous anhydride, N_{2}O_{3}, as a
substance which dissociates into NO + NO_{2}, the density should vary
between 38 (so long as the N_{2}O_{3} remains unchanged) and 19 (when
NO + NO_{2} is obtained). There are no figures of constant density for
H_{2}O_{2}, NHO_{3}, N_{2}O_{4}, and many similar compounds which are
either wholly or partially decomposed in passing into vapour. Salts and
similar substances either have no vapour density because they do not
pass into vapour (for instance, potassium nitrate, KNO_{3}) without
decomposition, or, if they pass into vapour without decomposing,
their vapour density is observed with difficulty only at very high
temperatures. The practical determination of the vapour density at these
high temperatures (for example, for sodium chloride, ferrous chloride,
stannous chloride, &c.) requires special methods which have been worked
out by Sainte-Claire Deville, Crafts, Nilson and Pettersson, Meyer,
Scott, and others. Having overcome the difficulties of experiment, it
is found that the law of Avogadro-Gerhardt holds good for such salts
as potassium iodide, beryllium chloride, aluminium chloride, ferrous
chloride, &c.--that is, the density obtained by experiment proves to
be equal to half the molecular weight--naturally within the limits of
experimental error or of possible deviation from the law.

Gerhardt deduced his law from a great number of examples of volatile
carbon compounds. We shall become acquainted with certain of them in
the following chapters; their entire study, from the complexity of
the subject, and from long-established custom, forms the subject of a
special branch of chemistry termed 'organic' chemistry. With all these
substances the observed and calculated densities are very similar.

When the consequences of a law are verified by a great number of
observations, it should be considered as confirmed by experiment. But
this does not exclude the possibility of _apparent_ deviations. They may
evidently be of two kinds: the fraction M/D may be found to be either
greater or less than 2--that is, the calculated density may be either
greater or less than the observed density. When the difference between
the results of experiment and calculation falls within the possible
errors of experiment (for example, equal to hundredths of the density),
or within a possible error owing to the laws of gases having an only
approximate application (as is seen from the deviations, for instance,
from the law of Boyle and Mariotte), then the fraction M/D proves but
slightly different from 2 (between 1·9 and 2·2), and such cases as these
may be classed among those which ought to be expected from the nature
of the subject. It is a different matter if the quotient of M/D be
several times, and in general a multiple, _greater_ or less than 2. The
application of the law must then be explained or it must be laid aside,
because the laws of nature admit of no exceptions. We will therefore
take two such cases, and first one in which the _quotient_ M/D _is
greater than 2, or the density obtained by experiment is less than is in
accordance with the law_.

It must be admitted, as a consequence of the law of Avogadro-Gerhardt,
that there is a decomposition in those cases where the volume of the
vapour corresponding with the weight of the amount of a substance
entering into reaction is greater than the volume of two parts by
weight of hydrogen. Suppose the density of the vapour of water to be
determined at a temperature above that at which it is decomposed,
then, if not all, at any rate a large proportion of the water will be
decomposed into hydrogen and oxygen. The density of such a mixture of
gases, or of detonating gas, will be less than that of aqueous vapour;
it will be equal to 6 (compared with hydrogen), because 1 volume of
oxygen weighs 16, and 2 volumes of hydrogen 2; and, consequently, 3
volumes of detonating gas weigh 18 and 1 volume 6, while the density
of aqueous vapour = 9. Hence, if the density of aqueous vapour be
determined after its decomposition, the quotient M/D would be found
to be 3 and not 2. This phenomenon might be considered as a deviation
from Gerhardt's law, but this would not be correct, because it may be
shown by means of diffusion through porous substances, as described in
Chapter II., that water is decomposed at such high temperatures. In the
case of water itself there can naturally be no doubt, because its vapour
density agrees with the law at all temperatures at which it has been
determined.[12] But there are many substances which decompose with great
ease directly they are volatilised, and therefore only exist as solids
or liquids, and not in a state of vapour. There are, for example, many
salts of this kind, besides all definite solutions having a constant
boiling point, all the compounds of ammonia for example, all ammonium
salts--&c. Their vapour densities, determined by Bineau, Deville, and
others, show that they do not agree with Gerhardt's law. Thus the
vapour density of sal-ammoniac, NH_{4}Cl, is nearly 14 (compared with
hydrogen), whilst its molecular weight is not less than 53·5, whence the
vapour density should be nearly 27, according to the law. The molecule
of sal-ammoniac cannot be less than NH_{4}Cl, because it is formed from
the molecules NH_{3} and HCl, and contains single atoms of nitrogen
and chlorine, and therefore cannot be divided; it further never enters
into reactions with the molecules of other substances (for instance,
potassium hydroxide, or nitric acid) in quantities of less than 53·5
parts by weight, &c. The calculated density (about 27) is here double
the observed density (about 13·4); hence M/D = 4 and not 2. For this
reason the vapour density of sal-ammoniac for a long time served as an
argument for doubting the truth of the law. But it proved otherwise,
after the matter had been fully investigated. The low density depends
on the decomposition of sal-ammoniac, on volatilising, into ammonia and
hydrogen chloride. The observed density is not that of sal-ammoniac, but
of a mixture of NH_{3} and HCl, which should be nearly 14, because the
density of NH_{3} = 8·5 and of HCl = 18·2, and therefore the density
of their mixture (in equal volumes) should be about 13·4.[13] The
actual decomposition of the vapours of sal-ammoniac was demonstrated
by Pebal and Than by the same method as the decomposition of water,
by passing the vapour of sal-ammoniac through a porous substance. The
experiment demonstrating the decomposition during volatilisation of
sal-ammoniac may be made very easily, and is a very instructive point
in the history of the law of Avogadro-Gerhardt, because without its
aid it would never have been imagined that sal-ammoniac decomposed
in volatilising, as this decomposition bears all the signs of simple
sublimation; consequently the knowledge of the decomposition itself
was forestalled by the law. The whole aim and practical use of the
discovery of the laws of nature consists in, and is shown by, the fact
that they enable the unknown to be foretold, the unobserved to be
foreseen. The arrangement of the experiment is based on the following
reasoning.[14] According to the law and to experiment, the density of
ammonia, NH_{3}, is 8-1/2, and of hydrochloric acid, HCl, 18-1/4, if
the density of hydrogen = 1. Consequently, in a mixture of NH_{3} and
HCl, the ammonia will penetrate much more rapidly through a porous
mass, or a fine orifice, than the heavier hydrochloric acid, just as
in a former experiment the hydrogen penetrated more rapidly than the
oxygen. Therefore, if the vapour of sal-ammoniac comes into contact with
a porous mass, the ammonia will pass through it in greater quantities
than the hydrochloric acid, and this excess of ammonia may be detected
by means of moist red litmus paper, which should be turned blue. If the
vapour of sal-ammoniac were not decomposed, it would pass through the
porous mass as a whole, and the colour of the litmus paper would not
be altered, because sal-ammoniac is a neutral salt. Thus, by testing
with litmus the substances passing through the porous mass, it may be
decided whether the sal-ammoniac is decomposed or not when passing into
vapour. Sal-ammoniac volatilises at so moderate a temperature that the
experiment may be conducted in a glass tube heated by means of a lamp,
an asbestos plug being placed near the centre of the tube.[15] The
asbestos forms a porous mass, which is unaltered at a high temperature.
A piece of dry sal-ammoniac is placed at one side of the asbestos
plug, and is heated by a Bunsen burner. The vapours formed are driven
by a current of air forced from a gasometer or bag through two tubes
containing pieces of moist litmus paper, one blue and one red paper
in each. If the sal-ammoniac be heated, then the ammonia appears on
the opposite side of the asbestos plug, and the litmus there turns
blue. And as an excess of hydrochloric acid remains on the side where
the sal-ammoniac is heated, it turns the litmus at that end red. This
proves that the sal-ammoniac, when converted into vapour, splits up into
ammonia and hydrochloric acid, and at the same time gives an instance of
the possibility of correctly conjecturing a fact on the basis of the law
of Avogadro-Gerhardt.[15 bis]

  [12] As the density of aqueous vapour remains constant within the
       limits of experimental accuracy, even at 1,000°, when
       dissociation has certainly commenced, it would appear that only a
       very small amount of water is decomposed at these temperatures.
       If even 10 p.c. of water were decomposed, the density would be
       8·57 and the quotient M/D = 2·1, but at the high temperatures
       here concerned the error of experiment is not greater than the
       difference between this quantity and 2. And probably at 1,000°
       the dissociation is far from being equal to 10 p.c. _Hence the
       variation in the vapour density of water does not give us the
       means of ascertaining the amount of its dissociation._

  [13] This explanation of the vapour density of sal-ammoniac, sulphuric
       acid, and similar substances which decompose in being distilled
       was the most natural to resort to as soon as the application of
       the law of Avogadro-Gerhardt to chemical relations was begun; it
       was, for instance, given in my work on _Specific Volumes_, 1856,
       p. 99. The formula, M/D = 2, which was applied later by many
       other investigators, had already been made use of in that work.

  [14] The beginner must remember that an experiment and the mode in
       which it is carried out must be determined by the principle or
       fact which it is intended to illustrate, and not _vice versa_,
       as some suppose. The idea which determines the necessity of an
       experiment is the chief consideration.

  [15] It is important that the tubes, asbestos, and sal-ammoniac should
       be dry, as otherwise the moisture retains the ammonia and
       hydrogen chloride.

  [15 bis] Baker (1894) showed that the decomposition of NH_{4}Cl in
       the act of volatilising only takes place in the presence of water,
       traces of which are amply sufficient, but that in the total
       absence of moisture (attained by carefully drying with P_{2}O_{5})
       there is no decomposition, and the vapour density of the
       sal-ammoniac is found to be normal, _i.e._, nearly 27. It is not
       yet quite clear what part the trace of moisture plays here, and it
       must be presumed that the phenomenon belongs to the category of
       electrical and contact phenomena, which have not yet been fully
       explained (_see_ Chapter IX., Note 29).

So also the fact of a decomposition may be proved in the other instances
where M/D proved greater than 2, and hence the apparent deviations
appear in reality as an excellent proof of the general application and
significance of the law of Avogadro-Gerhardt.

In those cases where the _quotient_ M/D proves to be _less_ than 2,
or the observed density _greater_ than that calculated, by a multiple
number of times, the matter is evidently more simple, and the fact
observed only indicates that the weight of the molecule is as many times
greater as that taken as the quotient obtained is less than 2. So, for
instance, in the case of ethylene, whose composition is expressed by
CH_{2}, the density was found by experiment to be 14, and in the case
of amylene, whose composition is also CH_{2}, the density proved to be
35, and consequently the quotient for ethylene = 1, and for amylene =
2/5. If the molecular weight of ethylene be taken, not as 14, as might
be imagined from its composition, but as twice as great--namely, as
28--and for amylene as five times greater--that is as 70--then the
molecular composition of the first will be C_{2}H_{4}, and of the
second C_{5}H_{10}, and for both of them M/D will be equal to 2. This
application of the law, which at first sight may appear perfectly
arbitrary, is nevertheless strictly correct, because the amount of
ethylene which reacts--for example, with sulphuric and other acids--is
not equal to 14, but to 28 parts by weight. Thus with H_{2}SO_{4},
Br_{2}, or HI, &c., ethylene combines in a quantity C_{2}H_{4}, and
amylene in a quantity C_{5}H_{10}, and not CH_{2}. On the other hand,
ethylene is a gas which liquefies with difficulty (absolute boiling
point = +10°), whilst amylene is a liquid boiling at 35° (absolute
boiling point = +192°), and by admitting the greater density of the
molecules of amylene (M = 70) its difference from the lighter molecules
of ethylene (M = 28) becomes clear. Thus, the smaller quotient M/D
is _an indication of polymerisation_, as the larger quotient is of
decomposition. The difference between the densities of oxygen and ozone
is a case in point.

On turning to the elements, it is found in certain cases, especially
with metals--for instance, mercury, zinc, and cadmium--that that weight
of the atoms which must be acknowledged in their compounds (of which
mention will be afterwards made) appears to be also the molecular
weight. Thus, the atomic weight of mercury must be taken as = 200,
but the vapour density = 100, and the quotient = 2. Consequently the
_molecule of mercury contains one atom_, Hg. It is the same with sodium,
cadmium, and zinc. This is the simplest possible molecule, which
necessarily is only possible in the case of elements, as the molecule
of a compound must contain at least two atoms. However, the molecules
of many of the elements prove to be complex--for instance, the weight
of an atom of oxygen = 16, and its density = 16, so that its molecule
must contain two atoms, O_{2}, which might already be concluded by
comparing its density with that of ozone, whose molecule contains O_{3}
(Chapter IV.) So also the molecule of hydrogen equals H_{2}, of chlorine
Cl_{2}, of nitrogen N_{2}, &c. If chlorine react with hydrogen, the
volume remains unaltered after the formation of hydrochloric acid,
H_{2} + Cl_{2} = HCl + HCl. It is a case of substitution between the
one and the other, and therefore the volumes remain constant. There
are elements whose molecules are much more complex--for instance,
sulphur, S_{6}--although, by heating, the density is reduced to a third,
and S_{2} is formed. Judging from the vapour density of phosphorus
(D = 62) the molecule contains four atoms P_{4}. Hence many elements
when polymerised appear in molecules which are more complex than the
simplest possible. In carbon, as we shall afterwards find, a very
complex molecule must be admitted, as otherwise its non-volatility and
other properties cannot be understood. And if compounds are decomposed
by a more or less powerful heat, and if polymeric substances are
depolymerised (that is, the weight of the molecule diminishes) by a rise
of temperature, as N_{2}O_{4} passes into NO_{2}, or ozone, O_{3}, into
ordinary oxygen, O_{2}, then we might expect to find the splitting-up of
the complex molecules of elements into the simplest molecule containing
a single atom only--that is to say, if O_{2} be obtained from O_{3},
then the formation of O might also be looked for. The possibility but
not proof of such a proposition is indicated by the vapour of iodine.
Its normal density = 127 (Dumas, Deville, and others), which corresponds
with the molecule I_{2}. At temperatures above 800° (up to which the
density remains almost constant), this density distinctly decreases, as
is seen from the verified results obtained by Victor Meyer, Crafts, and
Troost. At the ordinary pressure and 1,000° it is about 100, at 1,250°
about 80, at 1,400° about 75, and apparently it strives to reduce itself
to one-half--that is, to 63. Under a reduced pressure this splitting-up,
or depolymerisation, of iodine vapour actually reaches a density[16]
of 66, as Crafts demonstrated by reducing the pressure to 100 mm. and
raising the temperature to 1,500°. From this it may be concluded that
at high temperatures and low pressures the molecule I_{2} gradually
passes into the molecule I containing one atom like mercury, and that
something similar occurs with other elements at a considerable rise of
temperature, which tends to bring about the disunion of compounds and
the decomposition of complex molecules.[17]

  [16] Just as we saw (Chapter VI. Note 46) an increase of the
       dissociation of N_{2}O_{4} and the formation of a large
       proportion of NO_{2}, with a decrease of pressure. The
       decomposition of I_{2} into I + I is a similar dissociation.

  [17] Although at first there appeared to be a similar phenomenon in the
       case of chlorine, it was afterwards proved that if there is
       a decrease of density it is only a small one. In the case of
       bromine it is not much greater, and is far from being equal to
       that for iodine.

       As in general we very often involuntarily confuse chemical
       processes with physical, it may be that a physical process
       of change in the coefficient of expansion with a change of
       temperature participates with a change in molecular weight,
       and partially, if not wholly, accounts for the decrease of the
       density of chlorine, bromine, and iodine. Thus, I have remarked
       (Comptes Rendus, 1876) that the coefficient of expansion of
       gases increases with their molecular weight, and (Chapter II.,
       Note 26) the results of direct experiment show the coefficient
       of expansion of hydrobromic acid (M = 81) to be 0·00386 instead
       of 0·00367, which is that of hydrogen (M = 2). Hence, in the
       case of the vapour of iodine (M = 254) a very large coefficient
       of expansion is to be expected, and from this cause alone
       the relative density would fall. As the molecule of chlorine
       Cl_{2} is lighter (= 71) than that of bromine (= 160), which is
       lighter than that of iodine (= 254), we see that the order in
       which the decomposability of the vapours of these haloids is
       observed corresponds with the expected rise in the coefficient of
       expansion. Taking the coefficient of expansion of iodine vapour
       as 0·004, then at 1,000° its density would be 116. Therefore
       the dissociation of iodine may be only an apparent phenomenon.
       However, on the other hand, the heavy vapour of mercury (M = 200,
       D = 100) scarcely decreases in density at a temperature of 1,500°
       (D = 98, according to Victor Meyer); but it must not be forgotten
       that the molecule of mercury contains only one atom, whilst that
       of iodine contains two, and this is very important. Questions
       of this kind which are difficult to decide by experimental
       methods must long remain without a certain explanation, owing to
       the difficulty, and sometimes impossibility, of distinguishing
       between physical and chemical changes.

Besides these cases of apparent discrepancy from the law of
Avogadro-Gerhardt there is yet a third, which is the last, and is very
instructive. In the investigation of separate substances they have to be
isolated in the purest possible form, and their chemical and physical
properties, and among them the vapour density, then determined. If
it be normal--that is, if D = M/2--it often serves as a proof of the
purity of the substance, _i.e._ of its freedom from all foreign matter.
If it be abnormal--that is, if D be not equal to M/2--then for those
who do not believe in the law it appears as a new argument against it
and nothing more; but to those who have already grasped the important
significance of the law it becomes clear that there is some error in
the observation, or that the density was determined under conditions in
which the vapour does not follow the laws of Boyle or Gay-Lussac, or
else that the substance has not been sufficiently purified, and contains
other substances. The law of Avogadro-Gerhardt in that case furnishes
convincing evidence of the necessity of a fresh and more exact research.
And as yet the causes of error have always been found. There are not a
few examples in point in the recent history of chemistry. We will cite
one instance. In the case of pyrosulphuryl chloride, S_{2}O_{5}Cl_{2},
M = 215, and consequently D should = 107·5, instead of which Ogier and
others obtained 53·8--that is, a density half as great; and further,
Ogier (1882) demonstrated clearly that the substance is not dissociated
by distillation into SO_{3} and SO_{2}Cl_{2}, or any other two products,
and thus the abnormal density of S_{2}O_{5}Cl_{2} remained unexplained
until D. P. Konovaloff (1885) showed that the previous investigators
were working with a mixture (containing SO_{3}HCl), and that
pyrosulphuryl chloride has a normal density of approximately 107. Had
not the law of Avogadro-Gerhardt served as a guide, the impure liquid
would have still passed as pure; the more so since the determination of
the amount of chlorine could not aid in the discovery of the impurity.
Thus, by following a true law of nature we are led to true deductions.

All cases which have been studied confirm the law of Avogadro-Gerhardt,
and as by it a deduction is obtained, from the determination of the
vapour density (a purely physical property), as to the weight of the
molecule or quantity of a substance entering into chemical reaction,
this law links together the two provinces of learning--physics and
chemistry--in the most intimate manner. Besides which, the law of
Avogadro-Gerhardt places the conceptions of _molecules_ and _atoms_ on
a firm foundation, which was previously wanting. Although since the
days of Dalton it had become evident that it was necessary to admit the
existence of the elementary atom (the chemical individual indivisible
by chemical or other forces), and of the groups of atoms (or molecules)
of compounds, indivisible by mechanical and physical forces; still
the relative magnitude of the molecule and atom was not defined with
sufficient clearness. Thus, for instance, the atomic weight of oxygen
might be taken as 8 or 16, or any multiple of these numbers, and nothing
indicated a reason for the acceptation of one rather than another of
these magnitudes;[18] whilst as regards the weights of the molecules of
elements and compounds there was no trustworthy knowledge whatever. With
the establishment of Gerhardt's law the idea of the molecule was fully
defined, as well as the relative magnitude of the elementary atom.

  [18] And so it was in the fifties. Some took O = 8, others O = 16.
       Water in the first case would be HO and hydrogen peroxide HO_{2},
       and in the second case, as is now generally accepted, water
       H_{2}O and hydrogen peroxide H_{2}O_{2} or HO. Disagreement
       and confusion reigned. In 1860 the chemists of the whole world
       met at Carlsruhe for the purpose of arriving at some agreement
       and uniformity of opinion. I was present at this Congress, and
       well remember how great was the difference of opinion, and how
       a compromise was advocated with great acumen by many scientific
       men, and with what warmth the followers of Gerhardt, at whose
       head stood the Italian professor, Canizzaro, followed up the
       consequences of the law of Avogadro. In the spirit of scientific
       freedom, without which science would make no progress, and would
       remain petrified as in the middle ages, and with the simultaneous
       necessity of scientific conservatism, without which the roots
       of past study could give no fruit, a compromise was not arrived
       at, nor ought it to have been, but instead of it truth, in the
       form of the law of Avogadro-Gerhardt, received by means of the
       Congress a wider development, and soon afterwards conquered all
       minds. Then the new so-called Gerhardt atomic weights established
       themselves, and in the seventies they were already in general use.

The chemical particle or _molecule must be considered as the quantity
of a substance which enters into chemical reaction with other molecules,
and occupies in a state of vapour the same volume as two parts by weight
of hydrogen_.

The molecular weight (which has been indicated by M) of a substance is
determined by its composition, transformations, and vapour density.

The molecule is not divisible by the mechanical and physical changes
of substances, but in chemical reaction it is either altered in its
properties, or quantity, or structure, or in the nature of the motion of
its parts.

An agglomeration of molecules, which are alike in all chemical respects,
makes up the masses of homogeneous substances in all states.[19]

  [19] A bubble of gas, a drop of liquid, or the smallest crystal,
       presents an agglomeration of a number of molecules, in a state of
       continual motion (like the stars of the Milky Way), distributing
       themselves evenly or forming new systems. If the aggregation of
       all kinds of heterogeneous molecules be possible in a gaseous
       state, where the molecules are considerably removed from each
       other, then in a liquid state, where they are already close
       together, such an aggregation becomes possible only in the
       sense of the mutual reaction between them which results from
       their chemical attraction, and especially in the aptitude of
       heterogeneous molecules for combining together. Solutions and
       other so-called indefinite chemical compounds should be regarded
       in this light. According to the principles developed in this
       work we should regard them as containing both the compounds of
       the heterogeneous molecules themselves and the products of their
       decomposition, as in peroxide of nitrogen, N_{2}O_{4} and NO_{2}.
       And we must consider that those molecules A, which at a given
       moment are combined with B in AB, will in the following moment
       become free in order to again enter into a combined form. The
       laws of chemical equilibrium proper to dissociated systems cannot
       be regarded in any other light.

Molecules consist of atoms in a certain state of distribution and
motion, just as the solar system[20] is made up of inseparable parts
(the sun, planets, satellites, comets, &c.) The greater the number of
atoms in a molecule, the more complex is the resultant substance. The
equilibrium between the dissimilar atoms may be more or less stable,
and may for this reason give more or less stable substances. Physical
and mechanical transformations alter the velocity of the motion and
the distances between the individual molecules, or of the atoms in the
molecules, or of their sum total, but they do not alter the original
equilibrium of the system; whilst chemical changes, on the other hand,
alter the molecules themselves, that is, the velocity of motion, the
relative distribution, and the quality and quantity of the atoms in the
molecules.

  [20] This strengthens the fundamental idea of the unity and harmony of
       type of all creation and is one of those ideas which impress
       themselves on man in all ages, and give rise to a hope of
       arriving in time, by means of a laborious series of discoveries,
       observations, experiments, laws, hypotheses, and theories, at
       a comprehension of the internal and invisible structure of
       concrete substances with that same degree of clearness and
       exactitude which has been attained in the visible structure of
       the heavenly bodies. It is not many years ago since the law of
       Avogadro-Gerhardt took root in science. It is within the memory
       of many living scientific men, and of mine amongst others. It is
       not surprising, therefore, that as yet little progress has been
       made in the province of molecular mechanics; but the theory of
       gases alone, which is intimately connected with the conception of
       molecules, shows by its success that the time is approaching when
       our knowledge of the internal structure of matter will be defined
       and established.

_Atoms are the smallest quantities_ or chemically indivisible masses _of
the elements forming the molecules_ of elements and compounds.

Atoms have weight, the sum of their weights forms the weight of the
molecule, and the sum of the weights of the molecules forms the weight
of masses, and is the cause of gravity, and of all the phenomena which
depend on the mass of a substance.

The elements are characterised, not only by their independent existence,
their incapacity of being converted into each other, &c., but also by
the weight of their atoms.

Chemical and physical properties depend on the weight, composition, and
properties of the molecules forming a substance, and on the weight and
properties of the atoms forming the molecules.

This is the substance of those principles of molecular mechanics which
lie at the basis of all contemporary physical and chemical constructions
since the establishment of the law of Avogadro-Gerhardt. The fecundity
of the principles enunciated is seen at every step in all the particular
cases forming the present store of chemical data. We will here cite a
few examples of the application of the law.

As the weight of an atom must be understood as the minimum quantity of
an element entering into the composition of all the molecules formed
by it, therefore, in order to find the weight of an atom of oxygen,
let us take the molecules of those of its compounds which have already
been described, together with the molecules of certain of those carbon
compounds which will be described in the following chapter:

                 Molecular  Amount of          Molecular  Amount of
                   Weight     Oxygen             Weight     Oxygen

         H_{2}O     18         16      HNO_{3}    63         48
         N_{2}O     44         16      CO         28         16
         NO         30         16      CO_{2}     44         32
         NO_{2}     46         32

The number of substances taken might be considerably increased, but
the result would be the same--that is, the molecules of the compounds
of oxygen would never be found to contain less than 16 parts by weight
of this element, but always _n_16, where _n_ is a whole number. The
molecular weights of the above compounds are found either directly from
the density of their vapour or gas, or from their reactions. Thus, the
vapour density of nitric acid (as a substance which easily decomposes
above its boiling point) cannot be accurately determined, but the
fact of its containing one part by weight of hydrogen, and all its
properties and reactions, indicate the above molecular composition and
no other. In this manner it is very easy to find the atomic weight of
all the elements, knowing the molecular weight and composition of their
compounds. It may, for instance, be easily proved that less than _n_12
parts of carbon never enters into the molecules of carbon compounds,
and therefore C must be taken as 12, and not as 6 which was the number
in use before Gerhardt. In a similar manner the atomic weights now
accepted for the elements oxygen, nitrogen, carbon, chlorine, sulphur,
&c., were found and indubitably established, and they are even now
termed the Gerhardt atomic weights. As regards the metals, many of
which do not give a single volatile compound, we shall afterwards
see that there are also methods by which their atomic weights may be
established, but nevertheless the law of Avogadro-Gerhardt is here
also ultimately resorted to, in order to remove any doubt which may be
encountered. Thus, for instance, although much that was known concerning
the compounds of beryllium necessitated its atomic weight being taken
as Be = 9--that is, the oxide as BeO and the chloride BeCl_{2}--still
certain analogies gave reason for considering its atomic weight to be Be
= 13·5, in which case its oxide would be expressed by the composition
Be_{2}O_{3}, and the chloride by BeCl_{3}.[21] It was then found that
the vapour density of beryllium chloride was approximately 40, when
it became quite clear that its molecular weight was 80, and as this
satisfies the formula BeCl_{2}, but does not suit the formula BeCl_{3},
it therefore became necessary to regard the atomic weight of Be as 9 and
not as 13-1/2.

  [21] If Be = 9, and beryllium chloride be BeCl_{2}, then for every 9
       parts of beryllium there are 71 parts of chlorine, and the
       molecular weight of BeCl_{2} = 80; hence the vapour density
       should be 40 or _n_{4}0. If Be = 13·5, and beryllium chloride be
       BeCl_{3}, then to 13·5 of beryllium there are 106·5 of chlorine;
       hence the molecular weight would be 120, and the vapour density
       60 or _n_60. The composition is evidently the same in both cases,
       because 9 : 71 :: 13·5 : 106·5. Thus, if the symbol of an element
       designate different atomic weights, apparently very different
       formulæ may equally well express both the percentage composition
       of compounds, and those properties which are required by the laws
       of multiple proportions and equivalents. The chemists of former
       days accurately expressed the composition of substances, and
       accurately applied Dalton's laws, by taking H = 1, O = 8, C = 6,
       Si = 14, &c. The Gerhardt equivalents are also satisfied by them,
       because O = 16, C = 12, Si = 28, &c., are multiples of them.
       The choice of one or the other multiple quantity for the atomic
       weight is impossible without a firm and concrete conception of
       the molecule and atom, and this is only obtained as a consequence
       of the law of Avogadro-Gerhardt, and hence the modern atomic
       weights are the results of this law (_see_ Note 28).

With the establishment of a true conception of molecules and atoms,
chemical formulæ became direct expressions, not only of composition,[22]
but also of molecular weight or _vapour density_, and consequently of a
series of fundamental chemical and physical data, inasmuch as a number
of the properties of substances are dependent on their vapour density,
or molecular weight and composition. The vapour density D = M/2. For
instance, the formula of ethyl ether is C_{4}H_{10}O, corresponding
with the molecular weight 74, and the vapour density 37, which is the
fact. Therefore, the density of vapours and gases has ceased to be
an empirical magnitude obtained by experiment only, and has acquired
a rational meaning. It is only necessary to remember that 2 grams
of hydrogen, or the molecular weight of this primary gas in grams,
occupies, at 0° and 760 mm. pressure, a volume of 22·3 litres (or
22,300 cubic centimetres), in order to directly determine the weights
of cubical measures of gases and vapours from their formulæ, because
_the molecular weights in grams of all other vapours at 0° and 760 mm.
occupy the same volume, 22·3 litres_. Thus, for example, in the case
of carbonic anhydride, CO_{2}, the molecular weight M = 44, hence 44
grams of carbonic anhydride at 0° and 760 mm. occupy a volume of 22·3
litres--consequently, a litre weighs 1·97 gram. By combining the laws of
gases--Gay-Lussac's, Mariotte's, and Avogadro-Gerhardt's--we obtain[23]
a general formula for gases

                        6200_s_(273 + _t_) = M_p_

where _s_ is the weight in grams of a cubic centimetre of a vapour or
gas at a temperature _t_ and pressure _p_ (expressed in centimetres of
mercury) if the molecular weight of the gas = M. Thus, for instance, at
100° and 760 millimetres pressure (_i.e._ at the atmospheric pressure)
the weight of a cubic centimetre of the vapour of ether (M = 74) is _s_
= 0·0024.[24]

  [22] The percentage amounts of the elements contained in a given
       compound may be calculated from its formula by a simple
       proportion. Thus, for example, to find the percentage amount of
       hydrogen in hydrochloric acid we reason as follows:--HCl shows
       that hydrochloric acid contains 35·5 of chlorine and 1 part of
       hydrogen. Hence, in 36·5 parts of hydrochloric acid there is 1
       part by weight of hydrogen, consequently 100 parts by weight of
       hydrochloric acid will contain as many more units of hydrogen as
       100 is greater than 36·5; therefore, the proportion is as
       follows--_x_ : 1 :: 100 : 36·5 or _x_ = 100/36·5 = 2·739.
       Therefore 100 parts of hydrochloric acid contain 2·739 parts of
       hydrogen. In general, when it is required to transfer a formula
       into its percentage composition, we must replace the symbols by
       their corresponding atomic weights and find their sum, and knowing
       the amount by weight of a given element in it, it is easy by
       proportion to find the amount of this element in 100 or any other
       quantity of parts by weight. If, on the contrary, it be required
       to find the formula from a given percentage composition, we must
       proceed as follows: Divide the percentage amount of each element
       entering into the composition of a substance by its atomic weight,
       and compare the figures thus obtained--they should be in simple
       multiple proportion to each other. Thus, for instance, from the
       percentage composition of hydrogen peroxide, 5·88 of hydrogen and
       94·12 of oxygen, it is easy to find its formula; it is only
       necessary to divide the amount of hydrogen by unity and the amount
       of oxygen by 16. The numbers 5·88 and 5·88 are thus obtained,
       which are in the ratio 1 : 1, which means that in hydrogen
       peroxide there is one atom of hydrogen to one atom of oxygen.

       The following is a proof of the practical rule given above _that
       to find the ratio of the number of atoms from the percentage
       composition, it is necessary to divide the percentage amounts
       by the atomic weights of the corresponding substances, and to
       find the ratio which these numbers bear to each other_. Let us
       suppose that two radicles (simple or compound), whose symbols
       and combining weights are A and B, combine together, forming
       a compound composed of _x_ atoms of A and _y_ atoms of B. The
       formula of the substance will be A_x_B_y_. From this formula we
       know that our compound contains _x_A parts by weight of the first
       element, and _y_B of the second. In 100 parts of our compound
       there will be (by proportion) (100. _x_A)/(_x_A + _y_B) of the
       first element, and (100. _y_B)/(_x_A + _y_B) of the second. Let
       us divide these quantities, expressing the percentage amounts by
       the corresponding combining weights; we then obtain 100_x_/(_x_A
       + _y_B) for the first element and 100_y_/(_x_A + _y_B) for the
       second element. And these numbers are in the ratio _x_ : _y_--that
       is, in the ratio of the number of atoms of the two substances.

       It may be further observed that even the very language or
       nomenclature of chemistry acquires a particular clearness and
       conciseness by means of the conception of molecules, because then
       the names of substances may directly indicate their composition.
       Thus the term 'carbon dioxide' tells more about and expresses
       CO_{2} better than carbonic acid gas, or even carbonic anhydride.
       Such nomenclature is already employed by many. But expressing
       the composition without an indication or even hint as to the
       properties, would be neglecting the advantageous side of the
       present nomenclature. Sulphur dioxide, SO_{2}, expresses the same
       as barium dioxide, BaO_{2}, but sulphurous anhydride indicates
       the acid properties of SO_{2}. Probably in time one harmonious
       chemical language will succeed in embracing both advantages.

  [23] This formula (which is given in my work on 'The Tension of
       Gases,' and in a somewhat modified form in the 'Comptes Rendus,'
       Feb. 1876) is deduced in the following manner. According to the
       law of Avogadro-Gerhardt, M = 2D for all gases, where M is the
       molecular weight and D the density referred to hydrogen. But it is
       equal to the weight _s__{0} of a cubic centimetre of a gas in
       grams at 0° and 76 cm. pressure, divided by 0·0000898, for this is
       the weight in grams of a cubic centimetre of hydrogen. But the
       weight _s_ of a cubic centimetre of a gas at a temperature _t_ and
       under a pressure _p_ (in centimetres) is equal to _s__{0}_p_/76(1
       + _at_). Therefore, _s__{0} = _s_.76(1 + _at_)/_p_; hence D =
       76._s_(1 + _at_)/0·0000898_p_, whence M = 152_s_(1 +
       _at_)/0·0000898_p_, which gives the above expression, because
       1/_a_ = 273, and 152 multiplied by 273 and divided by 0·0000898 is
       nearly 6200. In place of _s_, _m/v_ may be taken, where _m_ is the
       weight and _v_ the volume of a vapour.

  [24] The above formula may be directly applied in order to ascertain
       the molecular weight from the data; weight of vapour _m_ grms.,
       its volume _v_ c.c., pressure _p_ cm., and temperature _t_°; for
       _s_ = the weight of vapour _m_, divided by the volume _v_, and
       consequently M = 6,200_m_(273 + _t_)/_pv_. Therefore, instead
       of the formula (_see_ Chapter II., Note 34), _pv_ = R(273 +
       _t_), where R varies with the mass and nature of a gas, we
       may apply the formula _pv_ = 6,200(_m_/M)(273 + _t_). These
       formulæ simplify the calculations in many cases. For example,
       required the volume _v_ occupied by 5 grms. of aqueous vapour
       at a temperature _t_ = 127° and under a pressure _p_ = 76 cm.
       According to the formula M = 6,200_m_(273 + _t_)/_pv_, we find
       that _v_ = 9,064 c.c., as in the case of water M = 18, _m_ in
       this instance = 5 grms. (These formulæ, however, like the laws of
       gases, are only approximate.)

As the molecules of many elements (hydrogen, oxygen, nitrogen, chlorine,
bromine, sulphur--at least at high temperatures) are of uniform
composition, the formulæ of the compounds formed by them directly
indicate the composition by volume. So, for example, the formula HNO_{3}
directly shows that in the decomposition of nitric acid there is
obtained 1 vol. of hydrogen, 1 vol. of nitrogen, and 3 vols. of oxygen.

And since a great number of mechanical, physical, and chemical
properties are directly dependent on the elementary and volumetric
composition, and on the vapour density, the accepted system of atoms
and molecules gives the possibility of simplifying a number of most
complex relations. For instance, it may be easily demonstrated _that
the vis viva of the molecules of all vapours and gases is alike_. For
it is proved by mechanics that the _vis viva_ of a moving mass = (1/2)
_mv_^2, where _m_ is the mass and _v_ the velocity. For a molecule,
_m_ = M, or the molecular weight, and the velocity of the motion of
gaseous molecules = a constant which we will designate by C, divided
by the square root of the density of the gas[25] = C/[sqrt]D, and as D
= M/2, the _vis viva_ of molecules = C^2--that is, a constant for all
molecules. _Q.E.D._[26] The specific heat of gases (Chapter XIV.), and
many other of their properties, are determined by their density, and
consequently by their molecular weight. Gases and vapours in passing
into a liquid state evolve the so-called _latent heat_, which also
proves to be in connection with the molecular weight. The observed
latent heats of carbon bisulphide, CS_{2} = 90, of ether, C_{4}H_{10}O,
= 94, of benzene, C_{6}H_{6}, = 109, of alcohol, C_{2}H_{6}O, = 200,
of chloroform, CHCl_{3}, = 67, &c., show the amount of heat expended
in converting one part by weight of the above substances into vapour.
A great uniformity is observed if the measure of this heat he referred
to the weight of the molecule. For carbon bisulphide the formula CS_{2}
expresses a weight 76, hence the latent heat of evaporation referred to
the molecular quantity CS_{2} = 76 x 90 = 6,840, for ether = 9,656, for
benzene = 8,502, for alcohol = 9,200, for chloroform = 8,007, for water
= 9,620, &c. That is, for molecular quantities, the latent heat varies
comparatively little, from 7,000 to 10,000 heat units, whilst for equal
parts by weight it is ten times greater for water than for chloroform
and many other substances.[27]

  [25] Chapter I., Note 34.

  [26] _The velocity of the transmission of sound through gases and
       vapours_ closely bears on this. It = [sqrt](_Kpg_)/D(1 +
       [Greek: a]_t_), where _K_ is the ratio between the two specific
       heats (it is approximately 1·4 for gases containing two atoms in a
       molecule), _p_ the pressure of the gas expressed by weight (that
       is, the pressure expressed by the height of a column of mercury
       multiplied by the density of mercury), _g_ the acceleration of
       gravity, D the weight of a cubic measure of the gas, [Greek: a] =
       0·00367, and _t_ the temperature. Hence, if _K_ be known, and as D
       can he found from the composition of a gas, we can calculate the
       velocity of the transmission of sound in that gas. Or if this
       velocity be known, we can find _K_. The relative velocities of
       sound in two gases can he easily determined (Kundt).

       If a horizontal glass tube (about 1 metre long and closed at
       both ends) be full of a gas, and be firmly fixed at its middle
       point, then it is easy to bring the tube and gas into a state
       of vibration, by rubbing it from centre to end with a damp
       cloth. The vibration of the gas is easily rendered visible, if
       the interior of the tube be dusted with lycopodium (the yellow
       powder-dust or spores of the lycopodium plant is often employed
       in medicine), before the gas is introduced and the tube fused
       up. The fine lycopodium powder arranges itself in patches, whose
       number depends on the velocity of sound in the gas. If there
       be 10 patches, then the velocity of sound in the gas is ten
       times slower than in glass. It is evident that this is an easy
       method of comparing the velocity of sound in gases. It has been
       demonstrated by experiment that the velocity of sound in oxygen
       is four times less than in hydrogen, and the square roots of the
       densities and molecular weights of hydrogen and oxygen stand in
       this ratio.

  [27] If the conception of the molecular weights of substances does not
       give an exact law when applied to the latent heat of evaporation,
       at all events it brings to light a certain uniformity in
       figures, which otherwise only represent the simple result of
       observation. Molecular quantities of liquids appear to expend
       almost equal amounts of heat in their evaporation. It may be said
       that the latent heat of evaporation of molecular quantities is
       approximately constant, because the _vis viva_ of the motion of
       the molecules is, as we saw above, a constant quantity. According
       to thermodynamics the latent heat of evaporation is equal to
       ((_t_ + 273)/E)(_n_´-_n_)_dp_/_d_T × 13·59, where _t_ is the
       boiling point, _n_´ the specific volume (_i.e._ the volume of a
       unit of weight) of the vapour, and _n_ the specific volume of
       the liquid, _dp_/_d_T the variation of the tension with a rise
       of temperature per 1°, and 13·89 the density of the mercury
       according to which the pressure is measured. Thus the latent heat
       of evaporation increases not only with a decrease in the vapour
       density (_i.e._ the molecular weight), but also with an increase
       in the boiling point, and therefore depends on different factors.

Generalising from the above, the weight of the molecule determines the
properties of a substance _independently of its composition_--_i.e._ of
the number and quality of the atoms entering into the molecule--whenever
the substance is in a gaseous state (for instance, the density of
gases and vapours, the velocity of sound in them, their specific heat,
&c.), or passes into that state, as we see in the latent heat of
evaporation. This is intelligible from the point of view of the atomic
theory in its present form, for, besides a rapid motion proper to the
molecules of gaseous bodies, it is further necessary to postulate that
these molecules are dispersed in space (filled throughout with the
luminiferous ether) like the heavenly bodies distributed throughout
the universe. Here, as there, it is only the degree of removal (the
distance) and the masses of substances which take effect, while
those peculiarities of a substance which are expressed in chemical
transformations, and only come into action on near approach or on
contact, are in abeyance by reason of the dispersal. Hence it is at once
obvious, in the first place, that in the case of solids and liquids,
in which the molecules are closer together than in gases and vapours,
a greater complexity is to be expected, _i.e._ a dependence of all
the properties not only upon the weight of the molecule but also upon
its composition and quality, or upon the properties of the individual
chemical atoms forming the molecule; and, in the second place, that,
in the case of a small number of molecules of any substance being
disseminated through a mass of another substance--for example, in the
formation of weak (dilute) solutions (although in this case there is
an act of chemical reaction--_i.e._ a combination, decomposition, or
substitution)--the dispersed molecules will alter the properties of
the medium in which they are dissolved, almost in proportion to the
molecular weight and almost independently of their composition. The
greater the number of molecules disseminated--_i.e._ the stronger the
solution--the more clearly defined will those properties become which
depend upon the composition of the dissolved substance and its relation
to the molecules of the solvent, for the distribution of one kind of
molecules in the sphere of attraction of others cannot but be influenced
by their mutual chemical reaction. These general considerations give
a starting point for explaining why, since the appearance of Van't
Hoff's memoir (1886), 'The Laws of Chemical Equilibrium in a Diffused
Gaseous or Liquid State' (_see_ Chapter I., Note 19), it has been found
more and more that _dilute_ (weak) solutions exhibit such variations
of properties as depend wholly upon the weight and number of the
molecules and not upon their composition, and even give the means of
determining the weight of molecules by studying the variations of the
properties of a solvent on the introduction of a small quantity of a
substance passing into solution. Although this subject has been already
partially considered in the first chapter (in speaking of solutions),
and properly belongs to a special (physical) branch of chemistry, we
touch upon it here because the meaning and importance of molecular
weights are seen in it in a new and peculiar light, and because it gives
a method for determining them whenever it is possible to obtain dilute
solutions. Among the numerous properties of dilute solutions which have
been investigated (for instance, the osmotic pressure, vapour tension,
boiling point, internal friction, capillarity, variation with change of
temperature, specific heat, electroconductivity, index of refraction,
&c.) we will select one--the 'depression' or fall of the temperature of
freezing (Raoult's cryoscopic method), not only because this method has
been the most studied, but also because it is the most easily carried
out and most frequently applied for determining the weight of the
molecules of substances in solution, although here, owing to the novelty
of the subject there are also many experimental discrepancies which
cannot as yet be explained by theory.[27 bis]

  [27 bis] The osmotic pressure, vapour tension of the solvent, and
       several other means applied like the cryoscopic method to dilute
       solutions for determining the molecular weight of a substance in
       solution, are more difficult to carry out in practice, and only
       the method of _determining the rise of the boiling point_ of
       dilute solutions can from its facility be placed parallel with the
       cryoscopic method, to which it bears a strong resemblance, as in
       both the solvent changes its state and is partially separated. In
       the boiling point method it passes off in the form of a vapour,
       while in cryoscopic determinations it separates out in the form of
       a solid body.

       Van't Hoff, starting from the second law of thermodynamics, showed
       that the dependence of the rise of pressure (_dp_) upon a rise of
       temperature (_d_T) is determined by the equation _dp_ =
       (_kmp_/2T^2)_d_T, where _k_ is the latent heat of evaporation of
       the solvent, _m_ its molecular weight, _p_ the tension of the
       saturated vapour of the solvent at T, and T the absolute
       temperature (T = 273 + _t_), while Raoult found that the quantity
       (_p_-_p´_)/_p_ (Chapter I., Note 50) or the measure of the
       relative fall of tension (_p_ the tension of the solvent or water,
       and _p´_ of the solution) is found by the ratio of the number of
       molecules, _n_ of the substance dissolved, and N of the solvent,
       so that (_p_-_p´_)/_p_ = C_n_/(N + _n_) where C is a constant.
       With very dilute solutions _p_ _-p´_ may be taken as equal to
       _dp_, and the fraction _n_/(N + _n_) as equal to _n_/N (because in
       that case the value of N is very much greater than _n_), and then,
       judging from experiment, C is nearly unity--hence: _dp/p_ = _n_/N
       or _dp_ = _np_/N, and on substituting this in the above equation
       we have (_kmp_/2T^2)_d_T = _np_/N. Taking a weight of the solvent
       _m_/N = 100, and of the substance dissolved (per 100 of the
       solvent) _q_, where _q_ evidently = _n_M, if M be the molecular
       weight of the substance dissolved, we find that _n_/N = _qm_/100M,
       and hence, according to the preceding equation, we have M =
       (0·02T^2/_k_)·(_q_/_d_T), that is, by taking a solution of _q_
       grms. of a substance in 100 grms. of a solvent, and determining by
       experiment the rise of the boiling point _d_T, we find the
       molecular weight M of the substance dissolved, because the
       fraction 0·02T^2/_k_ is (for a given pressure and solvent) a
       constant; for water at 100° (T = 373°) when _k_ = 534 (Chapter I.,
       Note 11), it is nearly 5·2, for ether nearly 21, for bisulphide of
       carbon nearly 24, for alcohol nearly 11·5, &c. As an example, we
       will cite from the determinations made by Professor Sakurai, of
       Japan (1893), that when water was the solvent and the substance
       dissolved, corrosive sublimate, HgCl_{2}, was taken in the
       quantity _q_ = 8·978 and 4·253 grms., the rise in the boiling
       point _d_T was = O°·179 and 0°·084, whence M = 261 and 263, and
       when alcohol was the solvent, _q_ = 10·873 and 8·765 and _d_T =
       0°·471 and 0°·380, whence M = 266 and 265, whilst the actual
       molecular weight of corrosive sublimate = 271, which is very near
       to that given by this method. In the same manner for aqueous
       solutions of sugar (M = 342), when _q_ varied from 14 to 2·4, and
       the rise of the boiling point from 0°·21 to 0°·035, M was found to
       vary between 339 and 364. For solutions of iodine I_{2} in ether,
       the molecular weight was found by this method to be between 255
       and 262, and I_{2} = 254. Sakurai obtained similar results
       (between 247 and 262) for solutions of iodine in bisulphide of
       carbon.

       We will here remark that in determining M (the molecular weight of
       the substance dissolved) at small but increasing concentrations
       (per 100 grms. of water), the results obtained by Julio Baroni
       (1893) show that the value of M found by the formula may either
       increase or decrease. An increase, for instance, takes place in
       aqueous solutions of HgCl_{2} (from 255 to 334 instead of 271),
       KNO_{3} (57-66 instead of 101), AgNO_{3} (104-107 instead of 170),
       K_{2}SO_{4} (55-89 instead of 174), sugar (328-348 instead of
       342), &c. On the contrary the calculated value of M decreases as
       the concentration increases, for solutions of KCl (40-39 instead
       of 74·5), NaCl (33-28 instead of 58·5), NaBr (60-49 instead of
       103), &c. In this case (as also for LiCl, NaI, C_{2}H_{3}NaO_{2},
       &c.) the value of _i_ (Chapter I., Note 49), or the ratio between
       the actual molecular weight and that found by the rise of the
       boiling point, was found to increase with the concentration,
       _i.e._ to be greater than 1, and to differ more and more from
       unity as the strength of the solution becomes greater. For
       example, according to Schlamp (1894), for LiCl, with a variation
       of from 1·1 to 6·7 grm. LiCl per 100 of water, _i_ varies from
       1·63 to 1·89. But for substances of the first series (HgCl_{2},
       &c.), although in very dilute solutions _i_ is greater than 1, it
       approximates to 1 as the concentration increases, and this is the
       normal phenomenon for solutions which do not conduct an electric
       current, as, for instance, of sugar. And with certain
       electrolytes, such as HgCl_{2}, MgSO_{4}, &c., _i_ exhibits a
       similar variation; thus, for HgCl_{2} the value of M is found to
       vary between 255 and 334; that is, _i_ (as the molecular weight =
       271) varies between 1·06 and 0·81. Hence I do not believe that the
       difference between _i_ and unity (for instance, for CaCl_{2}, _i_
       is about 3, for KI about 2, and decreases with the concentration)
       can at present be placed at the basis of any general chemical
       conclusions, and it requires further experimental research. Among
       other methods by which the value of _i_ is now determined for
       dilute solutions is the study of their electroconductivity,
       admitting that _i_ = 1 + _a_(_k_-1), where _a_ = the ratio of the
       molecular conductivity to the limiting conductivity corresponding
       to an infinitely large dilution (_see_ Physical Chemistry), and
       _k_ is the number of ions into which the substance dissolved can
       split up. Without entering upon a criticism of this method of
       determining _i_, I will only remark that it frequently gives
       values of _i_ very close to those found by the depression of the
       freezing point and rise of the boiling point; but that this
       accordance of results is sometimes very doubtful. Thus for a
       solution containing 5·67 grms. CaCl_{2} per 100 grms. of water,
       _i_, according to the vapour tension = 2·52, according to the
       boiling point = 2·71, according to the electroconductivity = 2·28,
       while for solutions in propyl alcohol (Schlamp 1894) _i_ is near
       to 1·33. In a word, although these methods of determining the
       molecular weight of substances in solution show an undoubted
       progress in the general chemical principles of the molecular
       theory, there are still many points which require explanation.

       We will add certain general relations which apply to these
       problems. Isotonic (Chapter I., Note 19) solutions exhibit not
       only similar osmotic pressures, but also the same vapour tension,
       boiling point and freezing temperature. The osmotic pressure bears
       the same relation to the fall of the vapour tension as the
       specific gravity of a solution does to the specific gravity of the
       vapour of the solvent. The general formulæ underlying the whole
       doctrine of the influence of the molecular weight upon the
       properties of solutions considered above, are: 1. Raoult in
       1886-1890 showed that

             ((_p_-_p_´)/_p_) × (100/_a_) × (M/_m_) = a constant C

       where _p_ and _p_´ are the vapour tensions of the solvent and
       substance dissolved, _a_ the amount in grms. of the substance
       dissolved per 100 grms. of solvent, M and _m_ the molecular
       weights of the substance dissolved and solvent. 2. Raoult and
       Recoura in 1890 showed that the constant above C = the ratio of
       the actual vapour density _d_´ of the solvent to the theoretical
       density _d_ calculated according to the molecular weight. This
       deduction may now be considered proved, because both the fall of
       tension and the ratio of the vapour densities _d_´/_d_ give, for
       water 1·03, for alcohol 1·02, for ether 1·04, for bisulphide of
       carbon 1·00, for benzene 1·02, for acetic acid 1·63. 3. By
       applying the principles of thermodynamics and calling L_{1} the
       latent heat of fusion and T_{1} the absolute (= _t_ + 273)
       temperature of fusion of the solvent, and L_{2} and T_{2} the
       corresponding values for the boiling point, Van't Hoff in
       1886-1890 deduced:--

         (Depression of freezing point)/(Rise of boiling point)
                                    = (L_{2}/L_{1}) × (T_{1}^2/T_{2}^2)

          Depression of freezing point = (AT_{1}^{2}_a_)/(L_{1}M_{1})

              Rise of boiling point = (AT_{2}^{2}_a_)/(L_{2}M_{1})

       where A = 0·01988 (or nearly 0·02 as we took it above), _a_ is
       the weight in grms. of the substance dissolved per 100 grms. of
       the solvent, M_{1} the molecular weight of the dissolved substance
       (in the solution), and M the molecular weight of this substance
       according to its composition and vapour density, then _i_ =
       M/M_{1}. The experimental data and theoretical considerations upon
       which these formulæ are based will be found in text-books of
       physical and theoretical chemistry.

If 100 gram-molecules of water, _i.e._ 1,800 grms, be taken and _n_
gram-molecules of sugar, C_{12}H_{22}O_{11}, _i.e._ _n_ 342 grms., be
dissolved in them, then the depression _d_, or fall (counting from
0°) of the temperature of the formation of ice will be (according to
Pickering)

          _n_ = 0    0·010    0·025    0·100    0·250    1·000
          _d_ = 0°  0°·0103  0°·0280  0°·1115  0°·2758  1°·1412

which shows that for high degrees of dilution (up to 0·25_n_) _d_
approximately (estimating the possible errors of experiment at ±0°·005)
= _n_1·10, because then _d_ = 0°, 0°·0110, 0°·0275, 0°·1100, 0°·2750,
1°·1000, and the difference between these figures and the results
of experiment for very dilute solutions is less than the possible
errors of experiment (for _n_ = 1 the difference is already greater)
and therefore for dilute solutions of sugar it may be said that _n_
molecules of sugar in dissolving in 100 molecules of water give a
depression of about 1°·1_n_. Similar data for acetone (Chapter I., Note
49) give a depression of 1°·006_n_ for _n_ molecules of acetone per 100
molecules of water. And in general, for indifferent substances (the
majority of organic bodies) the depression per 100H_{2}O is _nearly
n_1°·1 to _n_1°·0 (ether, for instance, gives the last number), and
consequently in dissolving in 100 grms. of water it is about 18°·0_n_
to 19°·0_n_, taking this rule to apply to the case of a small number
of _n_ (not over 0·2_n_). If instead of water, other liquid or fused
solvents (for example, benzene, acetic acid, acetone, nitrobenzene or
molten naphthaline, metals, &c.) be taken and in the proportion of 100
molecules of the solvent to _n_ molecules of a dissolved indifferent
(neither acid nor saline) substance, then the depression is found to be
equal to from 0°·62_n_ to 0°·65_n_ and in general K_n_. If the molecular
weight of the solvent = _m_, then 100 gram-molecules will weigh 100_m_
grms., and the depression will be approximately (taking 0·63_n_) equal
to _m_0·63_n_ degrees for _n_ molecules of the substance dissolved in
100 grms. of the solvent, or in general the depression for 100 grms.
of a given solvent = _kn_ where _k_ is almost a constant quantity (for
water nearly 18, for acetone nearly 37, &c.) for all dilute solutions.
Thus, having found a convenient solvent for a given substance and
prepared a definite (by weight) solution (_i.e._ knowing how many grms.
_r_ of the solvent there are to _q_ grms. of the substance dissolved)
and having determined the depression _d_--_i.e._ the fall in temperature
of freezing for the solvent--it is possible to determine the molecular
weight of the substance dissolved, because _d_ = _kn_ where _d_ is found
by experiment and _k_ is determined by the nature of the solvent, and
therefore _n_ or the number of molecules of the substance dissolved can
be found. But if _r_ grms. of the solvent and _q_ grms. of the substance
dissolved are taken, then there are 100_q_/_r_ of the latter per 100
grms. of the former, and this quantity = _n_X, where _n_ is found from
the depression and = _d_/_k_ and X is the molecular weight of the
substance dissolved. Hence X = 100_qk_/_rd_, which gives the molecular
weight, naturally only approximately, but still with sufficient accuracy
to easily indicate, for instance, whether in peroxide of hydrogen the
molecule contains HO or H_{2}O_{2} or H_{3}O_{3}, &c. (H_{2}O_{2} is
obtained). Moreover, attention should be drawn to the fact that a great
many substances taken as solvents give per 100 molecules a depression of
about 0·63_n_, whilst water gives about 1·05_n_, _i.e._ a larger
quantity, as though the molecules of liquid water were more complex than
is expressed by the formula H_{2}O.[28] A similar phenomenon which
repeats itself in the osmotic pressure, vapour tension of the solvent,
&c. (_see_ Chapter I., Notes 19 and 49), _i.e._ a variation of the
constant (_k_ for 100 grms. of the solvent or K for 100 molecules of it),
is also observed in passing from indifferent substances to saline (to
acids, alkalis and salts) both in aqueous and other solutions as we will
show (according to Pickering's data 1892) for solutions of NaCl and
CuSO_{4} in water. For

                    _n_ = 0·01  0·03  0·05  0·1  0·5

molecules of NaCl the depression is

            _d_ = 0°·0177  0°·0598  0°·0992  0°·1958  0°·9544

which corresponds to a depression per molecule

                    K = 1·77  1·96  1·98  1·96  1·91

_i.e._ here in the most dilute solutions (when _n_ is nearly 0) _d_ is
obtained about 1·7_n_, while in the case of sugar it was about 1·1_n_.
For CuSO_{4} for the same values of _n_, experiment gave:

            _d_ = 0°·0164  0°·0451  0°·0621  0°·1321  0°·5245
              K =  1·64     1·50     1·44     1·32     1·05

_i.e._ here again _d_ for very dilute solutions is nearly 1·7_n_, but
the value of K falls as the solution becomes more concentrated, while
for NaCl it at first increased and only fell for the more concentrated
solutions. The value of K in the solution of _n_ molecules of a body
in 100H_{2}O, when _d_ = K_n_, for very dilute solutions of CaCl_{2}
is nearly 2·6, for Ca(NO_{3})_{2} nearly 2·5, for HNO_{3}, KI and KHO
nearly 1·9-2·O, for borax Na_{2}B_{4}O_{7} nearly 3·7, &c., while for
sugar and similar substances it is, as has been already mentioned,
nearly 1·0-1·1. Although these figures are very different[28 bis] still
_k_ and K may be considered constant for analogous substances, and
therefore the weight of the molecule of the body in solution can be
found from _d_. And as the vapour tension of solutions and their boiling
points (_see_ Note 27 bis and Chapter I., Note 51) vary in the same
manner as the freezing point depression, so they also may serve as means
for determining the molecular weight of a substance in solution.[29]

  [28] A similar conclusion respecting the molecular weight of liquid
       water (_i.e._ that its molecule in a liquid state is more
       complex than in a gaseous state, or polymerized into H_{8}O_{4},
       H_{6}O_{3} or in general into _n_H_{2}O) is frequently met
       in chemico-physical literature, but as yet there is no basis
       for its being fully admitted, although it is possible that a
       polymerization or aggregation of several molecules into one
       takes place in the passage of water into a liquid or solid
       state, and that there is a converse depolymerization in the
       act of evaporation. Recently, particular attention has been
       drawn to this subject owing to the researches of Eötvös (1886)
       and Ramsay and Shields (1893) on the variation of the surface
       tension N with the temperature (N = the capillary constant _a_^2
       multiplied by the specific gravity and divided by 2, for example,
       for water at 0° and 100° the value of _a_^2 = 15·41 and 12·58
       sq. mm., and the surface tension 7·92 and 6·04). Starting from
       the absolute boiling point (Chapter II., Note 29) and adding 6°,
       as was necessary from all the data obtained, and calling this
       temperature T, it is found that AS = _k_T, where S is the surface
       of a gram-molecule of the liquid (if M is its weight in grams,
       _s_ its sp. gr., then its sp. volume = M/_s_, and the surface
       S = [3root](M/_s_)^2), A the surface tension (determined by
       experiment at T), and _k_ a constant which is independent of
       the composition of the molecule. The equation AS = _k_T is in
       complete agreement with the well-known equation for gases _vp_
       = RT (p. 140) which serves for deducing the molecular weight
       from the vapour density. Ramsay's researches led him to the
       conclusion that the liquid molecules of CS_{2}, ether, benzene,
       and of many other substances, have the same value as in a state
       of vapour, whilst with other liquids this is not the case, and
       that to obtain an accordance, that is, that _k_ shall be a
       constant, it is necessary to assume the molecular weight in the
       liquid state to be _n_ times as great. For the fatty alcohols and
       acids _n_ varies from 1-1/2 to 3-1/2, for water from 2-1/4 to 4,
       according to the temperature (at which the depolymerization takes
       place). Hence, although this subject offers a great theoretical
       interest, it cannot be regarded as firmly established, the more
       so since the fundamental observations are difficult to make and
       not sufficiently numerous; should, however, further experiments
       confirm the conclusions arrived at by Professor Ramsay, this will
       give another method of determining molecular weights.

  [28 bis] Their variance is expressed in the same manner as was done by
       Van't Hoff (Chapter I., Notes 19 and 49) by the quantity _i_,
       taking it as = 1 when _k_ = 1·05, in that case for KI, _i_ is
       nearly 2, for borax about 4, &c.

  [29] We will cite one more example, showing the direct dependence of
       the properties of a substance on the molecular weight. If one
       molecular part by weight of the various chlorides--for instance,
       of sodium, calcium, barium, &c.--be dissolved in 200 molecular
       parts by weight of water (for instance, in 3,600 grams) then
       it is found that the greater the molecular weight of the salt
       dissolved, the greater is the specific gravity of the resultant
       solution.

                  Molecular    Sp. gr.                Molecular   Sp. gr.
                   weight       at 15°                 weight      at 15°
         HCl         36·5       1·0041       CaCl_{2}   111        1·0236
         NaCl        58·5       1·0106       NiCl_{2}   130        1·0328
         KCl         74·5       1·0121       ZnCl_{2}   136        1·0331
         BeCl_{2}    80         1·0138       BaCl_{2}   208        1·0489
         MgCl_{2}    95         1·0203

Thus not only in vapours and gases, but also in dilute solutions
of solid and liquid substances, we see that if not all, still many
properties are wholly dependent upon the molecular weight and not upon
the quality of a substance, and that this gives the possibility of
determining the weight of molecules by studying these properties (for
instance, the vapour density, depression of the freezing point, &c.) It
is apparent from the foregoing that the physical and even more so the
chemical properties of homogeneous substances, more especially solid and
liquid, do not depend exclusively upon the weights of their molecules,
but that many are in definite (_see_ Chapter XV.) dependence upon the
weights of the atoms of the elements entering into their composition,
and are determined by their quantitative and individual peculiarities.
Thus the density of solids and liquids (as will afterwards be shown) is
chiefly determined by the weights of the atoms of the elements entering
into their composition, inasmuch as dense elements (in a free state) and
compounds are only met with among substances containing elements with
large atomic weights, such as gold, platinum, and uranium. And these
elements themselves, in a free state, are the heaviest of all elements.
Substances containing such light elements as hydrogen, carbon, oxygen
and nitrogen (like many organic substances) never have a high specific
gravity; in the majority of cases it scarcely exceeds that of water. The
density generally decreases with the increase of the amount of hydrogen,
as the lightest element, and a substance is often obtained lighter than
water. The refractive power of substances also entirely depends on the
composition and the properties of the component elements.[29 bis] The
history of chemistry presents a striking example in point--Newton
foresaw from the high refractive index of the diamond that it would
contain a combustible substance since so many combustible oils have a
high refractive power. We shall afterwards see (Chapter XV.) that many
of those properties of substances which are in direct dependence not
upon the weight of the molecules but upon their composition, or, in
other words, upon the properties and quantities of the elements entering
into them, stand in a peculiar (periodic) dependence upon the atomic
weight of the elements; that is, the mass (of molecules and atoms),
proportional to the weight, determines the properties of substances
as it also determines (with the distance) the motions of the heavenly
bodies.

  [29 bis] With respect to the optical refractive power of substances, it
       must first be observed that the coefficient of refraction is
       determined by two methods: (_a_) either all the data are referred
       to one definite ray--for instance, to the Fraunhofer (sodium) line
       D of the solar spectrum--that is, to a ray of definite wave
       length, and often to that red ray (of the hydrogen spectrum) whose
       wave length is 656 millionths of a millimetre; (_b_) or Cauchy's
       formula is used, showing the relation between the coefficient of
       refraction and dispersion to the wave length _n_ = A +
       (B/([Greek: l]^2)), where A and B are two constants varying for
       every substance but constant for all rays of the spectrum, and
       [Greek: l] is the wave length of that ray whose coefficient of
       refraction is _n_. In the latter method the investigation usually
       concerns the magnitudes of A, which are independent of dispersion.
       We shall afterwards cite the data, investigated by the first
       method, by which Gladstone, Landolt, and others established the
       conception of the refraction equivalent.

       It has long been known that the _coefficient of refraction n_
       for a given substance decreases with the density of a substance D,
       so that the magnitude (_n_-1) ÷ D = C is almost constant for a
       given ray (having a definite wave length) and for a given
       substance. This constant is called the _refractive energy_, and
       its product with the atomic or molecular weight of a substance the
       _refraction equivalent_. The coefficient of refraction of oxygen
       is 1·00021, of hydrogen 1·00014, their densities (referred to
       water) are 0·00143 and 0·00009, and their atomic weights, O = 16,
       H = 1; hence their refraction equivalents are 3 and 1·5. Water
       contains H_{2}O, consequently the sum of the equivalents of
       refraction is (2 × 1·5) + 3 = 6. But as the coefficient of
       refraction of water = 1·331, its refraction equivalent = 5·958, or
       nearly 6. Comparison shows that, approximately, the sum of the
       refraction equivalents of the atoms forming compounds (or
       mixtures) is equal to the refraction equivalent of the compound.
       According to the researches of Gladstone, Landolt, Hagen, Brühl
       and others, the refraction equivalents of the elements are--H =
       1·3, Li = 3·8, B = 4·0, C = 5·0, N = 4·1 (in its highest state of
       oxidation, 5·3), O = 2·9, F = 1·4, Na = 4·8, Mg = 7·0, Al = 8·4,
       Si = 6·8, P = 18·3, S = 16·0, Cl = 9·9, K = 8·1, Ca = 10·4, Mn =
       12·2, Fe = 12·0 (in the salts of its higher oxides, 20·1), Co =
       10·8, Cu = 11·6, Zn = 10·2, As = 15·4, Bi = 15·3, Ag = 15·7, Cd =
       13·6, I = 24·5, Pt = 26·0, Hg = 20·2, Pb = 24·8, &c. The
       refraction equivalents of many elements could only be calculated
       from the solutions of their compounds. The composition of a
       solution being known it is possible to calculate the refraction
       equivalent of one of its component parts, those for all its other
       components being known. The results are founded on the acceptance
       of a law which cannot be strictly applied. Nevertheless the
       representation of the refraction equivalents gives an easy means
       for directly, although only approximately, obtaining the
       coefficient of refraction from the chemical composition of a
       substance. For instance, the composition of carbon bisulphide is
       CS_{2} = 76, and from its density, 1·27, we find its coefficient
       of refraction to be 1·618 (because the refraction equivalent = 5 +
       2 × 16 = 37), which is very near the actual figure. It is evident
       that in the above representation compounds are looked on as simple
       mixtures of atoms, and the physical properties of a compound as
       the sum of the properties present in the elementary atoms forming
       it. If this representation of the presence of simple atoms in
       compounds had not existed, the idea of combining by a few figures
       a whole mass of data relating to the coefficient of refraction of
       different substances could hardly have arisen. For further details
       on this subject, see works on _Physical Chemistry_.




                              CHAPTER VIII

                       CARBON AND THE HYDROCARBONS


It is necessary to clearly distinguish between the two closely-allied
terms, charcoal and carbon. Charcoal is well known to everybody,
although it is no easy matter to obtain it in a chemically pure state.
Pure charcoal is a simple, insoluble, infusible, combustible substance
produced by heating organic matter, and has the familiar aspect of
a black mass, devoid of any crystalline structure, and completely
insoluble. Charcoal is a substance possessing a peculiar combination
of physical and chemical properties. This substance, whilst in a state
of ignition, combines directly with oxygen; in organic substances it
is found in combination with hydrogen, oxygen, nitrogen, and sulphur.
But in all these combinations there is no real charcoal, as in the same
sense there is no ice in steam. What is found in such combinations
is termed 'carbon'--that is, an element common to charcoal, to those
substances which can be formed from it, and also to those substances
from which it can be obtained. Carbon may take the form of charcoal, but
occurs also as diamond and as graphite. Truly no other element has such
a wide terminology. Oxygen is always called 'oxygen,' whether it is in a
free gaseous state, or in the form of ozone, or oxygen in water, or in
nitric acid or in carbonic anhydride. But here there is some confusion.
In water it is evident that there is no oxygen in a gaseous form, such
as can be obtained in a free state, no oxygen in the form of ozone, but
a substance which is capable of producing both oxygen, ozone, and water.
As an element, oxygen possesses a known chemical individuality, and an
influence on the properties of those combinations into which it enters.
Hydrogen gas is a substance which reacts with difficulty, but the
element hydrogen represents in its combinations an easily displaceable
component part. Carbon may be considered as an atom of carbon matter,
and charcoal as a collection of such atoms forming a whole substance,
or mass of molecules of the substance. The accepted atomic weight
of carbon is 12, because that is the least quantity of carbon which
enters into combination in molecules of its compounds; but the weight
of the molecules of charcoal is probably very much greater. This weight
remains unknown because charcoal is capable of but few direct reactions
and those only at a high temperature (when the weight of its molecules
probably changes, as when ozone changes into oxygen), and it does not
turn into vapour. Carbon exists in nature, both in a free and combined
state, in most varied forms and aspects. Carbon in a free state is
found in at least three different forms, as charcoal, graphite, and the
diamond. In a combined state it enters into the composition of what are
called organic substances--a multitude of substances which are found in
all plants and animals. It exists as carbonic anhydride both in air and
in water, and in the soil and crust of the earth as salts of carbonic
acid and as organic remains.

The variety of the substances of which the structure of plants and
animals is built up is familiar to all. Wax, oil, turpentine, and
tar, cotton and albumin, the tissue of plants and the muscular fibre
of animals, vinegar and starch, are all vegetable and animal matters,
and all carbon compounds.[1] The class of carbon compounds is so vast
that it forms a separate branch of chemistry, known under the name of
organic chemistry--that is, the chemistry of carbon compounds, or, more
strictly, of the hydrocarbons and their derivatives.

  [1] Wood is the non-vital part of ligneous plants: the vital part of
      ordinary trees is situated between the bark and the lignin. Every
      year a layer of lignin is deposited on this part by the juices
      which are absorbed by the roots and drawn up by the leaves; for
      this reason the age of trees may be determined by the number of
      lignin layers deposited. The woody matter consists principally
      of fibrous tissue on to which the lignin or so-called incrusting
      matter has been deposited. The tissue has the composition
      C_{6}H_{10}O_{5}, the substance deposited on it contains more
      carbon and hydrogen and less oxygen. This matter is saturated
      with moisture when the wood is in a fresh state. Fresh birch wood
      contains about 31 p.c. of water, lime wood 47 p.c., oak 35 p.c.,
      pine and fir about 37 p.c. When dried in the air the wood loses a
      considerable quantity of water and not more than 19 p.c. remains.
      By artificial means this loss of water may be increased. If water
      be driven into the pores of wood the latter becomes heavier than
      water, as the lignin of which it is composed has a density of
      about 1·6. One cubic centimetre of birch wood does not weigh more
      than 0·901 gram, fir 0·894, lime tree 0·817, poplar 0·765 when
      in a fresh state; when in a dry state birch weighs 0·622, pine
      0·550, fir 0·355, lime 0·430, guaiacum 1·342, ebony 1·226. On one
      hectare (2·7 acres) of woodland the yearly growth averages the
      amount of 3,000 kilograms (or about 3 tons) of wood, but rarely
      reaches as much as 5,000 kilos. The average chemical composition
      of wood dried in air may be expressed as follows:--Hygroscopic
      water 15 p.c., carbon 42 p.c., hydrogen 5 p.c., oxygen and
      nitrogen 37 p.c., ash 1 p.c. Wood parts with its hygroscopic water
      at 150°, and decomposes at about 300°, giving a brown, brittle,
      so-called red charcoal; above 350° black charcoal is produced.
      As the hydrogen contained in wood requires for its combustion
      about forty parts by weight of oxygen, which is present to the
      amount of about 36 p.c., all that burns of the wood is the carbon
      which it contains, 100 parts of wood only giving out as much heat
      as forty parts of charcoal, and therefore it would be far more
      profitable to use charcoal for heating purposes than wood, if it
      were possible to obtain it in such quantities as correspond with
      its percentage ratio--that is forty parts per 100 parts of wood.
      Generally, however, the quantity produced is far less, not more
      than 30 p.c., because part of the carbon is given off as gas, tar,
      &c. If wood has to be transported great distances, or if it is
      necessary to obtain a very high temperature by burning it, then
      even as little as 25 p.c. of charcoal from 100 parts of wood may be
      advantageous. Charcoal (from wood) develops on burning 8,000 heat
      units, whilst wood dried in air does not develop more than 2,800
      units of heat; therefore seven parts of charcoal give as much heat
      as twenty parts of wood. As regards the temperature of combustion,
      it is far higher with charcoal than with wood, because twenty parts
      of burning wood give, besides the carbonic anhydride which is also
      formed together with charcoal, eleven parts of water, the
      evaporation of which requires a considerable amount of heat.

      [Illustration: FIG. 57.--Apparatus for the dry distillation of
      wood. The retort _a_ containing the wood is heated by the flues _c
      e_. The steam and volatile products of distillation pass along the
      tube g through the condenser _m_, where they are condensed. The
      form, distribution, and dimensions of the apparatus vary.]

      The composition of the growing parts of plants, the leaves, young
      branches, shoots, &c., differs from the composition of the wood
      in that these vital parts contain a considerable quantity of sap
      which contains much nitrogenous matter (in the wood itself there
      is very little), mineral salts, and a large amount of water.
      Taking, for example, the composition of clover and pasture hay in
      the green and dry state; in 100 parts of green clover there is
      about 80 p.c. of water and 20 p.c. of dry matter, in which there
      are about 3·5 parts of nitrogenous albuminous matter, about 9·5
      parts of soluble and about 5 parts of insoluble non-nitrogenous
      matter, and about 2 p.c. of ash. In dry clover or clover-hay there
      is about 15 p.c. of water, 13 p.c. of nitrogenous matter, and 7
      p.c. of ash. This composition of grassy substances shows that they
      are capable of forming the same sort of charcoal as wood itself.
      It also shows the difference of nutritive properties existing
      between wood and the substances mentioned. These latter serve as
      food for animals, because they contain those substances which are
      capable of being dissolved (entering into the blood) and forming
      the body of animals; such substances are proteids, starch, &c. Let
      us remark here that with a good harvest an acre of land gives in
      the form of grass as much organic substance as it yields in the
      form of wood.

      One hundred parts of dry wood are capable of giving, on dry
      distillation, besides 25 p.c. of charcoal and 10 p.c. or more of
      tar, 40 p.c. of watery liquid, containing acetic acid and wood
      spirit, and about 25 p.c. of gases, which may be used for heating
      or lighting purposes, because they do not differ from ordinary
      illuminating gas, which can indeed be obtained from wood. As
      wood-charcoal and tar are valuable products, in some cases the
      dry distillation of wood is carried on principally for producing
      them. For this purpose those kinds of woods are particularly
      advantageous which contain resinous substances, especially
      coniferous trees, such as fir, pine, &c.; birch, oak, and ash
      give much less tar, but on the other hand they yield more aqueous
      liquor. The latter is used for the manufacture of wood spirit,
      CH_{4}O, and acetic acid, C_{2}H_{4}O_{2}. In such cases, the
      dry distillation is carried on in stills. The stills are nothing
      more than horizontal or vertical cylindrical retorts, made of
      boiler plate, heated with fuel and having apertures at the top
      and sometimes also at the bottom for the exit of the light and
      heavy products of distillation. The dry distillation of wood in
      stoves is carried on in two ways, either by burning a portion of
      the wood inside the stove in order to submit the remainder to dry
      distillation by means of the heat obtained in this manner, or by
      placing the wood in a stove the thin sides of which are surrounded
      with a flue leading from the fuel, placed in a space below.

      The first method does not give such a large amount of liquid
      products of the dry distillation as the latter. In the latter
      process there is generally an outlet below for emptying out the
      charcoal at the close of the operation. For the dry distillation
      of 100 parts of wood from forty to twenty parts of fuel are used.

      In the north of Russia wood is so plentiful and cheap that this
      locality is admirably fitted to become the centre of a general
      trade in the products of its dry distillation. Coal (Note 6),
      sea-weed, turf, animal substances (Chapter VI.), &c., are also
      submitted to the process of dry distillation.

If any one of these organic compounds be strongly heated without free
access of air--or, better still, in a vacuum--it decomposes with more or
less facility. If the supply of air be insufficient, or the temperature
be too low for combustion (_see_ Chapter III.), and if the first
volatile products of transformation of the organic matter are subjected
to condensation (for example, if the door of a stove be opened), an
imperfect combustion takes place, and smoke, with charcoal or soot,
is formed.[2] The nature of the phenomenon, and the products arising
from it, are the same as those produced by heating alone, since that
part which is in a state of combustion serves to heat the remainder of
the fuel. The decomposition which takes place on heating a compound
composed of carbon, hydrogen, and oxygen is as follows:--A part of the
hydrogen is separated in a gaseous state, another part in combination
with oxygen, and a third part separates in combination with carbon, and
sometimes in combination with carbon and oxygen in the form of gaseous
or volatile products, or, as they are also called, the products of
dry distillation. If the vapours of these products are passed through
a strongly heated tube, they are changed again in a similar manner
and finally resolve themselves into hydrogen and charcoal. Altogether
these various products of decomposition contain a smaller amount of
carbon than the original organic matter; part of the carbon remains
in a free state, forming charcoal.[3] It remains in that space where
the decomposition took place, in the shape of the black, infusible,
non-volatile charcoal familiar to all. The earthy matter and all
non-volatile substances (ash) forming a part of the organic matter,
remain behind with the charcoal. The tar-like substances, which require
a high temperature in order to decompose them, also remain mixed with
charcoal. If a volatile organic substance, such as a gaseous compound
containing oxygen and hydrogen, be taken, the carbon separates on
passing the vapour through a tube heated to a high temperature. Organic
substances when burning with an insufficient supply of air give off
soot--that is, charcoal--proceeding from carbon compounds in a state of
vapour, the hydrogen of which has, by combustion, been converted into
water; so, for instance, turpentine, naphthalene, and other hydrocarbons
which are with difficulty decomposed by heat, easily yield carbon in the
form of soot during combustion. Chlorine and other substances which,
like oxygen, are capable of taking up hydrogen, and also substances
which are capable of taking up water, can also separate carbon from (or
char) most organic substances.

  [2] The result of imperfect combustion is not only the loss of a part
      of the fuel and the production of smoke, which in some respects
      is inconvenient and injurious to health, but also a low flame
      temperature, which means that a less amount of heat is transmitted
      to the object heated. Imperfect combustion is not only always
      accompanied by the formation of soot or unburnt particles of
      charcoal, but also by that of carbonic oxide, CO, in the smoke
      (Chapter IX.) which burns, emitting much heat. In works and
      factories where large quantities of fuel are consumed, many
      appliances are adopted to ensure perfect combustion, and to combat
      against such a ruinous practice as the imperfect combustion of
      fuel. The most effective and radical means consists in employing
      combustible gases (producer and water gases), because by their
      aid perfect combustion can be easily realised without a loss of
      heat-producing power and the highest temperature can be reached.
      When solid fuel is used (such as coal, wood, and turf), imperfect
      combustion is most liable to occur when the furnace doors are
      opened for the introduction of fresh fuel. The step furnace may
      often prove a remedy for this defect. In the ordinary furnace
      fresh fuel is placed on the burning fuel, and the products of dry
      distillation of the fresh fuel have to burn at the expense of
      the oxygen remaining uncombined with the burnt fuel. Imperfect
      combustion is observed in this case also from the fact that the
      dry distillation and evaporation of the water of the fresh fuel
      lying on the top of that burnt, lowers the temperature of the
      flame, because part of the heat becomes latent. On this account a
      large amount of smoke (imperfect combustion) is observed when a
      fresh quantity of fuel is introduced into the furnace. This may
      be obviated by constructing the furnace (or managing the stoking)
      in such a way that the products of distillation pass through
      the red-hot charcoal remaining from the burnt fuel. It is only
      necessary in order to ensure this to allow a sufficient quantity
      of air for perfect combustion. All this may be easily attained
      by the use of step fire-bars. The fuel is fed into a hopper and
      falls on to the fire-bars, which are arranged in the form of a
      staircase. The burning charcoal is below, and hence the flame
      formed by the fresh fuel is heated by the contact of the red-hot
      burning charcoal. An air supply through the fire grate, an equal
      distribution of the fuel on the fire-bars (otherwise the air will
      blow through empty spaces and lower the temperature), a proper
      proportion between the supply of air and the chimney draught,
      and a perfect admixture of air with the flame (without an undue
      excess of air), are the means by which we can contend against the
      imperfect combustion of such kinds of fuel as wood, peat, and
      ordinary (smoky) coal. Coke, charcoal, anthracite, burn without
      smoke, because they do not contain hydrogenous substances which
      furnish the products of dry distillation, but imperfect combustion
      may occur with them also; in that case the smoke contains carbonic
      oxide.

  [3] Under the action of air, organic substances are capable of
      oxidising to such an extent that all the carbon and all the
      hydrogen they contain will be transformed into carbonic anhydride
      and water. The refuse of plants and that of animals are subjected
      to such a change whether they slowly decompose and putrefy, or
      rapidly burn with direct access of air. But if the supply of air
      be limited, there can be no complete transformation into water and
      carbonic anhydride, there will be other volatile matters (rich in
      hydrogen), while charcoal must remain as a non-volatile substance.
      All organic substances are unstable, they do not resist heat, and
      change even at ordinary temperatures, particularly if water be
      present. It is therefore easy to understand that charcoal may in
      many cases be obtained through the transformation of substances
      entering into the composition of organisms, but that it is never
      found in a pure state.

      However, water and carbonic anhydride are not the only products
      separated from organic substances. Carbon, hydrogen, and oxygen
      are capable of giving a multitude of compounds; some of these
      are volatile compounds, gaseous, soluble in water--they are
      carried off from organic matter, undergoing change without access
      of air. Others, on the contrary, are non-volatile, rich in
      carbon, unaffected by heat and other agents. The latter remain
      in admixture with charcoal in the place where the decomposition
      takes place; such, for example, are tarry substances. The quantity
      of those bodies which are found mixed with the charcoal is very
      varied, and depends on the energy and duration of the decomposing
      agent. The annexed table shows, according to the data of Violette,
      those changes which wood undergoes at various temperatures when
      submitted to dry distillation by means of superheated steam:--

        +----------------------------+------------------------------+
        |                 Residue    |                              |
        |Temperature  from 100 parts | In 100 parts of the residual |
        |             of alder wood  |           charcoal           |
        +----------------------------+---------------------+--------+
        |                            |   C     H   O and N |  Ash   |
        |    150°         100·0      | 47·5   6·1   46·3   |  0·1   |
        |    350°          29·7      | 76·6   4·1   18·4   |  0·6   |
        |   1032°          18·7      | 81·9   2·3   14·1   |  1·6   |
        |   1500°          17·3      | 95·0   0·7    3·8   |  1·7   |
        +----------------------------+---------------------+--------+

Wood charcoal is prepared in large quantities in a similar manner--that
is, by the partial combustion of wood.[4] In nature a similar process
of carbonisation of vegetable refuse takes place in its transformation
under water, as shown by the marshy vegetation which forms peat.[5] In
this manner doubtless the enormous masses of coal were formed[6] which,
following the example set by England, are now utilised everywhere as
the principal material for heating steam boilers, and in general for
all purposes of heating and burning.[7] Russia possesses many very
rich coalfields, amongst which the Donetz district is most worthy of
remark.[8]

  [4] The object of producing charcoal from wood has been explained in
      Note 1. _Wood charcoal_ is obtained in so-called stacks by
      partially burning the wood, or by means of dry distillation (Note
      1) without the access of air. It is principally manufactured for
      metallurgical processes, especially for smelting and forging
      iron. The preparation of charcoal in stacks has one advantage,
      and that is that it may be done on any spot in the forest. But
      in this way all the products of dry distillation are lost.
      For charcoal burning, a pile or stack is generally built, in
      which the logs are placed close together, either horizontally,
      vertically, or inclined, forming a stack of from six to fifty
      feet in diameter and even larger. Under the stack are several
      horizontal air passages, and an opening in the middle to let out
      the smoke. The surface of the stack is covered with earth and
      sods to a considerable thickness, especially the upper part, in
      order to hinder the free passage of air and to concentrate the
      heat inside. When the stack is kindled, the pile begins to settle
      down by degrees, and it is then necessary to look after the turf
      casing and keep it in repair. As the combustion spreads throughout
      the whole pile, the temperature rises and real dry distillation
      commences. It is then necessary to stop the air holes, in order as
      much as possible to prevent unnecessary combustion. The nature of
      the process is, that part of the fuel burns and develops the heat
      required for subjecting the remainder to dry distillation. The
      charring is stopped when the products of dry distillation, which
      are emitted, no longer burn with a brilliant flame, but the pale
      blue flame of carbonic oxide appears. Dry wood in stacks yields
      about one-fourth of its weight of charcoal.

  [5] When dead vegetable matter undergoes transformation in air, in the
      presence of moisture and lower organisms, there remains a
      substance much richer in carbon--namely, humus, black earth or
      mould. 100 parts of humus in a dry state contain about 70 p.c.
      of carbon. The roots, leaves, and stems of plants which wither
      and fall to the ground form a soil rich in humus. The non-vital
      vegetable substances (ligneous tissue) first form brown matter
      (ulmic compounds), and then black matter (humic substances),
      which are both insoluble in water; after this a brown acid is
      produced, which is soluble in water (apocrenic acid), and lastly
      a colourless acid also soluble in water (crenic acid). Alkali
      dissolves a part of the original brown and black substances,
      forming solutions of a brown tint (ulmic and humic acids) which
      sometimes communicate their colour to springs and rivers. The
      proportion of humus in soil generally has a direct influence on
      its fertility; firstly, because putrefying plants develop carbonic
      anhydride and ammonia, and yield the substances forming the ashes
      of plants, which are necessary to vegetation; secondly, because
      humus is capable of attracting the moisture of the air and of
      absorbing water (twice its weight) and in this way keeps the soil
      in a damp condition, which is indispensable for nourishment;
      thirdly, humus renders the soil porous, and, fourthly, it renders
      it more capable of absorbing the heat of the sun's rays. On this
      account black earth is often most remarkable for its fertility.
      One object of manuring is to increase the quantity of humus in the
      soil, and any easily changeable vegetable or any animal matter
      (composts) may be used. The boundless tracts of black earth soil
      in Russia are capable of bestowing countless wealth on the country.

      The origin and extent of black earth soil are treated in detail in
      Professor Dokouchaeff's works.

      If those substances which produce humus undergo decomposition
      under water, less carbonic anhydride is formed, a quantity of
      marsh gas, CH_{4}, is evolved, and the solid residue forms
      an acid humus found in great quantities in marshy places and
      called _peat_. Peat is especially abundant in the lowlands of
      Holland, North Germany, Ireland, and Bavaria. In Russia it is
      likewise found in large quantities, especially in the North-West
      districts. The old hard forms of peat resemble in composition and
      properties brown coal; the newest formations, as yet unhardened
      by pressure, form very porous masses which retain traces of the
      vegetable matter from which they have been formed. Dried (and
      sometimes pressed) peat is used as fuel. The composition of peat
      varies considerably with the locality in which it is found. When
      dried in air it does not contain less than 15 p.c. of water and
      8 p.c. of ash; the remainder consists of 45 p.c. of carbon, 4
      p.c. of hydrogen, 1 p.c. of nitrogen, and 28 p.c. of oxygen. Its
      heating power is about equivalent to that of wood. The brown
      earthy varieties of coal were probably formed from peat. In other
      cases they have a marked woody structure, and are then known as
      lignites. The composition of the brown sorts of coal resembles in
      a marked degree that of peat--namely, in a dried state brown coal
      contains on an average 60 p.c. of carbon, 5 p.c. of hydrogen, 26
      p.c. of oxygen and nitrogen, and 9 p.c. of ash. In Russia brown
      coal is met with in many districts near Moscow, in the Governments
      of Toula and Tver and the neighbourhood; it is very usually used
      as fuel, particularly when found in thick seams. The brown coals
      usually burn with a flame like wood and peat, and are akin to them
      in heating power, which is half or a third that of the best coal.

  [6] Grass and wood, the vegetation of primæval seas and similar refuse
      of all geological periods, must have been in many cases subjected
      to the same changes they now undergo--that is, under water they
      formed peat and lignites. Such substances, preserved or a long
      time underground, subjected to the action of water, compressed by
      the new strata formed above them, transformed by the separation
      of their more volatile component parts (peat and lignites, even
      in their last condition, still continue to evolve nitrogen,
      carbonic anhydride, and marsh gases) form _coal_. Coal is a
      dense homogeneous mass, black, with an oily or glassy lustre, or
      more rarely dull without any evident vegetable structure; this
      distinguishes it in appearance from the majority of lignites.
      The density of coal (not counting the admixture of pyrites,
      &c.) varies from 1·25 (dry bituminous coal) to 1·6 (anthracite,
      flameless), and even reaches 1·9 in the very dense variety of
      coal found in the Olonetzky government (termed shungite), which
      according to the investigations of Professor Inostrantzeff may be
      regarded as the extreme member of the various forms of coal.

      In order to explain the formation of coal from vegetable matter,
      Cagniard de la Tour enclosed pieces of dried wood in a tube and
      heated them to the boiling point of mercury, when the wood was
      changed into a semi-liquid black mass from which a substance
      exceedingly like coal separated. In this manner some kinds of
      wood formed coal which on being heated left caking coke, others
      non-caking; precisely as we find with the natural varieties of
      coal. Violette repeated these experiments with wood dried at
      150°, and showed that when wood is decomposed in this way, a
      gas, an aqueous liquor, and a residue are formed. The latter
      at a temperature of 200° has the properties of wood charcoal
      incompletely burnt; at 300° and higher a homogeneous mass like
      coal is formed which at 340° is dense and without cavities. At
      400° the residue resembles anthracite. In nature probably the
      decomposition was in rare cases effected by heat alone; more
      generally it was effected by means of water and heat, but in
      either case the result ought to be almost the same.

      The average composition of coal compiled from many analyses,
      disregarding the ash, is as follows: 84 parts of carbon, 5 parts
      of hydrogen, 1 part of nitrogen, 8 parts of oxygen, 2 of sulphur.
      The quantity of ash is on an average 5 p.c., but there are coals
      which contain a larger quantity, and naturally they are not so
      advantageous for use as fuel. The amount of water does not usually
      exceed more than 10 p.c. The _anthracites_ form a remarkable
      variety of coals, they do not give any volatile products, or but
      a very small amount, as they contain but little hydrogen compared
      to oxygen. In the average composition of coal we saw that for 5
      parts of hydrogen there were 8 parts of oxygen; therefore 4 parts
      by weight of hydrogen are capable of forming hydrocarbons, because
      1 part of hydrogen is necessary in order to form water with the 8
      parts of oxygen. These 4 parts by weight of hydrogen can convert
      48 parts of carbon into volatile products, because 1 part of
      hydrogen by weight in these substances combines with 12 parts of
      carbon. The anthracites differ essentially from this: neglecting
      the ash, their average composition is as follows: 94 parts of
      carbon, 3 of hydrogen, and 3 of oxygen and nitrogen. According to
      the analyses of A. A. Voskresensky, the Grousheffsky anthracite
      (Don district) contains: C = 93·8, H = 1·7, ash = 1·5. Therefore
      the anthracites contain but little hydrogen capable of combining
      with the carbon to form hydrocarbons which burn with a flame.
      Anthracites are the oldest forms of coal. The newest and least
      transformed coals, which resemble some of the brown varieties, are
      the _dry_ coals. They burn with a flame like wood, and leave a
      coke having the appearance of lumps of coal, half their component
      parts being absorbed by the flame (they contain much hydrogen and
      oxygen). The remaining varieties of coal (gas coal, smithy coal,
      coking, and anthracite) according to Grüner in all respects form
      connecting links between the _dry_ coals and the anthracites.
      These coals burn with a very smoky flame, and on being heated
      leave _coke_, which bears the same relation to coal that charcoal
      does to wood. The quantity and quality of coke vary considerably
      with the different sorts of coal from which it is formed. In
      practice coals are most often distinguished by the properties
      and quantity of the coke which they give. In this particular the
      so-called bituminous coals are especially valuable, as even the
      slack of this kind gives on dry distillation large spongy masses
      of coke. If large pieces of these kinds of coal are subjected
      to dry distillation, they, as it were, melt, flow together, and
      form caking masses of coke. The best coking coals give 65 p.c. of
      dense caking coke. Such coal is very valuable for metallurgical
      purposes (_see_ Note 8). Besides coke, the dry distillation of
      coal produces gas (_see_ further, illuminating gas, p. 361),
      coal-tar (which gives benzene, carbolic acid, naphthalene, tar
      for artificial asphalt, &c.) and also an aqueous alkaline liquor
      (with wood and lignites the liquid is acid from acetic acid) which
      contains ammonium carbonate (_see_ Note 6).

  [7] In England in 1850 the output of coal was as much as 48 million
      tons, and in latter years it has risen to about 190 millions.
      Besides this other countries contribute 300 millions--Russia about
      6 millions. The United States of America come next to England with
      an output of 160 million tons, then Germany 90 millions; France
      produces but little (25 millions), and takes about 5 million tons
      from England. Thus the world consumes about 500 million tons of
      coal yearly. Besides household purposes, coal is chiefly used as
      fuel for steam-engines. As every horse-power (= 75 kilogrammetres
      per second) of a steam-engine expends on the average more than
      25 kilograms in 24 hours, or in a year (counting stoppages) not
      less than 5 tons per horse-power, and there are not less than 40
      million horse-power at work in the world, the consumption of coal
      for motive-power is at least equal to half the whole production.
      For this reason coal serves as a criterion of the industrial
      development of a country. About 15 p.c. of coal is used for the
      manufacture of cast iron, wrought iron, steel, and articles made
      of them.

  [8] The principal coal beds of Russia under exploitation are: The Don
      basin (150 million poods per annum, 62 poods = 1 ton), the Polish
      basin (Dombrovo and others 120 million poods per annum), the Toula
      and Riazan beds of the Moscow basin (up to 25 million poods),
      the Ural basin (10 million poods), the Caucasian (Kviboul, near
      Kutais), the Khirjhis steppes, the smithy coal basin (Gov. of
      Tomsk), the Sahaline, &c. The Polish and Moscow basins do not give
      any coking coals. The presence of every variety of coal (from
      the dry coal near Lisichansk on the Donetz to the anthracites of
      the entire south-east basin), the great abundance of excellent
      metallurgical coal (coking, _see_ Note 6) in the western part
      of the basin, its vast extent (as much as 25,000 sq. versts),
      the proximity of the seams to the surface (the shafts are now
      from 20 to 100 fathoms deep, and in England and Belgium as deep
      as 500 fathoms), the fertility of the soil (black earth), the
      proximity of the sea (about 100 versts from the Sea of Azoff) and
      of the rivers Donetz, Don, and Dneiper, the most abundant seams
      of excellent iron ore (Korsan Mogila, Krivoy Rog, Soulin, &c.,
      &c.), copper ore, mercury ore (near Nikitovka, in the Bakhmouth
      district of the Ekaterinoslav Gov.), and other ores, the richest
      probably in the whole world, the beds of rock-salt (near the
      stations of the Stoupka and Brianzovka) the excellent clay of all
      kinds (china, fire-clay), gypsum, slate, sandstone, and other
      _wealth of the Don coal basin_, give complete assurance of the
      fact that with the growth of industrial activity in Russia this
      bountiful land of the Cossacks and New Russia will become the
      centre of the most extensive productive enterprise, not for the
      requirements of Russia alone, but of the whole world, because in
      no other place can be found such a concentration of favourable
      conditions. The growth of enterprise and knowledge, together with
      the extinction of the forests which compels Russia to foster the
      production of coal, will help to bring about this desired result.
      England with a whole fleet of merchant vessels exports annually
      about 25 million tons of coal, the price of which is higher than
      on the Donetz (where a pood of worked coal costs less than 5
      copecks on the average), where anthracites and semi-anthracites
      (like Cardiff or steam coal, which burns without smoke) and coking
      and metallurgical coals are able both in quantity and quality to
      satisfy the most fastidious requirements of the industry already
      existing and rapidly increasing everywhere. The coal mines of
      England and Belgium are approaching a state of exhaustion, whilst
      in those of the Don basin, only at a depth of 100 fathoms,
      1,200,000 million poods of coal lie waiting to be worked.

During the imperfect combustion of volatile substances containing
carbon and hydrogen, the hydrogen and part of the carbon first burn,
and the remainder of the carbon forms soot. Tar, pitch, and similar
substances for this reason burn with a smoky flame. Thus soot is
finely-divided charcoal separated during the imperfect combustion
of the vapours and gases of carbonaceous substances rich in carbon.
Specially-prepared soot (lampblack) is very largely used as a black
paint and a large quantity goes for the manufacture of printers' ink. It
is prepared by burning tar, oil, natural gas, naphtha, &c. The quantity
of organic matter remaining undecomposed in the charcoal depends on the
temperature to which it has been submitted. Charcoal prepared at the
lowest temperature still contains a considerable quantity of hydrogen
and oxygen--even as much as 4 p.c. of hydrogen and 20 p.c. of oxygen.
Such charcoal still preserves the structure of the substance from
which it was obtained. Ordinary charcoal, for instance, in which the
structure of the tree is still visible, is of this kind. On submitting
it to further heating, a fresh quantity of hydrogen with carbon and
oxygen (in the form of gases or volatile matter) may be separated, and
the purest charcoal will be obtained on submitting it to the greatest
heat.[9] If it be required to prepare pure charcoal from soot it is
necessary first to wash it with alcohol and ether in order to remove the
soluble tarry products, and then submit it to a powerful heat to drive
off the impurities containing hydrogen and oxygen. Charcoal however when
completely purified does not change in appearance. Its porosity,[10]
bad conducting power for heat, capability of absorbing the luminous
rays (hence its blackness and opacity), and many other qualities, are
familiar from everyday experience.[11] The specific gravity of charcoal
varies from 1·4 to 1·9, and that it floats on water is due to the air
contained in its pores. If charcoal is reduced to a powder and moistened
with spirit, it immediately sinks in water. It is _infusible_ in the
furnace and even at the temperature of the oxyhydrogen flame. In the
heat generated by means of a powerful galvanic current charcoal only
softens but does not completely melt, and on cooling it is found to
have undergone a complete change both in properties and appearance,
and is more or less transformed into graphite. The physical stability
of charcoal is without doubt allied to its chemical stability. It
is evidently a substance devoid of energy, for it is insoluble in
all known liquids, and _at an ordinary temperature does not combine
with anything_; it is an inactive substance, like nitrogen.[12] But
these properties of charcoal change with a rise of temperature; thus,
unlike nitrogen, charcoal, at a high temperature, combines directly
with oxygen. This is well known, as charcoal burns in air. Indeed, not
only does oxygen _combine with charcoal at a red heat_, but sulphur,
hydrogen, silicon, and also iron and some other metals[12 bis] do
so at a very high temperature--that is, when the molecules of the
charcoal have reached a state of great instability--whilst at ordinary
temperatures neither oxygen, sulphur, nor metals act on charcoal in
any way. When burning in oxygen, charcoal forms carbonic anhydride,
CO_{2}, whilst in the vapours of sulphur, carbon bisulphide, CS_{2}, is
formed, and wrought iron, when acted on by carbon, becomes cast iron.
At the great heat obtained by passing the galvanic current through
carbon electrodes, charcoal combines with hydrogen, forming acetylene,
C_{2}H_{2}. Charcoal does not combine directly with nitrogen, but in the
presence of metals and alkaline oxides, nitrogen is absorbed, forming
a metallic cyanide, as, for instance, potassium cyanide, KCN. From
these few direct combinations which charcoal is capable of entering
into, may be derived those numerous carbonaceous compounds which enter
into the composition of plants and animals, and can be thus obtained
artificially. Certain substances containing oxygen give up a part of
it to charcoal at a relatively low temperature. For instance, nitric
acid when boiled with charcoal gives carbonic anhydride and nitric
peroxide. Sulphuric acid is reduced to sulphurous anhydride when heated
with carbon. When heated to redness charcoal absorbs oxygen from a large
number of the oxides. Even such oxides as those of sodium and potassium,
when heated to redness, yield their oxygen to charcoal although they
do not part with it to hydrogen. Only a few of the oxides, like silica
(oxide of silicon) and lime (calcium oxide) resist the reducing
action of charcoal. Charcoal is capable of changing its physical
condition without undergoing any alteration in its essential chemical
properties--that is, it passes into _isomeric_ or _allotropic forms_.
The two other particular forms in which carbon appears are the _diamond_
and _graphite_. The identity of composition of these with charcoal is
proved by burning an equal quantity of all three separately in oxygen
(at a very high temperature), when each gives the same quantity of
carbonic anhydride--namely, 12 parts of charcoal, diamond, or graphite
in a pure state, yield on burning 44 parts by weight of carbonic
anhydride. The physical properties present a marked contrast; the
densest sorts of charcoal have a density of only 1·9, whilst the density
of graphite is about 2·3, and that of the diamond 3·5. A great many
other properties depend on the density, for instance combustibility.
The lighter charcoal is, the more easily it burns; graphite burns with
considerable difficulty even in oxygen, and the diamond burns only
in oxygen and at a very high temperature. On burning, charcoal, the
diamond, and graphite develop different quantities of heat. One part by
weight of wood charcoal converted by burning into carbonic anhydride
develops 8,080 heat units; dense charcoal separated in gas retorts
develops 8,050 heat units; natural graphite, 7,800 heat units; and the
diamond 7,770. The greater the density the less the heat evolved by the
combustion of the carbon.[13]

  [9] As it is difficult to separate from the charcoal the admixture of
      ash--that is, the earthy matter contained in the vegetable
      substance used for producing charcoal--in order to obtain it
      in its purest condition it is necessary to use such organic
      substances as do not contain any ash, for example completely
      refined or purified crystallised sugar, crystallised tartaric
      acid, &c.

  [10] The cavities in charcoal are the passages through which those
       volatile products formed at the same time as the charcoal have
       passed. The degree of porosity of charcoal varies considerably,
       and has a technical significance, in different kinds of charcoal.
       The most porous charcoal is very light; a cubic metre of wood
       charcoal weighs about 200 kilograms. Many of the properties of
       charcoal which depend exclusively on its porosity are shared by
       many other porous substances, and vary with the density of the
       charcoal and depend on the way it was prepared. The property
       which charcoal has of absorbing gases, liquids, and many
       substances in solution, is a case in point. The densest kind
       of charcoal is formed by the action of great heat on sugar and
       other fusible substances. The lustrous grey dense coke formed in
       gas retorts is also of this character. This dense coke collects
       on the internal walls of the retorts subjected to great heat,
       and is produced by the vapours and gases separated from the
       heated coal in the retorts. In virtue of its density such coke
       becomes a good conductor of the galvanic current and approaches
       graphite. It is principally used in galvanic batteries. Coke, or
       the charcoal remaining from the imperfect combustion of coal and
       tarry substances, is also but slightly porous, brilliant, does
       not soil or mark paper, is dense, almost devoid of the faculty
       of retaining liquids and solids, and does not absorb gases. The
       light sorts of charcoal produced from charred wood, on the other
       hand, show this absorptive power in a most marked degree. This
       property is particularly developed in that very fine and friable
       charcoal prepared by heating animal substances such as hides and
       bones. _The absorptive power of charcoal_ for gases is similar
       to the condensation of gases in spongy platinum. Here evidently
       there is a case of the adherence of gases to a solid, precisely
       as liquids have the property of adhering to various solids. One
       volume of charcoal will absorb the following volumes of gases
       (charcoal is capable of absorbing an immense amount of chlorine,
       almost equal to its own weight):--

         --------------------------------------------------------
              Saussure.          Favre.           Heat emitted
          Boxwood Charcoal  Cocoanut Charcoal   per gram of gas
         -------------------------------------------------------
             NH_{3}  90          172  vols.        494  units
             CO_{2}  35           97    "          158    "
             N_{2}O  40           99    "          169    "
             HCl     85          165    "          274    "
         -------------------------------------------------------

       The quantity of gas absorbed by the charcoal increases with the
       pressure, and is approximately proportional to it. The quantity
       of heat given out by the absorption nearly approaches that set
       free on dissolving, or passing into a liquid condition.

       Charcoal absorbs not only gases, but a number of other
       substances. For instance, alcohol which contains disagreeably
       smelling fusel oil, on being mixed with charcoal or filtered
       through it, loses most of the fusel oil. The practice of
       filtering substances through charcoal in order to get rid of
       foreign matters is often applied in chemical and manufacturing
       processes. Oils, spirits, various extracts, and vegetable and
       other solutions are filtered through charcoal in order to purify
       them. The bleaching power of charcoal may be tested by using
       various  solutions--such as aniline dyes, litmus, &c.
       Charcoal, which has absorbed one substance to saturation is still
       capable of absorbing certain other substances. Animal charcoal,
       produced in a very finely-divided state, especially by heating
       bones, makes the best sort for the purposes of absorption. Bone
       charcoal is used in large quantities in sugar works for filtering
       syrups and all saccharine solutions, in order to purify them,
       not only from colouring and odorous matter, but also from the
       lime which is mixed with the syrups in order to render them
       less unstable during boiling. The absorption of lime by animal
       charcoal depends, in all probability, in a great degree on the
       mineral component parts of bone charcoal.

  [11] Charcoal is a very bad conductor of heat, and therefore forms an
       excellent insulator or packing to prevent the transmission of
       heat. A charcoal lining is often used in crucibles for heating
       many substances, as it does not melt and resists a far greater
       heat than many other substances.

  [12] The unalterability of charcoal under the action of atmospheric
       agencies, which produce changes in the majority of stony and
       metallic substances, is often made use of in practice. For
       example, charcoal is frequently strewn in boundary ditches. The
       surface of wood is often charred to render it durable in those
       places where the soil is damp and wood itself would soon rot. The
       chambers (or in some works towers) through which acids pass (for
       example, sulphuric and hydrochloric) in order to bring them into
       contact with gases or liquids, are filled with charcoal or coke,
       because at ordinary temperatures it resists the action of even
       the strongest acids.

  [12 bis] Maquenne (1892) discovered that carbon is capable of combining
       with the alkali metals. A 20 p.c. amalgam of the metals was heated
       to a red heat with charcoal powder in a stream of hydrogen. The
       compounds so obtained possessed, after the mercury had been driven
       off, the compositions BaC_{2}, SrC_{2}, CaC_{2}. All these
       compounds react with water forming acetylene, for example:

             BaC_{2} + 2H_{2}O = C_{2}H_{2} + Ba(OH)_{2}

       Maquenne proposes the barium carbide as a source of acetylene. He
       obtained this compound by heating carbonate of barium, magnesium
       powder, and retort carbon in a Perreau furnace (BaCO_{3} + 3Mg + C
       = 3MgO + BaC_{2}). One hundred grams of BaC_{2} evolve 5,200 to
       5,400 c.c. of acetylene, mixed with about 2-3 p.c. of hydrogen.

       The relation of acetylene, C_{2}H_{2}, to these metallic carbides
       is evident from the fact that these metals (Ca, Sr, Ba) replace 2
       atoms of hydrogen, and therefore C_{2}Ba corresponds to
       C_{2}H_{2}, so that they may be regarded as metallic derivatives
       of acetylene. Moissan (1893) obtained similar carbides directly
       from the oxides by subjecting them to the action of the voltaic
       arc, in the presence of carbon, for instance, BaO + 3C = CO +
       C_{2}Ba, although at a furnace heat carbon has no action on the
       oxides CaO, BaO, SrO. Concerning Al_{4}C_{5}, _see_ Chapter XVII.
       Note 38.

  [13] When subjected to pressure, charcoal loses heat, hence the densest
       form stands to the less dense as a solid to a liquid, or as a
       compound to an element. From this the conclusion may be drawn
       that the molecules of graphite are more complex than those of
       charcoal, and those of the diamond still more so. The specific
       heat shows the same variation, and as we shall see further on,
       the increased complexity of a molecule leads to a diminution of
       the specific heat. At ordinary temperatures the specific heat of
       charcoal is 0·24, graphite 0·20, the diamond 0·147. For retort
       carbon Le Chatelier (1893) found that the product of the sp. heat
       and atomic weight varies, between 0° and 250°, according to the
       formula: = 1·92 + 0·0077_t_, and between 250° and 1000°, = 3·54 +
       0·00246_t_ (_see_ Chapter XIV. Note 4).

By means of intense heat charcoal may be transformed into graphite.
If a charcoal rod 4 mm. in diameter and 5 mm. long be enclosed in an
exhausted receiver and the current from 600 Bunsen's elements, placed
in parallel series of 100, be passed through it, the charcoal becomes
strongly incandescent, partially volatilises, and is deposited in the
form of graphite. If sugar be placed in a charcoal crucible and a
powerful galvanic current passed through it, it is baked into a mass
similar to graphite. If charcoal be mixed with wrought iron and heated,
cast iron is formed, which contains as much as five per cent. of
charcoal. If molten cast iron be suddenly chilled, the carbon remains in
combination with the iron, forming so called white cast iron; but if the
cooling proceeds slowly, the greater part of the carbon separates in the
form of graphite, and if such cast iron (so called grey cast iron) be
dissolved in acid, the carbon remains in the form of graphite. Graphite
is met with in nature, sometimes in the form of large compact masses,
sometimes permeating rocky formations like the schists or slates, and
in fact is met with in those places which, in all probability, have
been subjected to the action of subterranean heat.[14] The graphite in
cast iron, and sometimes also natural graphite, occasionally appears in
a crystalline form in the shape of six-sided plates, but more often it
occurs as a compact amorphous mass having the characteristic properties
of the familiar black-lead pencil.[15]

  [14] There are places where anthracite gradually changes into graphite
       as the strata sink. I myself had the opportunity of observing
       this gradual transformation in the valley of Aosta.

  [15] Pencils are made of graphite worked up into a homogeneous mass by
       disintegrating, powdering, and cleansing it from earthy
       impurities; the best kinds are made of completely homogeneous
       graphite sawn up into the requisite sticks. Graphite is found in
       many places. In Russia the so-called Aliberoffsky graphite is
       particularly renowned; it is found in the Altai mountains near
       the Chinese frontier; in many places in Finland and likewise
       on the banks of the Little Tungouska, Sidoroff also found a
       considerable quantity of graphite. When mixed with clay, graphite
       is used for making crucibles and pots for melting metals.

       Graphite, like most forms of charcoal, still contains a certain
       quantity of hydrogen, oxygen, and ash, so that in its natural
       state it does not contain more than 98 _p.c._ of carbon.

       In practice, graphite is purified simply by washing it when in
       a finely-ground state, by which means the bulk of the earthy
       matter may be separated. The following process, proposed by
       Brodie, consists in mixing the powdered graphite with 1/14 part
       of its weight of potassium chlorate. The mixture is then heated
       with twice its weight of strong sulphuric acid until no more
       odoriferous gases are emitted; on cooling, the mixture is thrown
       into water and washed; the graphite is then dried and heated
       to a red heat; after this it shrinks considerably in volume
       and forms a very fine powder, which is then washed. By acting
       on graphite several times with a mixture of potassium chlorate
       and nitric acid heated up to 60°, Brodie transformed it into a
       yellow insoluble acid substance which he called graphitic acid,
       C_{11}H_{4}O_{5}. The diamond remains unchanged when subjected to
       this treatment, whilst amorphous charcoal is completely oxidised.
       Availing himself of this possibility of distinguishing graphite
       from the diamond or amorphous charcoal, Berthelot showed that
       when compounds of carbon and hydrogen are decomposed by heat,
       amorphous charcoal is mainly formed, whilst when compounds of
       carbon with chlorine, sulphur, and boron are decomposed, graphite
       is principally deposited.

The diamond is a crystalline and transparent form of carbon. It is of
rare occurrence in nature, and is found in the alluvial deposits of
the diamond mines of Brazil, India, South Africa, &c. It has also been
found in meteorites.[15 bis] It crystallises in octahedra, dodecahedra,
cubes, and other forms of the regular system.[16] The efforts which
have been made to produce diamonds artificially, although they have not
been fruitless, have not as yet led to the production of large-sized
crystals, because those means by which crystals are generally formed
are inapplicable to carbon. Indeed, carbon in all its forms being
insoluble and infusible does not pass into a liquid condition by means
of which crystallisation could take place. Diamonds have several
times been successfully produced in the shape of minute crystals
having the appearance of a black powder, but when viewed under the
microscope appearing transparent, and possessing that hardness which
is the peculiar characteristic of the diamond. This diamond powder is
deposited on the negative electrode, when a weak galvanic current is
passed through liquid chloride of carbon.[16 bis]

  [15 bis] Diamonds are found in a particular dense rock, known by the
       name of itacolumite, and are dug out of the _débris_ produced by
       the destruction of the itacolumite by water. When the _débris_ is
       washed the diamonds remain behind; they are principally found in
       Brazil, in the provinces of Rio and Bahia, and at the Cape of Good
       Hope. The _débris_ gives the black or amorphous diamond,
       carbonado, and the ordinary colourless or yellow translucent
       diamond. As the diamond possesses a very marked cleavage, the
       first operation consists in splitting it, and then roughly and
       finely polishing it with diamond powder. It is very remarkable
       that Professors P. A. Latchinoff and Eroféeff found (1887) diamond
       powder in a meteoric stone which fell in the Government of Penza,
       in the district of Krasnoslobodsk, near the settlement of Novo
       Urei (Sept. 10, 1886). Up to that time charcoal and graphite (a
       special variety, cliftonite) had been found in meteorites and the
       diamond only conjectured to occur therein. The Novo Urei meteorite
       was composed of siliceous matter and metallic iron (with nickel)
       like many other meteorites.

  [16] Diamonds are sometimes found in the shape of small balls, and in
       that case it is impossible to cut them because directly the
       surface is ground or broken they fall into minute pieces.
       Sometimes minute diamond crystals form a dense mass like sugar,
       and this is generally reduced to diamond powder and used for
       grinding. Some known varieties of the diamond are almost opaque
       and of a black colour. Such diamonds are as hard as the ordinary
       ones, and are used for polishing diamonds and other precious
       stones, and also for rock boring and tunnelling.

  [16 bis] Hannay, in 1880, obtained diamonds by heating a mixture of
       heavy liquid hydrocarbons (paraffin oils) with magnesium in a
       thick iron tube. This investigation, however, was not repeated.

Moissan (Paris, 1893) produced diamonds artificially by means of the
high temperature attained in the electrical furnace[17] by dissolving
carbon in molten cast iron, and allowing the solution with an excess of
carbon, to cool under the powerful pressure exerted by rapidly cooling
the metal.[17 bis] K. Chroustchoff attained the same end by means of
silver, which dissolves carbon to the extent of 6 p.c. at a high
temperature. Rousseau, for the same purpose, heated carbide of calcium
in the electric furnace. There is no doubt that all these investigators
obtained the diamond as a transparent body, which burnt into CO_{2}, and
possessed an exceptional hardness, but only in the form of a fine powder.

  [17] The _electrical furnace_ is an invention of recent times, and
       gives the possibility of obtaining a temperature of 3,500°, which
       is not only not obtainable in ordinary furnaces, but even in the
       oxyhydrogen flame, whose temperature does not exceed 2,000°. The
       electrical furnace consists of two pieces of lime, laid one on
       the other. A cavity is made in the lower piece for the reception
       of the substance to be melted between two thick electrodes of
       dense carbon. On passing a current of 70 volts and 450 ampères
       a temperature of 3,000° is easily obtained. At a temperature of
       2,500° (100 ampères and 40 volts) not only do all metals melt,
       but even lime and magnesia (when placed in the space between
       the carbon electrodes, _i.e._ in the voltaic arc) become soft
       and crystallise on cooling. At 3,000° lime becomes very fluid,
       metallic calcium partially separates out and a carbon compound,
       which remains liquid for a long time. At this temperature oxide
       of uranium is reduced to the suboxide and metal, zirconia and
       rock crystal fuse and partially volatilise, as also does alumina;
       platinum, gold, and even carbon distinctly volatilise; the
       majority of the metals form carbides. At such a temperature also
       cast iron and carbon give graphite, while according to Rousseau,
       between 2,000° and 3,000° the diamond passes into graphite and
       conversely graphite into the diamond, so that this is a kind of
       reversible reaction.

  [17 bis] Moissan first investigated the solution of carbon in molten
       metals (and the formation of the carbides) such as magnesium,
       aluminium, iron, manganese, chromium, uranium, silver, platinum,
       and silicon. At the same time Friedel, owing to the discovery of
       the diamond in meteoric iron, admitted that the formation of the
       diamond is dependent upon the influence of iron and sulphur. With
       this object, that is to obtain the diamond, Friedel caused sulphur
       to react upon samples of cast iron rich in carbon, in a closed
       vessel at a maximum temperature of 500°, and after dissolving the
       sulphide of iron formed, he obtained a small quantity of a black
       powder which scratched corundum, i.e. diamond. Moissan's
       experiments (1893) were more successful, probably owing to his
       having employed the electrical furnace. If iron be saturated with
       carbon at a temperature between 1,100° and 3,000°, then at
       1,100°-1,200° a mixture of amorphous carbon and graphite is
       formed, while at 3,000° graphite alone is obtained in very
       beautiful crystals. Thus under these conditions the diamond is not
       formed, and it can only be obtained if the high temperature be
       aided by powerful pressures. For this purpose Moissan took
       advantage of the pressure produced in the passage of a mass of
       molten cast iron from a liquid into a solid state. He first melted
       150-200 grams of iron in the electrical furnace, and quickly
       introduced a cylinder of carbon into the molten iron. He then
       removed the crucible with the molten iron from the furnace and
       plunged it into a reservoir containing water. After treating with
       boiling hydrochloric acid, three varieties of carbon were
       obtained: (1) a small amount of graphite (if the cooling be
       rapid); (2) carbon of a chestnut colour in very fine twisted
       threads, showing that it had been subjected to a very high
       pressure (a similar variety was met with in various samples of the
       Canon Diabolo), and lastly (3) an inconsiderable quantity of an
       exceeding dense mass which was freed from the admixture of the
       lighter modifications by treatment with _aqua regia_, sulphuric
       and hydrofluoric acids, and from which Moissan, by means of liquid
       bromoform (sp. gr. 2·900), succeeded in separating some small
       pieces, having a greater density than bromoform, which scratched
       the ruby and had the properties of the diamond. Some of these
       pieces were black, others were transparent and refracted light
       strongly. The dark grey tint of the former resembled that of the
       black diamonds (carbonado). Their density was between 3 and 3·5.
       The transparent specimens had a greasy appearance and seemed to
       be, as it were, surrounded by an envelope of carbon. At 1,050°
       they did not burn entirely in a current of air, so that the
       imperfectly burnt particles, and a peculiar form of grains of a
       light ochre colour, which retained their crystalline form, could
       be examined under the microscope. Similar grains also remain after
       the imperfect combustion of the ordinary diamond. Moissan obtained
       the same results by rapidly cooling in a stream of coal gas a
       piece of cast iron, saturated with carbon obtained from sugar and
       first heated to 2,000°. In this instance he obtained small
       crystals of diamonds. K. Chroustchoff showed that at its boiling
       point silver dissolves 6 p.c. of carbon. This silver was rapidly
       cooled, so that a crust formed on the surface and prevented the
       metal expanding, and so produced a powerful pressure. A portion of
       the carbon which separates out under these conditions exhibits the
       properties of the diamond.

Judging from the fact that carbon forms a number of gaseous bodies
(carbonic oxide, carbonic anhydride, methane, ethylene, acetylene, &c.)
and volatile substances (for example, many hydrocarbons and their most
simple derivatives), and considering that the atomic weight of carbon, C
= 12, approaches that of nitrogen, N = 14, and that of oxygen, O = 16,
and that the compounds CO (carbonic oxide) and N_{2}C_{2} (cyanogen)
are gases, it may be argued that if carbon formed the molecule C_{2},
like N_{2} and O_{2}, it would be a gas. And as through polymerism or
the combination of like molecules (as O_{2} passes into O_{3} or NO_{2}
into N_{2}O_{4}) the temperatures of ebullition and fusion rise (which
is particularly clearly proved with the hydrocarbons of the C_{n}H_{2n}
series), it ought to be considered that _the molecules of charcoal,
graphite, and the diamond are very complex_, seeing that they are
insoluble, non-volatile, and infusible. The aptitude which the atoms of
carbon show for combining together and forming complex molecules appears
in all carbon compounds. Among the volatile compounds of carbon many are
well known the molecules of which contain C_{5} ... C_{10} ... C_{20}
... C_{30}, &c., in general C_{n} where n may be very large, and in none
of the other elements is this faculty of complexity so developed as in
carbon.[18] Up to the present time there are no grounds for determining
the degree of polymerism of the charcoal, graphite, or diamond
molecules, and it can only be supposed that they contain C_{n} where n
is a large quantity. Charcoal and those complex non-volatile organic
substances which represent the gradual transitions to charcoal[19] and
form the principal solid substances of organisms, contain a store or
accumulation of internal power in the form of the energy binding the
atoms into complex molecules. When charcoal or complex compounds of
carbon burn, the energy of the carbon and oxygen is turned into heat,
and this fact is taken advantage of at every turn for the generation of
heat from fuel.[20]

  [18] The existence of a molecule S_{6} is known (up to 600°), and it
       must he beld that this accounts for the formation of hydrogen
       persulphide, H_{2}S_{5}. Phosphorus appears in the molecule
       P_{4} and gives P_{4}H_{2}. When expounding the data on specific
       heat we shall have occasion to return to the question of the
       complexity of the carbon molecule.

  [19] The hydrocarbons poor in hydrogen and containing many atoms of
       carbon, like chrysene and carbopetrocene, &c.,
       C_{_n_}H_{2(_n_-_m_)}, are solids, and less fusible as _n_ and
       _m_ increase. They present a marked approach to the properties
       of the diamond. And in proportion to the diminution of the water
       in the carbohydrates C_{_n_}H_{2_m_}O_{_m_}--for example in the
       humic compounds (Note 5)--the transition of complex organic
       substances to charcoal is very evident. That residue resembling
       charcoal and graphite which is obtained by the separation (by
       means of copper sulphate and sodium chloride) of iron from white
       cast-iron containing carbon chemically combined with the iron,
       also seems, especially after the researches of G. A. Zaboudsky,
       to be a complex substance containing C_{12}H_{6}O_{3}. The
       endeavours which have been directed towards determining the
       measure of complexity of the molecules of charcoal, graphite, and
       the diamond will probably at some period lead to the solution
       of this problem and will most likely prove that the various
       forms of charcoal, graphite, and the diamond contain molecules
       of different and very considerable complexity. The constancy of
       the grouping of benzene, C_{6}H_{6}, and the wide diffusion and
       facility of formation of the carbohydrates containing C_{6} (for
       example, cellulose, C_{6}H_{10}O_{5}, glucose, C_{6}H_{12}O_{6})
       give reason for thinking that the group C_{6} is the first and
       simplest of those possible to free carbon, and it may be hoped
       that some time or other it may be possible to get carbon in this
       form. Perhaps in the diamond there may be found such a relation
       between the atoms as in the benzene group, and in charcoal such
       as in carbohydrates.

  [20] When charcoal burns, the complex molecule C_{_n_} is resolved into
       the simple molecules _n_CO_{2}, and therefore part of the
       heat--probably no small amount--is expended in the destruction of
       the complex molecule C_{_n_}. Perhaps by burning the most complex
       substances, which are the poorest as regards hydrogen, it may be
       possible to form an idea of the work required to split up C_{_n_}
       into separate atoms.

No other two elements are capable of combining together in such variety
as carbon and hydrogen. The hydrocarbons of the C_{_n_}H_{2_m_} series
in many cases differ widely from each other, although they have some
properties in common. All hydrocarbons, whether gaseous, liquid or
solid, are combustible substances sparingly soluble or insoluble in
water. The liquefied gaseous hydrocarbons, as well as those which are
liquid at ordinary temperatures, and those solid hydrocarbons which
have been liquefied by fusion, have the appearance and property of oily
liquors, more or less viscid, or fluid.[21] The solid hydrocarbons
more or less resemble wax in their properties, although ordinary oils
and wax generally contain oxygen in addition to carbon and hydrogen,
but in relatively small proportion. There are also many hydrocarbons
which have the appearance of tar--as, for instance, metacinnamene and
gutta-percha. Those liquid hydrocarbons which boil at a high temperature
are like oils, and those which have a low boiling point resemble ether,
whilst the gaseous hydrocarbons in many of their properties are akin
to hydrogen. All this tends to show that in hydrocarbons physically
considered the properties of solid non-volatile charcoal are strongly
modified and hidden, whilst those of the hydrogen predominate. All
hydrocarbons are neutral substances (neither basic nor acid), but under
certain conditions they enter into peculiar reactions. It has been seen
in those hydrogen compounds which have been already considered (water,
nitric acid, ammonia) that the hydrogen in almost all cases enters into
reaction, being displaced by metals. The hydrogen of the hydrocarbons,
it may be said, has no metallic character that is to say, it is not
directly[22] displaced by metals, even by such as sodium and potassium.
On the application of more or less heat all hydrocarbons decompose[23]
forming charcoal and hydrogen. The majority of hydrocarbons do not
combine with the oxygen of the air or oxidise at ordinary temperatures,
but under the action of nitric acid and many other oxidising substances
most of them undergo oxidation, in which either a portion of the
hydrogen and carbon is separated, or the oxygen enters into combination,
or else the elements of hydrogen peroxide enter into combination with
the hydrocarbon.[24] When heated in air, hydrocarbons burn, and,
according to the amount of carbon they contain, their combustion is
attended more or less with a separation of soot--that is, finely divided
charcoal--which imparts great brilliancy to the flame, and on this
account many of them are used for the purposes of illumination--as,
for instance, kerosene, coal gas, oil of turpentine. As hydrocarbons
contain reducing elements (that is, those capable of combining with
oxygen), they often act as reducing agents--as, for instance, when
heated with oxide of copper, they burn, forming carbonic anhydride and
water, and leave metallic copper. Gerhardt proved that all hydrocarbons
contain an even number of hydrogen atoms. Therefore, the general formula
for all hydrocarbons is C_{_n_}H_{2_m_} where _n_ and _m_ are whole
numbers. This fact is known as _the law of even numbers_. Hence, the
simplest possible hydrocarbons ought to be: CH_{2}, CH_{4}, CH_{6} ...
C_{2}H_{2}, C_{2}H_{4}, C_{2}H_{6}, C_{2}H_{8} ... but they do not all
exist, since the quantity of H which can combine with a certain amount
of carbon is limited, as we shall learn directly.

  [21] The viscosity, or degree of mobility, of liquids is determined
       by their internal friction. It is estimated by passing the
       liquids through narrow (capillary) tubes, the mobile liquids
       passing through with greater facility and speed than the viscid
       ones. The viscosity varies with the temperature and nature of the
       liquids, and in the case of solutions changes with the amount of
       the substance dissolved, but is not proportional to it. So that,
       for example, with alcohol at 20° the viscosity will be 69, and
       for a 50 p.c. solution 160, the viscosity of water being taken as
       100. The volume of the liquid which passes through by experiment
       (Poiseuille) and theory (Stokes) is proportional to the time, the
       pressure, and the fourth power of the diameter of the (capillary)
       tube, and inversely proportional to the length of the tube;
       this renders it possible to form comparative estimates of the
       coefficients of internal friction and viscosity.

       As the complexity of the molecules of hydrocarbons and their
       derivatives increases by the addition of carbon (or CH_{2}), so
       does the degree of viscosity also rise. The extensive series
       of investigations referring to this subject still await the
       necessary generalisation. That connection which (already partly
       observed) ought to exist between the viscosity and the other
       physical and chemical properties, forces us to conclude that
       the magnitude of internal friction plays an important part in
       molecular mechanics. In investigating organic compounds and
       solutions, similar researches ought to stand foremost. Many
       observations have already been made, but not much has yet been
       done with them; the bare facts and some mechanical data exist,
       but their relation to molecular mechanics has not been cleared up
       in the requisite degree. It has already been seen from existing
       data that the viscosity at the temperature of the absolute
       boiling point becomes as small as in gases.

  [22] In a number of hydrocarbons and their derivatives such a
       substitution of metals for the hydrogen may be attained by
       indirect means. The property shown by acetylene, C_{2}H_{2},
       and its analogues, of forming metallic derivatives is in this
       respect particularly characteristic. Judging from the fact that
       carbon is an acid element (that is, gives an acid anhydride with
       oxygen), though comparatively slightly acid (for carbonic acid is
       not at all a strong acid and compounds of chlorine and carbon,
       even CCl_{4}, are not decomposed by water as is the case with
       phosphorus chloride and even silicic chloride and boric chloride,
       although they correspond with acids of but little energy),
       one might expect to find in the hydrogen of hydrocarbons this
       faculty for being substituted by metals. The metallic compounds
       which correspond with hydrocarbons are known under the name of
       organo-metallic compounds. Such, for instance, is zinc ethyl,
       Zn(C_{2}H_{5})_{2}, which corresponds with ethyl hydride or
       ethane, C_{2}H_{6}, in which two atoms of hydrogen have been
       exchanged for one of zinc.

  [23] Gaseous and volatile hydrocarbons decompose when passed through a
       heated tube. When hydrocarbons are decomposed by heating, the
       primary products are generally other more stable hydrocarbons,
       among which are acetylene, C_{2}H_{2}, benzene, C_{6}H_{6},
       naphthalene, C_{10}H_{8}, &c.

  [24] Wagner (1888) showed that when unsaturated hydrocarbons are shaken
       with a weak (1 p.c.) solution of potassium permanganate,
       KMnO_{4}, at ordinary temperatures, they form glycols--for
       example, C_{2}H_{4} yields C_{2}H_{6}O_{2}.

Some of the hydrocarbons are capable of combination, whilst others
do not show that power. Those which contain less hydrogen belong to
the former category, and those which, for a given quantity of carbon,
contain the maximum amount of hydrogen, belong to the latter. The
composition of those last mentioned is expressed by the general formula
C_{_n_}H_{2_n_ + 2}. These so-called _saturated hydrocarbons_ are
incapable of combination.[25] The hydrocarbons CH_{6}, C_{2}H_{8},
C_{3}H_{10}, &c.... do not exist. Those containing the maximum amount
of hydrogen will be represented by CH_{4} (_n_ = 1, 2_n_ + 2 = 4),
C_{2}H_{6} (_n_ = 2), C_{3}H_{8} (n = 3), C_{4}H_{10}, &c. This may be
termed the _law of limits_. Placing this in juxtaposition with the law
of even numbers, it is easy to perceive that the possible hydrocarbons
can be ranged in series, the terms of which may be expressed by the
general formulæ C_{_n_}H_{2_n_+2}, C_{_n_}H_{2_n_}, C_{_n_}H_{2_n_-2},
&c.... Those hydrocarbons which belong to any one of the series
expressible by a general formula are said to be _homologous_0 with
one another. Thus, the hydrocarbons CH_{4}, C_{2}H_{6}, C_{3}H_{8},
C_{4}H_{10}, &c.... are members of the limiting (saturated) homologous
series C_{_n_}H_{2_n_+2}. That is, the difference between the members
of the series is CH_{2}.[26] Not only the composition but also the
properties of the members of a series tend to classification in one
group. For instance, the members of the series C_{_n_}H_{2_n_+2}
are not capable of forming additive compounds, whilst those of the
series C_{_n_}H_{2_n_} are capable of combining with chlorine,
sulphuric anhydride, &c.; and the members of the C_{_n_}H_{2_n_-6}
group, belonging to the coal tar series, are easily nitrated (give
nitro-compounds, Chapter VI.), and have other properties in common.
The physical properties of the members of a given homologous series
vary in some such manner as this; the boiling point generally rises and
the internal friction increases as _n_ increases[27]--that is, with an
increase in the relative amount of carbon and the atomic weight; the
specific gravity also regularly changes as _n_ becomes greater.[28]

  [25] My article on this subject appeared in the Journal of the
       St. Petersburg Academy of Sciences in 1861. Up to that time,
       although many additive combinations with hydrocarbons and their
       derivatives were known, they had not been generalised, and were
       even continually quoted as cases of substitution. Thus the
       combination of ethylene, C_{2}H_{4}, with chlorine, Cl_{2}, was
       often regarded as a formation of the products of the substitution
       of C_{2}H_{5}Cl and HCl, which it was supposed were held together
       as the water of crystallisation is in salts. Even earlier than
       this (1857, _Journal of the Petroffsky Academy_) I considered
       similar cases as true compounds. In general, according to the
       law of limits, an unsaturated hydrocarbon, or its derivative, on
       combining with _r_X_{2}, gives a substance which is saturated or
       else approaching the limit. The investigations of Frankland with
       many organo-metallic compounds clearly showed the limit in the
       case of metallic compounds, which we shall constantly refer to
       later on.

  [26] The conception of homology has been applied by Gerhardt to all
       organic compounds in his classical work, 'Traité de Chimie
       Organique,' finished in 1855 (4 vols.), in which he divided
       all organic compounds into _fatty_ and _aromatic_, which is in
       principle still adhered to at the present time, although the
       latter are more often called benzene derivatives, on account of
       the fact that Kekulé, in his beautiful investigations on the
       structure of aromatic compounds, showed the presence in them all
       of the 'benzene nucleus,' C_{6}H_{6}.

  [27] This is always true for hydrocarbons, but for derivatives of the
       lower homologues the law is sometimes different; for instance,
       in the series of saturated alcohols, C_{_n_}H_{2_n_+1}(OH),
       when _n_ = 0, we obtain water, H(OH), which boils at 100°, and
       whose specific gravity at 15° = 0·9992; when _n_ = 1, wood
       spirit CH_{3}(OH), which boils at 66°, and at 15° has a specific
       gravity = 0·7964; when _n_ = 2, ordinary alcohol, C_{2}H_{5}(OH),
       boiling at 78°, specific gravity at 15° = 0·7936, and with
       further increase of CH_{2} the specific gravity increases. For
       the glycols C_{_n_}H_{2_n_}(OH)_{2} the phenomenon of a similar
       kind is still more striking; at first the temperature of the
       boiling point and the density increase, and then for higher (more
       complex) members of the series diminish. The reason for this
       phenomenon, it is evident, must be sought for in the influence
       and properties of water, and that strong affinity which, acting
       between hydrogen and oxygen, determines many of the exceptional
       properties of water (Chapter I.).

  [28] As, for example, in the saturated series of hydrocarbons
       C_{_n_}H_{2_n_+2}, the lowest member (_n_ = 0) must be taken as
       hydrogen H_{2}, a gas which (_t.c._ below -190°) is liquefied
       with great difficulty, and when in a liquid state has doubtless
       a very small density. Where _n_ = 1, 2, 3, the hydrocarbons
       CH_{4}, C_{2}H_{6}, C_{3}H_{8} are gases, more and more readily
       liquefiable. The temperature of the absolute boiling point for
       CH_{4} =-100°, and for ethane C_{2}H_{6}, and in the higher
       members it rises. The hydrocarbon C_{4}H_{10}, liquefies at about
       0°. C_{5}H_{12} (there are several isomers) boils at from +9°
       (Lvoff) to 37°, C_{6}H_{14} from 58° to 78°, &c. The specific
       gravities in a liquid state at 15° are:--

       C_{5}H_{12}  C_{6}H_{14}  C_{7}H_{16}  C_{10}H_{22}  C_{16}H_{34}
          0·63         0·66         0·70          0·75          0·85

Many of the hydrocarbons met with in nature are the products of
organisms, and do not belong to the mineral kingdom. A still greater
number are produced artificially. These are formed by what is termed
the combination of residues. For instance, if a mixture of the vapours
of hydrogen sulphide and carbon bisulphide be passed through a tube
in which copper is heated, this latter absorbs the sulphur from both
the compounds, and the liberated carbon and hydrogen combine to form
a hydrocarbon, methane. If carbon be combined with any metal and this
compound MC_{_n_} be treated with an acid HX, then the haloid X will
give a salt with the metal and the residual carbon and hydrogen will
give a hydrocarbon. Thus cast iron which contains a compound of iron
and carbon gives liquid hydrocarbons like naphtha under the action
of acids. If a mixture of bromo-benzene, C_{6}H_{5}Br, and ethyl
bromide, C_{2}H_{5}Br, be heated with metallic sodium, the sodium
combines with the bromine of both compounds, forming sodium bromide,
NaBr. From the first combination the group C_{6}H_{5} remains, and
from the second C_{2}H_{5}. Having an odd number of hydrogen atoms,
they, in virtue of the law of even numbers, cannot exist alone, and
therefore combine together forming the compound C_{6}H_{5}.C_{2}H_{5}
or C_{8}H_{10} (ethylbenzene). Hydrocarbons are also produced by the
breaking up of more complex organic or hydrocarbon compounds, especially
by heating--that is, by dry distillation. For instance, gum-benzoin
contains an acid called benzoic acid, C_{7}H_{6}O_{2}, the vapours
of which, when passed through a heated tube, split up into carbonic
anhydride, CO_{2}, and benzene, C_{6}H_{6}. Carbon and hydrogen only
unite directly in one ratio of combination--namely, to form acetylene,
having the composition C_{2}H_{2}, which, as compared with other
hydrocarbons, exhibits a very great stability at a somewhat high
temperature.[29]

  [29] If, at the ordinary temperature (assuming therefore that the
       water formed will be in a liquid state) a gram molecule (26
       grams) of acetylene, C_{2}H_{2}, be burnt, 310 thousand calories
       will be emitted (Thomsen), and as 12 grams of charcoal produce 97
       thousand calories, and 2 grams of hydrogen 69 thousand calories,
       it follows that, if the hydrogen and carbon of the acetylene were
       burnt there would be only 2 × 97 + 69, or 263 thousand calories
       produced. It is evident, then, that acetylene in its formation
       absorbs 310-263, or 47 thousand calories.

       For considerations relative to the combustion of carbon
       compounds, we will first enumerate the quantity of heat separated
       by the combustion of definite chemical carbon compounds, and then
       give a few figures bearing on the kinds of fuel used in practice.

       For molecular quantities in perfect combustion the following
       amounts of heat are given out (when gaseous carbonic anhydride
       and liquid water are formed), according to Thomsen's data (1)
       for gaseous C_{_n_}H_{2_n_ + 2}: 52·8 + 158·8_n_ thousand
       calories; (2) for C_{_n_}H_{2_n_}: 17·7 + 158·1_n_ thousand
       calories; (3) according to Stohmann (1888) for liquid saturated
       alcohols, C_{_n_}H_{2_n_ + 2}O: 11·8 + 156·3_n_, and as the
       latent heat of evaporation = about 8·2 + 0·6_n_, in a gaseous
       state, 20·0 + 156·9_n_; (4) for monobasic saturated liquid
       acids, C_{_n_}H_{2_n_}O_{2}:--95·3 + 154·3_n_, and as their
       latent heat of evaporation is about 5·0 + 1·2_n_, in a gaseous
       form, about--90 + 155_n_; (5) for solid saturated bibasic acids,
       C_{_n_}H_{2_n_-2}O_{4}:--253·8 + 152·6_n_, if they are expressed
       as C_{_n_}H_{2_n_}C_{2}H_{2}O_{4}, then 51·4 + 152·6_n_; (6)
       for benzene and its liquid homologues (still according to
       Stohmann) C_{_n_}H_{2_n_-6}:--158·6 + 156·3_n_, and in a gaseous
       form about--155 + 157_n_; (7) for the gaseous homologues of
       acetylene, C_{_n_}H_{2_n_-2} (according to Thomsen)--5 + 157_n_.
       It is evident from the preceding figures that the group CH_{2},
       or CH_{3} substituted for H, on burning gives out from 152 to
       159 thousand calories. This is less than that given out by C +
       H_{2}, which is 97 + 69 or 166 thousand; the reason for this
       difference (it would be still greater if carbon were gaseous)
       is the amount of heat separated during the formation of CH_{2}.
       According to Stohmann, for dextroglucose, C_{6}H_{12}O_{6}, it
       is 673·7; for common sugar, C_{12}H_{22}O_{11}, 1325·7; for
       cellulose, C_{6}H_{10}O_{5}, 678·0; starch, 677·5; dextrin,
       666·2; glycol, C_{2}H_{6}O_{2}, 281·7; glycerine, 397·2, &c.
       The heat of combustion of the following solids (determined
       by Stohmann) is expressed per unit of weight: naphthalene,
       C_{10}H_{8}, 9,621; urea, CN_{2}H_{4}O, 2,465; white of egg,
       5,579; dry rye bread, 4,421; wheaten bread, 4,302; tallow, 9,365;
       butter, 9,192; linseed oil, 9,323. The most complete collection
       of arithmetical data for the heats of combustion will be found
       in V. F. Longinin's work, 'Description of the Various Methods
       of Determining the Heats of Combustion of Organic Compounds'
       (Moscow, 1894).

       The number of units of heat given out by _unit weight_ during the
       complete combustion and cooling of the following ordinary kinds
       of fuel in their usual state of dryness and purity are:--(1) for
       wood charcoal, anthracite, semi-anthracite, bituminous coal and
       coke, from 7,200 to 8,200; (2) dry, long flaming coals, and the
       best brown coals, from 6,200 to 6,800; (3) perfectly dry wood,
       3,500; hardly dry, 2,500; (4) perfectly dry peat, best kind,
       4,500; compressed and dried, 3,000; (5) petroleum refuse and
       similar liquid hydrocarbons, about 11,000; (6) illuminating gas
       of the ordinary composition (about 45 vols. H, 40 vols. CH_{4}, 5
       vols. CO, and 5 vols. N), about 12,000; (7) producer gas (_see_
       next Chapter), containing 2 vols. carbonic anhydride, 30 vols.
       carbonic oxide, and 68 vols. nitrogen _for one part by weight of
       the whole carbon burnt_, 5,300, and for one part by weight of the
       gas, 910, units of heat; and (8) water gas (_see_ next chapter)
       containing 4 vols. carbonic anhydride, 8 vols. N_{2}, 24 vols.
       carbonic oxide, and 46 vols. H_{2}, for one part by weight of
       the carbon consumed in the _generator_ 10,900, and for one part
       by weight of the gas, 3,600 units of heat. In these figures,
       as in all calorimetric observations, the water produced by the
       combustion of the fuel is supposed to be liquid. As regards the
       temperature reached by the fuel, it is important to remark that
       for solid fuel it is indispensable to admit (to ensure complete
       combustion) twice the amount of air required, but liquid, or
       pulverised fuel, and especially gaseous fuel, does not require an
       excess of air; therefore, a kilogram of charcoal, giving 8,000
       units of heat, requires about 24 kilograms of air (3 kilograms
       of air per thousand calories) and a kilogram of producer gas
       requires only 0·77 kilogram of air (0·85 kilo. of air per 1,000
       calories), 1 kilogram of water gas about 4·5 of air (1·25 kilo.
       of air per 1,000 calories).

There is one substance known among the saturated hydrocarbons composed
of 1 atom of carbon and 4 atoms of hydrogen; this is a compound
containing the highest percentage of hydrogen (CH_{4} contains 25 per
cent. of hydrogen), and at the same time it is the only hydrocarbon
whose molecule contains but a single atom of carbon. This saturated
hydrocarbon, CH_{4}, is called _marsh gas_ or _methane_. If vegetable
or animal refuse suffers decomposition in a space where the air has not
free access, or no access at all, then the decomposition is accompanied
with the formation of marsh gas, and this either at the ordinary
temperature, or at a comparatively much higher one. On this account
_plants_, when decomposing under water in _marshes_, give out this
gas.[29 bis] It is well known that if the mud in bogs be stirred up,
the act is accompanied with the evolution of a large quantity of gas
bubbles; these may, although slowly, also separate of their own accord.
The gas which is evolved consists principally of marsh gas.[30] If
wood, coal, or many other vegetable or animal substances are decomposed
by the _action of heat_ without access of air--that is, are subjected
to dry distillation--they, in addition to many other gaseous products
of decomposition (carbonic anhydride, hydrogen, and various other
substances), evolve a great deal of methane. Generally the gas which
is used for lighting purposes is obtained by this means and therefore
always contains marsh gas, mixed with dry hydrogen and other vapours and
gases, although it is subsequently purified from many of them.[31]
As the decomposition of the organic matter, which forms coal, is still
going on underground, the evolution of large quantities of marsh gas
frequently occurs in coal-mines.[32] When mixed with air it forms
an explosive mixture, which forms one of the great dangers of coal
mining, as subterranean work has always to be carried on by lamp-light.
This danger is, however, overcome by the use of Humphry Davy's safety
lamp.[33] Sir Humphry Davy observed that on introducing a piece of
wire gauze into a flame, it absorbs so much heat that combustion does
not proceed beyond it (the unburnt gases which pass through it may
be ignited on the other side). In accordance with this, the flame
of the Davy lamp is surrounded with a thick glass (as shown in the
drawing), and has no communication whatever with the explosive mixture
except through a wire gauze which prevents it igniting the mixture
of the marsh-gas issuing from the coal with air. In some districts,
particularly in those where petroleum is found--as, for instance, near
Baku, where a temple of the Indian fire-worshippers was built, and in
Pennsylvania, and other places--marsh gas in abundance issues from the
earth, and it is used, like coal gas, for the purposes of lighting and
warming.[34] Tolerably pure marsh gas[35] may be obtained by heating
a mixture of an acetate with an alkali. Acetic acid, C_{2}H_{4}O_{2},
on being heated is decomposed into marsh gas and carbonic anhydride,
C_{2}H_{4}O_{2} = CH_{4} + CO_{2}.

  [29 bis] Manure which decomposes under the action of bacteria gives off
       CO_{2} and CH_{4}.

  [30] It is easy to collect the gas which is evolved in marshy places
       if a glass bottle be inverted in the water and a funnel put into
       it (both filled with water); if the mud of the bottom be now
       agitated, the bubbles which rise may be easily caught by the
       inverted funnel.

  [31] [Illustration: FIG. 58.--General view of gas works. _B_, retorts;
       _f_, hydraulic main; _H_ and _I_, tar well; _i_, condensers; _L_,
       purifiers; _P_, gasholder.]

       Illuminating gas is generally prepared by heating gas coal
       (_see_ Note 6) in oval cylindrical horizontal cast-iron or clay
       retorts. Several such retorts _BB_ (fig. 58) are disposed in the
       furnace _A_, and heated together. When the retorts are heated
       to a red heat, lumps of coal are thrown into them, and they
       are then closed with a closely fitting cover. The illustration
       shows the furnace, with five retorts. Coke (_see_ Note 1, dry
       distillation) remains in the retorts, and the volatile products
       in the form of vapours and gases travel along the pipe _d_,
       rising from each retort. These pipes branch above the stove, and
       communicate with the receiver _f_ (hydraulic main) placed above
       the furnace. Those products of the dry distillation which most
       easily pass from the gaseous into the liquid and solid states
       collect in the hydraulic main. From the hydraulic main the
       vapours and gases travel along the pipe _g_ and the series of
       vertical pipes _j_ (which are sometimes cooled by water trickling
       over the surface), where the vapours and gases cool from the
       contact of the colder surface, and a fresh quantity of vapour
       condenses. The condensed liquids pass from the pipes _g_ and _j_
       and into the troughs _H_. These troughs always contain liquid
       at a constant level (the excess flowing away) so that the gas
       cannot escape, and thus they form, as it is termed, a hydraulic
       joint. In the state in which it leaves the condensers the gas
       consists principally of the following vapours and gases: (1)
       vapour of water, (2) ammonium carbonate, (3) liquid hydrocarbons,
       (4) hydrogen sulphide, H_{2}S, (5) carbonic anhydride, CO_{2},
       (6) carbonic oxide, CO, (7) sulphurous anhydride, SO_{2}, but
       a great part of the illuminating gas consists of (8) hydrogen,
       (9) marsh gas, (10) olefiant gas, C_{2}H_{4}, and other gaseous
       hydrocarbons. The hydrocarbons (3, 9, and 10), the hydrogen,
       and carbonic oxide are capable of combustion, and are useful
       component parts, but the carbonic anhydride, the hydrogen
       sulphide, and sulphurous anhydride, as well as the vapours of
       ammonium carbonate, form an injurious admixture, because they
       do not burn (CO_{2}, SO_{2}) and lower the temperature and
       brilliancy of the flame, or else, although capable of burning
       (for example, H_{2}S, CS_{2}, and others), they give out during
       combustion sulphurous anhydride which has a disagreeable smell,
       is injurious when inhaled, and spoils many surrounding objects.
       In order to separate the injurious products, the gas is washed
       with water, a cylinder (not shown in the illustration) filled
       with coke continually moistened with water serving for this
       purpose. The water coming into contact with the gas dissolves
       the ammonium carbonate; hydrogen sulphide, carbonic anhydride,
       and sulphurous anhydride, being only partly soluble in water,
       have to be got rid of by a special means. For this purpose the
       gas is passed through moist lime or other alkaline liquid, as
       the above-mentioned gases have acid properties and are therefore
       retained by the alkali. In the case of lime, calcium carbonate,
       sulphite and sulphide, all solid substances, are formed. It is
       necessary to renew the purifying material as its absorbing power
       decreases. A mixture of lime and sulphate of iron, FeSO_{4},
       acts still better, because the latter, with lime, Ca(HO)_{2},
       forms ferrous hydroxide, Fe(HO)_{2} and gypsum, CaSO_{4}. The
       suboxide (partly turning into oxide) of iron absorbs H_{2}S,
       forming FeS and H_{2}O, and the gypsum retains the remainder of
       the ammonia, the excess of lime absorbing carbonic anhydride
       and sulphuric anhydride. [In English works a native hydrated
       ferric hydroxide is used for removing hydrogen sulphide.] This
       purification of the gas takes place in the apparatus _L_, where
       the gas passes through perforated trays _m_, covered with
       sawdust mixed with lime and sulphate of iron. It is necessary
       to remark that in the manufacture of gas it is indispensable
       to draw off the vapours from the retorts, so that they should
       not remain there long (otherwise the hydrocarbons would in a
       considerable degree be resolved into charcoal and hydrogen),
       and also to avoid a great pressure of gas in the apparatus,
       otherwise a quantity of gas would escape at all cracks such as
       must inevitably exist in such a complicated arrangement. For
       this purpose there are special pumps (exhausters) so regulated
       that they only pump off the quantity of gas formed (the pump is
       not shown in the illustration). The purified gas passes through
       the pipe _n_ into the gasometer (gasholder) _P_, a dome made of
       iron plate. The edges of the dome dip into water poured into a
       ring-shaped channel _g_, in which the sides of the dome rise and
       fall. The gas is collected in this holder, and distributed to its
       destination by pipes communicating with the pipe _o_, issuing
       from the dome. The pressure of the dome on the gas enables it, on
       issuing from a long pipe, to penetrate through the small aperture
       of the burner. A hundred kilograms of coal give about 20 to 30
       cubic metres of gas, having a density from four to nine times
       greater than that of hydrogen. A cubic metre (1,000 litres) of
       hydrogen weighs about 87 grams; therefore 100 kilograms of coal
       give about 18 kilograms of gas, or about one-sixth of its weight.
       Illuminating gas is generally lighter than marsh gas, as it
       contains a considerable amount of hydrogen, and is only heavier
       than marsh gas when it contains much of the heavier hydrocarbons.
       Thus olefiant gas, C_{2}H_{4}, is fourteen times, and the
       vapours of benzene thirty-nine times, heavier than hydrogen, and
       illuminating gas sometimes contains 15 p.c. of its volume of
       them. The brilliancy of the flame of the gas increases with the
       quantity of olefiant gas and similar heavy hydrocarbons, as it
       then contains more carbon for a given volume and a greater number
       of carbon particles are separated. Gas usually contains from 35
       to 60 p.c. of its volume of marsh gas, from 30 to 50 p.c. of
       hydrogen, from 3 to 5 p.c. of carbonic oxide, from 2 to 10 p.c.
       heavy hydrocarbons, and from 3 to 10 p.c. of nitrogen. Wood gives
       almost the same sort of gas as coal and almost the same quantity,
       but the wood gas contains a great deal of carbonic anhydride,
       although on the other hand there is an almost complete absence
       of sulphur compounds. Tar, oils, naphtha, and such materials
       furnish a large quantity of good illuminating gas. An ordinary
       burner of 8 to 12 candle-power burns 5 to 6 cubic feet of coal
       gas per hour, but only 1 cubic foot of naphtha gas. One pood (36
       lbs. Eng.) of naphtha gives 500 cubic feet of gas--that is, one
       kilogram of naphtha produces about one cubic metre of gas. The
       formation of combustible gas by heating coal was discovered in
       the beginning of the last century, but only put into practice
       towards the end by Le-Bon in France and Murdoch in England. In
       England, Murdoch, together with the renowned Watt, built the
       first gas works in 1805.

       [Illustration: FIG. 59.--Blowpipe. Air is blown in at the
       trumpet-shaped mouthpiece, and escapes in a fine stream from the
       platinum jet placed at the extremity of the side tube.]

       [Illustration: FIG. 60.--Davy safety-lamp. [Modern form.]]

       In practice illuminating gas is not only used for lighting
       (electricity and kerosene are cheaper in Russia), but also as
       the motive power for gas engines (_see_ p. 175), which consume
       about half a cubic metre per horse-power per hour; gas is also
       used in laboratories for heating purposes. When it is necessary
       to concentrate the heat, either the ordinary blowpipe (fig. 59)
       is applied, placing the end in the flame and blowing through
       the mouthpiece; or, in other forms, gas is passed through the
       blowpipe; when a large, hot, smokeless flame is required for
       heating crucibles or glass-blowing, a foot-blower is used.
       High temperatures, which are often required for laboratory and
       manufacturing purposes, are most easily attained by the use of
       gaseous fuel (illuminating gas, producer gas, and water gas,
       which will be treated of in the following chapter), because
       complete combustion may be effected without an access of air. It
       is evident that in order to obtain high temperatures means must
       be taken to diminish the loss of heat by radiation, and to ensure
       perfect combustion.

  [32] The gas which is set free in coal mines contains a good deal
       of nitrogen, some carbonic anhydride, and a large quantity of
       marsh gas. The best means of avoiding an explosion consists
       in efficient ventilation. It is best to light coal mines with
       electric lamps.

  [33] The Davy lamp, of which an improved form is represented in the
       accompanying figure, is used for lighting coal and other mines
       where combustible gas is found. The wick of the lamp is enclosed
       in a thick glass cylinder which is firmly held in a metallic
       holder. Over this a metallic cylinder and the wire gauze are
       placed. The products of combustion pass through the gauze, and
       the air enters through the space between the cylinder and the
       wire gauze. To ensure greater safety the lamp cannot be opened
       without extinguishing the flame.

  [34] In Pennsylvania (beyond the Alleghany mountains) many of the
       shafts sunk for petroleum only emitted gas, but many useful
       applications for it were found and it was conducted in metallic
       pipes to works hundreds of miles distant, principally for
       metallurgical purposes.

  [35] The purest gas is prepared by mixing the liquid substance called
       zinc methyl, Zn(CH_{3})_{2}, with water, when the following
       reaction occurs:

             Zn(CH_{3})_{2} + 2HOH = Zn(HO)_{2} + 2CH_{3}H.

An alkali--for instance, NaHO--gives with acetic acid a salt,
C_{2}H_{3}NaO_{2}, which on decomposition retains carbonic anhydride,
forming a carbonate, Na_{2}CO_{3}, and marsh gas is given off:

            C_{2}H_{3}NaO_{2} + NaHO = Na_{2}CO_{3} + CH_{4}

Marsh gas is difficult to liquefy; it is almost insoluble in water,
and is without taste or smell. The most important point in connection
with its chemical reactions is that it does not combine directly with
anything, whilst the other hydrocarbons which contain less hydrogen than
expressed by the formula C_{_n_}H_{2_n_ + 2} are capable of combining
with hydrogen, chlorine, certain acids, &c.

If the law of substitution gives a very simple explanation of the
formation of hydrogen peroxide as a compound containing two aqueous
residues (OH)(OH), then on the basis of this law all hydrocarbons
ought to be derived from methane, CH_{4}, as being the simplest
hydrocarbon.[36] The increase in complexity of a molecule of methane
is brought about by the faculty of mutual combination which exists in
the atoms of carbon, and, as a consequence of the most detailed study
of the subject, much that might have been foreseen and conjectured
from the law of substitution has been actually brought about in such
a manner as might have been predicted, and although this subject on
account of its magnitude really belongs, as has been already stated,
to the sphere of organic chemistry, it has been alluded to here in
order to show, although only in part, the best investigated example of
the application of the law of substitution. According to this law, a
molecule of methane, CH_{4}, is capable of undergoing substitution in
the four following ways:--(1) Methyl substitution, when the radicle,
equivalent to hydrogen, called _methyl_ CH_{3}, replaces hydrogen.
In CH_{4} this radicle is combined with H and therefore can replace
it, as (OH) replaces H because with it it gives water; (2) methylene
substitution, or the exchange between H_{2} and CH_{2} (this radicle
is called methylene), is founded on a similar division of the molecule
CH_{4} into two equivalent parts, H_{2} and CH_{2}; (3) acetylene
substitution, or the exchange between CH on the one hand and H_{3}
on the other; and (4) carbon substitution--that is, the substitution
of H_{4} by an atom of carbon C, which is founded on the law of
substitution just as is the methyl substitution. These four cases of
substitution render it possible to understand the principal relations of
the hydrocarbons. For instance, the _law of even numbers_ is seen from
the fact that in all the cases of substitution mentioned the hydrogen
atoms increase or decrease by an even number; but as in CH_{4} they are
likewise even, it follows that no matter how many substitutions are
effected there will always be obtained an even number of hydrogen atoms.
When H is replaced by CH_{3} there is an increase of CH_{2}; when H_{2}
is replaced by CH_{2} there is no increase of hydrogen; in the acetylene
substitution CH replaces H_{3}, therefore there is an increase of C and
a decrease of H_{2}; in the carbon substitution there is a decrease
of H_{4}. In a similar way the _law of limit_ may be deduced as a
corollary of the law of substitution. For the largest possible quantity
of hydrogen is introduced by the methyl substitution, since it leads
to the addition of CH_{2}; starting from CH_{4} we obtain C_{2}H_{6},
C_{3}H_{8}, and in general, C_{_n_}H_{2_n_+2}, and these contain
the greatest possible amount of hydrogen. Unsaturated hydrocarbons,
containing less hydrogen, are evidently only formed when the increase
of the new molecule derived from methane proceeds from one of the
other forms of substitution. When the methyl substitution alone takes
place in methane, CH_{4}, it is evident that the saturated hydrocarbon
formed is C_{2}H_{6} or (CH_{3})(CH_{3}).[37] This is called _ethane_.
By means of the methylene substitution alone, _ethylene_, C_{2}H_{4},
or (CH_{2})(CH_{2}) may be directly obtained from CH_{4}, and by the
acetylene substitution C_{2}H_{2} or (CH)(CH), or _acetylene_, both the
latter being unsaturated hydrocarbons. Thus we have all the possible
hydrocarbons with two atoms of carbon in the molecule, C_{2}H_{6},
ethane, C_{2}H_{4}, ethylene, and C_{2}H_{2}, acetylene. But in them,
according to the law of substitution, the same forms of substitution may
be repeated--that is, the methyl, methylene, acetylene, and even carbon
substitutions (because C_{2}H_{6} will still contain hydrogen when C
replaces H_{4}) and therefore further substitutions will serve as a
source for the production of a fresh series of saturated and unsaturated
hydrocarbons, containing more and more carbon in the molecule and,
in the case of the acetylene substitution and carbon substitution,
containing less and less hydrogen. Thus _by means of the law of
substitution we can foresee_ not only the limit C_{_n_}H_{2_n_+2},
but an unlimited number of unsaturated hydrocarbons, C_{_n_}H_{2_n_},
C_{_n_}H_{2_n_-2} ... C_{_n_}H_{2(_n-m_)}, where _m_ varies from 0
to _n_-1,[38] and where _n_ increases indefinitely. From these facts
not only does the existence of a multitude of polymeric hydrocarbons,
differing in molecular weight, become intelligible, but it is also
seen that there is a possibility of cases of isomerism with the same
molecular weight. This _polymerism_ so common to hydrocarbon compounds
is already apparent in the first unsaturated series C_{_n_}H_{2_n_},
because all the terms of this series C_{2}H_{4}, C_{3}H_{6}, C_{4}H_{8}
... C_{30}H_{60} ... have one and the same composition CH_{2}, but
different molecular weights, as has been already explained in Chapter
VII. The differences in the vapour density, boiling points, and
melting points, of the quantities entering into reactions,[39] and the
methods of preparation[40] also so clearly tally with the conception
of polymerism, that this example will always be the clearest and most
conclusive for the illustration of polymerism and molecular weight.
Such a case is also met with among other hydrocarbons. Thus benzene,
C_{6}H_{6}, and cinnamene, C_{8}H_{8}, correspond with the composition
of acetylene or to a compound of the composition CH.[41] The first boils
at 81°, the second at 144°; the specific gravity of the first is 0·899;
that of the second, 0·925, at 0°--that is, here also the boiling point
rises with the increase of molecular weight, and so also, as might be
expected, does the density.

  [36] Methylene, CH_{2}, does not exist. When attempts are made to
       obtain it (for example, by removing X_{2} from CH_{2}X_{2}),
       C_{2}H_{4} or C_{3}H_{6} are produced--that is to say, it
       undergoes polymerisation.

  [37] Although the methods of formation and the reactions connected with
       hydrocarbons are not described in this work, because they are
       dealt with in organic chemistry, yet in order to clearly show
       the mechanism of those transformations by which the carbon
       atoms are built up into the molecules of the carbon compounds,
       we here give a general example of reactions of this kind. From
       marsh gas, CH_{4}, on the one hand the substitution of chlorine
       or iodine, CH_{3}Cl, CH_{3}I, for the hydrogen may be effected,
       and on the other hand such metals as sodium may be substituted
       for the hydrogen, _e.g._ CH_{3}Na. These and similar products of
       substitution serve as a means of obtaining other more complex
       substances from given carbon compounds. If we place the two
       above-named products of substitution of marsh gas (metallic and
       haloid) in mutual contact, the metal combines with the halogen,
       forming a very stable compound--namely, common salt, NaCl, and
       the carbon groups which were in combination with them separate in
       mutual combination, as shown by the equation:

                CH_{3}Cl + CH_{3}Na = NaCl + C_{2}H_{6}.

       This is the most simple example of the formation of a complex
       hydrocarbon from these radicles. The cause of the reaction must
       be sought for in the property which the haloid (chlorine) and
       sodium have of entering into mutual combination.

  [38] When _m_ = _n_-1, we have the series C_{_n_}H_{2}. The lowest
       member is acetylene, C_{2}H_{2}. These are hydrocarbons
       containing a minimum amount of hydrogen.

  [39] For instance, ethylene, C_{2}H_{4}, combines with Br_{2}, HI,
       H_{2}SO_{4}, as a whole molecule, as also does amylene,
       C_{5}H_{10}, and, in general, C_{_n_}H_{2_n_}.

  [40] For instance, ethylene is obtained by removing the water from
       ethyl alcohol, C_{2}H_{5}(OH), and amylene, C_{5}H_{10}, from
       amyl alcohol, C_{5}H_{11}(OH), or in general C_{_n_}H_{2_n_},
       from C_{_n_}H_{2_n_+1}(OH).

  [41] Acetylene and its polymerides have an empirical composition CH,
       ethylene and its homologues (and polymerides) CH_{2}, ethane
       CH_{3}, methane CH_{4}. This series presents a good example of
       the law of multiple proportions, but such diverse proportions
       are met with between the number of atoms of the carbon and
       hydrogen in the hydrocarbons already known that the accuracy of
       Dalton's law might be doubted. Thus the substances C_{30}H_{62}
       and C_{30}H_{60} differ so slightly in their composition by
       weight as to be within the limits of experimental error, but
       their reactions and properties are so distinct that they can be
       distinguished beyond a doubt. Without Dalton's law chemistry
       could not have been brought to its present condition, but it
       cannot alone express all those gradations which are quite clearly
       understood and predicted by the law of Avogadro-Gerhardt.

Cases of isomerism in the restricted sense of the word--that is,
when with an identity of composition and of molecular weight, the
properties of the substances are different--are very numerous among
the hydrocarbons and their derivatives. Such cases are particularly
important for the comprehension of molecular structure and they also,
like the polymerides, may be predicted from the above-mentioned
conceptions, expressing the principles of the structure of the carbon
compounds[42] based on the law of substitution. According to it, for
example, it is evident that there can be no isomerism in the cases
of the saturated hydrocarbons C_{2}H_{6} and C_{3}H_{8}, because the
former is CH_{4}, in which methyl has taken the place of H, and as
all the hydrogen atoms of methane must be supposed to have the same
relation to the carbon, it is all the same which of them be subjected
to the methyl substitution--the resulting product can only be ethane,
CH_{3}CH_{3};[43] the same argument also applies in the case of propane,
CH_{3}CH_{2}CH_{3}, where one compound only can be imagined. It is to
be expected, however, that there should be two butanes, C_{4}H_{10}, and
this is actually the case. In one, methyl may be considered as replacing
the hydrogen of one of the methyls, CH_{3}CH_{2}CH_{2}CH_{3}; and in
the other CH_{3} may be considered as substituted for H in /CH_{3}
CH_{2}, and there it will consist of CH_{3}CH. The latter may \CH_{3}
also be regarded as methane in which three of hydrogen are exchanged
for three of methyl. On going further in the series it is evident that
the number of possible isomerides will be still greater, but we have
limited ourselves to the simplest examples, showing the possibility and
actual existence of isomerides. C_{2}H_{4} and CH_{2}CH_{2} are, it is
evident, identical; but there ought to be, and are, two hydrocarbons of
the composition C_{3}H_{6}, propylene and trimethylene; the first is
ethylene, CH_{2}CH_{2}, in which one atom of hydrogen is exchanged for
methyl, CH_{2}CHCH_{3}, and trimethylene is ethane, CH_{3}CH_{3}, with
the substitution of methylene for two hydrogen atoms from two methyl
groups--that /CH_{2} is, CH_{2},[44] where the methylene introduced
is united to both \CH_{2} the atoms of carbon in CH_{3}CH_{3}. It is
evident that the cause of isomerism here is, on the one hand, the
difference of the amount of hydrogen in union with the particular
atoms of carbon, and, on the other, the different connection between
the several atoms of carbon. In the first case they may be said to
be chained together (more usually to form an 'open chain'), and in
the second case, to be locked together (to form a 'closed chain'
or 'ring'). Here also it is easily understood that on increasing
the quantity of carbon atoms the number of possible and existing
isomerides will greatly increase. If, at the same time, in addition
to the substitution of one of the radicles of methane for hydrogen a
further exchange of part of the hydrogen for some of the other groups
of elements X, Y ... occurs, the quantity of possible isomerides still
further increases in a considerable degree. For instance, there are
even two possible isomerides for the derivatives of ethane, C_{2}H_{6}:
if two atoms of the hydrogen be exchanged for X_{2}, one will have
the ethylene structure, CH_{2}XCH_{2}X, and the other an ethylidene
structure, CH_{3}CHX_{2}; such are, for instance, ethylene chloride,
CH_{2}ClCH_{2}Cl, and ethylidene chloride, CH_{3}CHCl_{2}. And as in the
place of the first atom of hydrogen not only metals may be substituted,
but Cl, Br, I, OH (the water radicle), NH_{2} (the ammonia radicle),
NO_{2} (the radicle of nitric acid), &c., so also in exchange for two
atoms of hydrogen O, NH, S, &c., may be substituted; hence it will be
understood that the quantity of isomerides is sometimes very great. It
is impossible here to describe how the isomerides are distinguished
from each other, in what reactions they occur, how and when one changes
into another, &c.; for this, taken together with the description of
the hydrocarbons already known, and their derivatives, forms a very
extensive and very thoroughly investigated branch of chemistry, called
_organic chemistry_. Enriched with a mass of closely observed phenomena
and strictly deduced generalisations, this branch of chemistry has
been treated separately for the reason that in it the hydrocarbon
groups are subjected to transformations which are not met with in such
quantity in dealing with any of the other elements or their hydrogen
compounds. It was important for us to show that notwithstanding the
great variety of the hydrocarbons and their products,[45] they are all
of them governed by the law of substitution, and referring our readers
for detailed information to works on organic chemistry, we will limit
ourselves to a short exposition of the properties of the two simplest
unsaturated hydrocarbons: ethylene, CH_{2}CH_{2}, and acetylene, CHCH,
and a short acquaintance with petroleum as the natural source of a mass
of hydrocarbons. _Ethylene, or olefiant gas_, C_{2}H_{4}, is the lowest
known member of the unsaturated hydrocarbon series of the composition
C_{_n_}H_{2_n_}. As in composition it is equal to two molecules of marsh
gas deprived of two molecules of hydrogen, it is evident that it might
be, and it actually can be, produced, although but in small quantities,
together with hydrogen, by heating marsh gas. On being heated, however,
olefiant gas splits up, first into acetylene and methane (3C_{2}H_{4} =
2C_{2}H_{2} + 2CH_{4}, Lewes, 1894), and at a higher temperature into
carbon and hydrogen; and therefore in those cases where marsh gas is
produced by heating, olefiant gas, hydrogen, and charcoal will also be
formed, although only in small quantities. The lower the temperature at
which complex organic substances are heated, the greater the quantity
of olefiant gas found in the gases given off; at a white heat it is
entirely decomposed into charcoal and marsh gas. If coal, wood, and more
particularly petroleum, tars, and fatty substances, are subjected to dry
distillation, they give off illuminating gas, which contains more or
less olefiant gas.

  [42] The conception of the structure of carbon compounds--that is, the
       expression of those unions and correlations which their atoms
       have in the molecules--was for a long time limited to the
       representation that organic substances contained complex
       radicles (for instance, ethyl C_{2}H_{5}, methyl CH_{3}, phenyl
       C_{6}H_{5}, &c.); then about the year 1840 the phenomena
       of substitution and the correspondence of the products of
       substitution with the primary bodies (nuclei and types) were
       observed, but it was not until about the year 1860 and later when
       on the one hand the teaching of Gerhardt about molecules was
       spreading, and on the other hand the materials had accumulated
       for discussing the transformations of the simplest hydrocarbon
       compounds, that conjectures began to appear as to the mutual
       connection of the atoms of carbon in the molecules of the complex
       hydrocarbon compounds. Then Kekulé and A. M. Butleroff began to
       formulate the connection between the separate atoms of carbon,
       regarding it as a quadrivalent element. Although in their
       methods of expression and in some of their views they differ
       from each other and also from the way in which the subject is
       treated in this work, yet the essence of the matter--namely,
       the comprehension of the causes of isomerism and of the union
       between the separate atoms of carbon--remains the same. In
       addition to this, starting from the year 1870, there appears a
       tendency which from year to year increases to discover the actual
       spacial distribution of the atoms in the molecules. Thanks to the
       endeavours of Le-Bel (1874), Van't Hoff (1874), and Wislicenus
       (1887) in observing cases of isomerism--such as the effect of
       different isomerides on the direction of the rotation of the
       plane of polarisation of light--this tendency promises much
       for chemical mechanics, but the details of the still imperfect
       knowledge in relation to this matter must be sought for in
       special works devoted to organic chemistry.

  [43] Direct experiment shows that however CH_{3}X is prepared (where
       X = for instance Cl, &c.) it is always one and the same
       substance. If, for example, in CX_{4}, X is gradually replaced by
       hydrogen until CH_{3}X is produced, or in CH_{4}, the hydrogen by
       various means is replaced by X, or else, for instance, if CH_{3}X
       be obtained by the decomposition of more complex compounds, the
       same product is always obtained.

       This was shown in the year 1860, or thereabout, by many methods,
       and is the fundamental conception of the structure of hydrocarbon
       compounds. If the atoms of hydrogen in methyl were not absolutely
       identical in value and position (as they are not, for instance,
       in CH_{3}CH_{2}CH_{3} or CH_{3}CH_{2}X), then there would be as
       many different forms of CH_{3}X as there were diversities in the
       atoms of hydrogen in CH_{4}. The scope of this work does not
       permit of a more detailed account of this matter. It is given in
       works on organic chemistry.

  [44] The union of carbon atoms in closed chains or rings was first
       suggested by Kekulé as an explanation of the structure and
       isomerism of the derivatives of benzene, C_{6}H_{6}, forming
       aromatic compounds (Note 26).

  [45] The following are the most generally known of the oxygenised but
       non-nitrogenous hydrocarbon derivatives. (1) the alcohols. These
       are hydrocarbons in which hydrogen is exchanged for hydroxyl
       (OH). The simplest of these is methyl alcohol, CH_{3}(OH),
       or wood spirit obtained by the dry distillation of wood. The
       common spirits of wine or ethyl alcohol, C_{2}H_{3}(OH), and
       glycol, C_{2}H_{4}(OH)_{2}, correspond with ethane. Normal
       propyl alcohol, CH_{3}CH_{2}CH_{2}(OH), and isopropyl alcohol,
       CH_{3}CH(OH)CH_{3}, propylene-glycol, C_{3}H_{6}(OH)_{2}, and
       glycerol, C_{3}H_{3}(OH)_{3} (which, with stearic and other
       acids, forms fatty substances), correspond with propane,
       C_{3}H_{8}. All alcohols are capable of forming water and
       ethereal salts with acids, just as alkalis form ordinary salts.
       (2) Aldehydes are alcohols minus hydrogen; for instance,
       acetaldehyde, C_{2}H_{4}O, corresponds with ethyl alcohol.
       (3) It is simplest to regard organic acids as hydrocarbons in
       which hydrogen has been exchanged for carboxyl (CO_{2}H), as
       will be explained in the following chapter. There are a number
       of intermediate compounds; for example, the aldehyde-alcohols,
       alcohol-acids (or hydroxy-acids), &c. Thus the hydroxy-acids are
       hydrocarbons in which some of the hydrogen has been replaced
       by hydroxyl, and some by carboxyl; for instance, lactic
       acid corresponds with C_{2}H_{6}, and has the constitution
       C_{2}H_{4}(OH)(CO_{2}H). If to these products we add the haloid
       salts (where H is replaced by Cl, Br, I), the nitro-compounds
       containing NO_{2} in place of H, the amides, cyanides, ketones,
       and other compounds, it will be readily seen what an immense
       number of organic compounds there are and what a variety of
       properties these substances have; this we see also from the
       composition of plants and animals.

Olefiant gas, almost free from other gases,[46] may be obtained from
ordinary alcohol (if possible, free from water) if it be mixed with
five parts of strong sulphuric acid and the mixture heated to slightly
above 100°. Under these conditions, the sulphuric acid removes the
elements of water from the alcohol, C_{2}H_{5}(OH), and gives olefiant
gas; C_{2}H_{6}O = H_{2}O + C_{2}H_{4}. The greater molecular weight
of olefiant gas compared with marsh gas indicates that it may be
comparatively easily converted into a liquid by means of pressure or
great cold; this may be effected, for example, by the evaporation
of liquid nitrous oxide. Its absolute boiling point is +10°, it
boils at -103° (1 atmosphere), liquefies at 0°, at a pressure of 43
atmospheres, and solidifies at -160°. Ethylene is colourless, has a
slight ethereal smell, is slightly soluble in water, and somewhat more
soluble in alcohol and in ether (in five volumes of spirit and six
volumes of ether).[47]

  [46] Ethylene bromide, C_{2}H_{4}Br_{2}, when gently heated in
       alcoholic solution with finely divided zinc, yields pure
       ethylene, the zinc merely taking up the bromine (Sabaneyeff).

  [47] Ethylene decomposes somewhat easily under the influence of the
       electric spark, or a high temperature. In this case the volume
       of the gas formed may remain the same when olefiant gas is
       decomposed into carbon and marsh gas, or may increase to double
       its volume when hydrogen and carbon are formed, C_{2}H_{4} =
       CH_{4} + C = 2C + 2H_{2}. A mixture of olefiant gas and oxygen is
       highly explosive; two volumes of this gas require six volumes of
       oxygen for its perfect combustion. The eight volumes thus taken
       then resolve themselves into eight volumes of the products of
       combustion, a mixture of water and carbonic anhydride, C_{2}H_{4}
       + 3O_{2} = 2CO_{2} + 2H_{2}O. On cooling after the explosion
       diminution of volume occurs because the water becomes liquid. For
       two volumes of the olefiant gas taken, the diminution will be
       equal to four volumes, and the same for marsh gas. The quantity
       of carbonic anhydride formed by both gases is not the same. Two
       volumes of marsh gas give only two volumes of carbonic anhydride,
       and two volumes of ethylene give four volumes of carbonic
       anhydride.

Like other unsaturated hydrocarbons, olefiant gas readily enters into
combination with certain substances, such as chlorine, bromine, iodine,
fuming sulphuric acid, or sulphuric anhydride, &c. If olefiant gas be
sealed up with a small quantity of sulphuric acid in a glass vessel, and
constantly agitated (as, for instance, by attaching it to the moving
part of a machine), the prolonged contact and repeated mixing causes
the olefiant gas, little by little, to combine with the sulphuric acid,
forming C_{2}H_{4}H_{2}SO_{4}. If, after this absorption, the sulphuric
acid be diluted with water and distilled, alcohol separates, which is
produced in this case by the olefiant gas combining with the elements of
water, C_{2}H_{4} + H_{2}O = C_{2}H_{6}O. In this reaction (Berthelot)
we see an excellent example of the fact that if a given substance, like
olefiant gas, is produced by the decomposition of another, then in the
reverse way this substance, entering into combination, is capable of
forming the original substance--in our example, alcohol. In combination
with various molecules, X_{2}, ethylene gives saturated compounds,
C_{2}H_{4}X_{2} or CH_{2}XCH_{2}X (for example, C_{2}H_{4}Cl_{2}), which
correspond with ethane, CH_{3}CH_{3} or C_{2}H_{6}.[48]

  [48] The homologues of ethylene, C_{_n_}H_{2_n_}, are also capable
       of direct combination with halogens, &c., but with various
       degrees of facility. The composition of these homologues can be
       expressed thus: (CH_{3})__x_(CH_{2})_{_y_}(CH)_{_z_}C_{_r_},
       where the sum of _x_ + _z_ is always an even number, and
       the sum of _x_ + _z_ + _r_ is equal to half the sum of
       3_x_ + _z_, whence _z_ + 2_r_ = _x_; by this means the possible
       isomerides are determined. For example, for butylenes,
       C_{4}H_{8}, (CH_{3})_{2}(CH)_{2}, (CH_{3})_{2}(CH_{2})C,
       (CH_{2})(CH_{2})_{2}CH, and (CH_{2})_{4} are possible.

_Acetylene_, C_{2}H_{2} = CHCH, is a gas; it was first prepared by
Berthelot (1857). It has a very pungent smell, is characterised by
its great stability under the action of heat, and is obtained as the
only product of the direct combination of carbon with hydrogen when a
luminous arc (voltaic) is formed between carbon electrodes. This arc
contains particles of carbon passing from one pole to the other. If the
carbons be surrounded with an atmosphere of hydrogen, the carbon in part
combines with the hydrogen, forming C_{2}H_{2}.[48 bis] Acetylene may be
formed from olefiant gas if two atoms of hydrogen be taken from it. This
may be effected in the following way: the olefiant gas is first made to
combine with bromine, giving C_{2}H_{4}Br_{2}; from this the hydrobromic
acid is removed by means of an alcoholic solution of caustic potash,
leaving the volatile product C_{2}H_{3}Br; and from this yet another
part of hydrobromic acid is withdrawn by passing it through anhydrous
alcohol in which metallic sodium has been dissolved, or by heating
it with a strong alcoholic solution of caustic potash. Under these
circumstances (Berthelot, Sawitsch, Miasnikoff) the alkali takes up the
hydrobromic acid from C_{_n_}H_{2_n_-1}Br, forming C_{_n_}H_{2_n_-2}.

  [48 bis] _See_ also method of preparing C_{2}H_{2} in Note 12 bis.

Acetylene is also produced in all those cases where organic substances
are decomposed by the action of a high temperature--for example, by dry
distillation. On this account a certain quantity is always found in coal
gas, and gives to it, at all events in part, its peculiar smell, but
the quantity of acetylene in coal gas is very small. If the vapour of
alcohol be passed through a heated tube a certain quantity of acetylene
is formed. It is also produced by the imperfect combustion of olefiant
and marsh gas--for example, if the flame of coal gas has not free access
to air.[49] The inner part of every flame contains gases in imperfect
combustion, and in them some amount of acetylene.

  [49] This is easily accomplished with those gas burners which are used
       in laboratories and mentioned in the Introduction. In these
       burners the gas is first mixed with air in a long tube, above
       which it is kindled. But if it be lighted inside the pipe it
       does not burn completely, but forms acetylene, on account of
       the cooling effect of the walls of the metallic tube; this is
       detected by the smell, and may be shown by passing the issuing
       gas (by aid of an aspirator) into an ammoniacal solution of
       cuprous chloride.

Acetylene, being further removed than ethylene from the limit
C_{_n_}H_{2_n_+2} of hydrocarbon compounds, has a still greater
faculty of combination than is shown by olefiant gas, and therefore can
be more readily separated from any mixture containing it. Actually,
acetylene not only combines with one and two molecules of I_{2}, HI,
H_{2}SO_{4}, Cl_{2}, Br_{2}, &c.... (many other unsaturated hydrocarbons
combine with them), but also with cuprous chloride, CuCl, forming a red
precipitate. If a gaseous mixture containing acetylene be passed through
an ammoniacal solution of cuprous chloride (or silver nitrate), the
other gases do not combine, but the acetylene gives a red precipitate
(or grey with silver), which detonates when struck with a hammer. This
red precipitate gives off acetylene under the action of acids. In this
manner pure acetylene may be obtained. Acetylene and its homologues also
readily react with corrosive sublimate, HgCl_{2} (Koucheroff, Favorsky).
Acetylene burns with a very brilliant flame, which is accounted for by
the comparatively large amount of carbon it contains.[50]

  [50] Amongst the homologues of acetylene C_{_n_}H_{2_n_-2}, the lowest
       is C_{3}H_{4}; allylene, CH_{3}CCH, and allene, CH_{2}CCH_{2},
       are known, but the closed structure, CH_{2}(CH)_{2}, is little
       investigated.

The formation and existence in nature of large masses of petroleum or a
mixture of liquid hydrocarbons, principally of the series C_{_n_}H_{2_n_
+2} and C_{_n_}H_{2_n_} is in many respects remarkable.[51] In some
mountainous districts--as, for instance, by the <DW72>s of the Caucasian
chain, on inclines lying in a direction parallel to the range--an oily
liquid issues from the earth together with salt water and hot gases
(methane and others); it has a tarry smell and dark brown colour,
and is lighter than water. This liquid is called naphtha or rock oil
(petroleum) and is obtained in large quantities by sinking wells and
deep bore-holes in those places where traces of naphtha are observed,
the naphtha being sometimes thrown up from the wells in fountains of
considerable height.[52] The evolution of naphtha is always accompanied
by salt water and marsh gas. Naphtha has from ancient times been
worked in Russia in the Apsheron peninsula near Baku, and is also
now worked in Burmah (India), in Galicia near the Carpathians, and
in America, especially in Pennsylvania and Canada, &c. Naphtha does
not consist of one definite hydrocarbon, but of a mixture of several,
and its density, external appearance, and other qualities vary with
the amount of the different hydrocarbons of which it is composed.
The light kinds of naphtha have a specific gravity about 0·8 and the
heavy kinds up to 0·98. The former are very mobile liquids, and more
volatile; the latter contain less of the volatile hydrocarbons and are
less mobile. When the light kinds of naphtha are distilled, the boiling
point taken in the vapours constantly changes, beginning at 0° and
going up to above 350°. That which passes over first is a very mobile,
colourless ethereal liquid (forming gazolene, ligroin, benzoline,
&c.), from which the hydrocarbons whose boiling points start from 0°
may be extracted--namely, the hydrocarbons C_{4}H_{10}, C_{5}H_{12}
(which boils at 30°), C_{6}H_{14} (boils at 62°), C_{7}H_{16} (boils
about 90°), &c. Those fractions of the naphtha distillate which boil
above 130°, and contain hydrocarbons with C_{9}, C_{10}, C_{11}, &c.,
enter into the composition of the oily substance, universally used
for lighting, called kerosene or photogen or photonaphthalene, and by
other names. The specific gravity of kerosene is from 0·78 to 0·84, and
it smells like naphtha. Those products of the distillation of naphtha
which pass off below 130° and have a specific gravity below 0·75,
enter into the composition of light petroleum (benzoline, ligroin,
petroleum spirit, &c.); which is used as a solvent for india-rubber,
for removing grease spots, &c. Those portions of naphtha (which can
only be distilled without change by means of superheated steam,
otherwise they are largely decomposed) which boil above 275° and up to
300° and have a specific gravity higher than 0·85, form an excellent
oil,[53] safe as regards inflammability (which is very important
as diminishing the risks of fire), and may be used in lamps as an
effective substitute for kerosene.[54] Those portions of naphtha which
pass over at a still higher temperature and have a higher specific
gravity than 0·9, which are found in abundance (about 30 p.c.) in the
Baku naphtha, make excellent lubricating or machine oils. Naphtha has
many important applications, and the naphtha industry is now of great
commercial importance, especially as naphtha and its refuse may be
used as fuel.[55] Whether naphtha was formed from organic matter is
very doubtful, as it is found in the most ancient Silurian strata
which correspond with epochs of the earth's existence when there was
little organic matter; it could not penetrate from the higher to the
lower (more ancient) strata as it floats on water (and water penetrates
through all strata). It therefore tends to rise to the surface of the
earth, and it is always found in highlands parallel to the direction of
the mountains.[56] Much more probably its formation may be attributed
to the action of water penetrating through the crevasses formed on
the mountain <DW72>s and reaching to the heart of the earth, to that
kernel of heated metallic matter which must be accepted as existing in
the interior of the earth. And as meteoric iron often contains carbon
(like cast iron), so, accepting the existence of such carburetted
iron at unattainable depths in the interior of the earth, it may be
supposed that naphtha was produced by the action of water penetrating
through the crevices of the strata during the upheaval of mountain
chains,[57] because water with iron carbide ought to give iron oxide
and hydrocarbons.[58] Direct experiment proves that the so-called
_spiegeleisen_ (manganiferous iron, rich in chemically combined
carbon) when treated with acids gives liquid hydrocarbons[59] which in
composition, appearance, and properties are completely identical with
naphtha.[60]

  [51] The saturated hydrocarbons predominate in American petroleum,
       especially in its more volatile parts; in Baku naphtha the
       hydrocarbons of the composition C_{_n_}H_{2_n_} form the
       main part (Lisenko, Markovnikoff, Beilstein) but doubtless
       (Mendeléeff) it also contains saturated ones, C_{_n_}H_{2_n_+2}.
       The structure of the naphtha hydrocarbons is only known for
       the lower homologues, but doubtless the distinction between the
       hydrocarbons of the Pennsylvanian and Baku naphthas, boiling at
       the same temperature (after the requisite refining by repeated
       fractional distillation, which can be very conveniently done
       by means of steam rectification--that is, by passing the steam
       through the dense mass), depends not only on the predominance
       of saturated hydrocarbons in the former, and naphthenes,
       C_{_n_}H_{2_n_}, in the latter, but also on the diversity of
       composition and structure of the corresponding portions of the
       distillation. The products of the Baku naphtha are richer in
       carbon (therefore in a suitably constructed lamp they ought to
       give a brighter light), they are of greater specific gravity, and
       have greater internal friction (and are therefore more suitable
       for lubricating machinery) than the American products collected
       at the same temperature.

  [52] The formation of naphtha fountains (which burst forth after the
       higher clay strata covering the layers of sands impregnated
       with naphtha have been bored through) is without doubt caused
       by the pressure or tension of the combustible hydrocarbon
       gases which accompany the naphtha, and are soluble in it under
       pressure. Sometimes these naphtha fountains reach a height of 100
       metres--for instance, the fountain of 1887 near Baku. Naphtha
       fountains generally act periodically and their force diminishes
       with the lapse of time, which might be expected, because the
       gases which cause the fountains find an outlet, as the naphtha
       issuing from the bore-hole carries away the sand which was
       partially choking it up.

  [53] This is a so-called intermediate oil (between kerosene and
       lubricating oils), solar oil, or pyronaphtha. Lamps are already
       being manufactured for burning it but still require improvement.
       Above all, however, it requires a more extended market, and
       this at present is wanting, owing to the two following reasons:
       (1) Those products of the American petroleum which are the
       most widely spread and almost universally consumed contain but
       little of this intermediate oil, and what there is is divided
       between the kerosene and the lubricating oils; (2) the Baku
       naphtha, which is capable of yielding a great deal (up to 30
       p.c.) of intermediate oil, is produced in enormous quantities,
       about 300 million poods, but has no regular markets abroad,
       and for the consumption in Russia (about 25 million poods of
       kerosene per annum) and for the limited export (60 million
       poods per annum) into Western Europe (by the Trans-Caucasian
       Railway) those volatile and more dangerous parts of the naphtha
       which enter into the composition of the American petroleum are
       sufficient, although Baku naphtha yields about 25 p.c. of such
       kerosene. For this reason pyronaphtha is not manufactured in
       sufficient quantities, and the whole world is consuming the
       unsafe kerosene. When a pipe line has been laid from Baku to the
       Black Sea (in America there are many which carry the raw naphtha
       to the sea-shore, where it is made into kerosene and other
       products) then the whole mass of the Baku naphtha will furnish
       safe illuminating oils, which without doubt will find an immense
       application. A mixture of the intermediate oil with kerosene
       or Baku oil (specific gravity 0·84 to 0·85) may be considered
       (on removing the benzoline) to be the best illuminating oil,
       because it is safe (flashing point from 40° to 60°), cheaper
       (Baku naphtha gives as much as 60 p.c. of Baku oil), and burns
       perfectly well in lamps differing but little from those made for
       burning American kerosene (unsafe, flashing point 20° to 30°).

  [54] The substitution of Baku pyronaphtha, or intermediate oil, or Baku
       oil (_see_ Note 53), would not only be a great advantage as
       regards safety from fire, but would also be highly economical.
       A ton (62 poods) of American crude petroleum costs at the coast
       considerably more than 24_s._ (12 roubles), and yields two-thirds
       of a ton of kerosene suitable for ordinary lamps. A ton of raw
       naphtha in Baku costs less than 4_s._ (1 rouble 80 copecks), and
       with a pipe line to the shore of the Black Sea would not cost
       more than 8 roubles, or 16_s._ Moreover, a ton of Baku naphtha
       will yield as much as two-thirds of a ton of kerosene, Baku oil,
       and pyronaphtha suitable for illuminating purposes.

  [55] Naphtha has been applied for heating purposes on a large scale
       in Russia, not only on account of the low cost of naphtha itself
       and of the residue from the preparation of kerosene, but also
       because the products of all the Baku naphtha do not find an
       outlet for general consumption. Naphtha itself and its various
       residues form excellent fuel, burning without smoke and giving
       a high temperature (steel and iron may be easily melted in the
       flame). A hundred poods of good coal (for instance, Don coal)
       used as fuel for heating boilers are equivalent to 36 cubic feet
       (about 250 poods) of dry wood, while only 70 poods of naphtha
       will be required; and moreover there is no need for stoking, as
       the liquid can be readily and evenly supplied in the required
       quantity. The economic and other questions relating to American
       and Baku petroleums have been discussed more in detail in some
       separate works of mine (D. Mendeléeff): (1) 'The Naphtha Industry
       of Pennsylvania and the Caucasus,' 1870; (2) 'Where to Build
       Naphtha Works,' 1880; (3) 'On the Naphtha Question,' 1883; (4)
       'The Baku Naphtha Question,' 1886; (5) the article on the naphtha
       industry in the account of the Russian industries printed for the
       Chicago Exhibition.

  [56] As during the process of the dry distillation of wood, sea-weed,
       and similar vegetable _débris_, and also when fats are decomposed
       by the action of heat (in closed vessels), hydrocarbons similar
       to those of naphtha are formed, it was natural that this fact
       should have been turned to account to explain the formation
       of the latter. But the hypothesis of the formation of naphtha
       from vegetable _débris_ inevitably assumes coal to be the chief
       element of decomposition, and naphtha is met with in Pennsylvania
       and Canada, in the Silurian and Devonian strata, which do not
       contain coal, and correspond to an epoch not abounding in organic
       matter. Coal was formed from the vegetable _débris_ of the
       Carboniferous, Jurassic, and other recent strata, but judging
       more from its composition and structure, it has been subjected
       to the same kind of decomposition as peat; nor could liquid
       hydrocarbons have been thus formed to such an extent as we see
       in naphtha. If we ascribe the derivation of naphtha to the
       decomposition of fat (adipose, animal fat) we encounter three
       almost insuperable difficulties: (1) Animal remains would furnish
       a great deal of nitrogenous matter, whilst there is but very
       little in naphtha; (2) the enormous quantity of naphtha already
       discovered as compared with the insignificant amount of fat in
       the animal carcase; (3) the sources of naphtha always running
       parallel to mountain chains is completely inexplicable. Being
       struck with this last-mentioned circumstance in Pennsylvania,
       and finding that the sources in the Caucasus surround the whole
       Caucasian range (Baku, Tiflis, Gouria, Kouban, Tamman, Groznoe,
       Dagestan), I developed in 1876 the hypothesis of the mineral
       origin of naphtha expounded further on.

  [57] During the upheaval of mountain ranges crevasses would be formed
       at the peaks with openings upwards, and at the foot of the
       mountains with openings downwards. These cracks in course of time
       fill up, but the younger the mountains the fresher the cracks
       (the Alleghany mountains are, without doubt, more ancient than
       the Caucasian, which were formed during the tertiary epoch);
       through them water must gain access deep into the recesses of the
       earth to an extent that could not occur on the level (on plains).
       The situation of naphtha at the foot of mountain chains is the
       principal argument in my hypothesis.

       Another fundamental reason is the consideration of the mean
       density of the earth. Cavendish, Airy, Cornu, Boys, and many
       others who have investigated the subject by various methods,
       found that, taking water = 1, the mean density of the earth
       is nearly 5·5. As at the surface water and all rocks (sand,
       clay, limestone, granite, &c.) have a density less than 3, it
       is evident (as solid substances are but slightly compressible
       even under the greatest pressure) that inside the earth there
       are substances of a greater density--indeed, not less than 7
       or 8. What conclusion, then, can be arrived at? Anything heavy
       contained in the bosom of the earth must be distributed not
       only on its surface, but throughout the whole solar system, for
       everything tends to show that the sun and planets are formed from
       the same material, and according to the hypothesis of Laplace
       and Kant it is most probable, and indeed must necessarily be
       held, that the earth and planets are but fragments of the solar
       atmosphere, which have had time to cool considerably and become
       masses semi-liquid inside and solid outside, forming both planets
       and satellites. The sun amongst other heavy elements contains
       a great deal of iron, as shown by spectrum analysis. There is
       also much of it in an oxidised condition on the surface of the
       earth. Meteoric stones, carried as fragmentary planets in the
       solar system and sometimes falling upon the earth, consisting of
       siliceous rocks similar to terrestrial ones, often contain either
       dense masses of iron (for example, the Pallosovo iron preserved
       in the St. Petersburg Academy of Sciences) or granular masses
       (for instance, the Okhansk meteorite of 1886). It is therefore
       possible that the interior of the earth contains much iron in a
       metallic state. This might be anticipated from the hypothesis of
       Laplace, for the iron must have been compressed into a liquid
       at that period when the other component parts of the earth
       were still strongly heated, and oxides of iron could not then
       have been formed. The iron was covered with slags (mixtures of
       silicates like glass fused with rocky matter) which did not allow
       it to burn at the expense of the oxygen of the atmosphere or
       of water, just at that time when the temperature of the earth
       was very high. Carbon was in the same state; its oxides were
       also capable of dissociation (Deville); it is also but slightly
       volatile, and has an affinity for iron, and iron carbide is found
       in meteoric stones (as well as carbon and even the diamond). Thus
       the supposition of the existence of iron carbides in the interior
       of the earth was derived by me from many indications, which
       are to some extent confirmed by the fact that granular pieces
       of iron have been found in some basalts (ancient lava) as well
       as in meteoric stones. The occurrence of iron in contact with
       carbon during the formation of the earth is all the more probable
       because those elements predominate in nature which have small
       atomic weights, and among them the most widely diffused, the most
       difficultly fusible, and therefore the most easily condensed
       (Chapter XV.) are carbon and iron. They passed into the liquid
       state when all compounds were at a temperature of dissociation.

  [58] The following is the typical equation for this formation:

       3Fe_{_m_}C_{_n_} + 4_{_m_}H_{2}O = _m_Fe_{3}O_{4} (magnetic
       oxide) + C_{5_n_}H_{8_m_} (_see_ Chapter XVII., Note 38).

  [59] Cloez investigated the hydrocarbons formed when cast-iron is
       dissolved in hydrochloric acid, and found C_{_n_}H_{2_n_} and
       others. I treated crystalline manganiferous cast-iron with the
       same acid, and obtained a liquid mixture of hydrocarbons exactly
       similar to natural naphtha in taste, smell, and reaction.

  [60] Probably naphtha was produced during the upheaval of all mountain
       chains, but only in some cases were the conditions favourable
       to its being preserved underground. The water penetrating below
       formed there a mixture of naphtha and watery vapours, and this
       mixture issued through fissures to the cold parts of the earth's
       crust. The naphtha vapours, on condensing, formed naphtha,
       which, if there were no obstacles, appeared on the surface
       of land and water. Here part of it soaked through formations
       (possibly the bituminous slates, schists, dolomites, &c., were
       thus formed), another part was carried away on the water,
       became oxidised, evaporated, and was driven to the shores (the
       Caucasian naphtha probably in this way, during the existence of
       the Aralo-Caspian sea, was carried as far as the Sisran banks
       of the Volga, where many strata are impregnated with naphtha
       and products of its oxidation resembling asphalt and pitch); a
       great part of it was burnt in one way or another--that is, gave
       carbonic anhydride and water. If the mixture of vapours, water,
       and naphtha formed inside the earth had no free outlet to the
       surface, it nevertheless would find its way through fissures to
       the superior and colder strata, and there become condensed. Some
       of the formations (clays) which do not absorb naphtha were only
       washed away by the warm water, and formed mud, which we also now
       observe issuing from the earth in the form of mud volcanoes. The
       neighbourhood of Baku and the whole of the Caucasus near the
       naphtha districts are full of such volcanoes, which from time to
       time are in a state of eruption. In old naphtha beds (such as
       the Pennsylvanian) even these blow-holes are closed, and the mud
       volcanoes have had time to be washed away. The naphtha and the
       gaseous hydrocarbons formed with it under the pressure of the
       overlying earth and water impregnated the layers of sand, which
       are capable of absorbing a great quantity of such liquid, and
       if above this there were strata impermeable to naphtha (dense,
       clayey, damp strata) the naphtha would accumulate in them. It is
       thus preserved from remote geological periods up to the present
       day, compressed and dissolved under the pressure of the gases
       which burst out in places forming naphtha fountains. If this
       be granted, it may be thought that in the comparatively new
       (geologically speaking) mountain chains, such as the Caucasian,
       naphtha is even now being formed. Such a supposition may explain
       the remarkable fact that, in Pennsylvania, localities where
       naphtha had been rapidly worked for five years have become
       exhausted, and it becomes necessary to constantly have recourse
       to sinking new wells in fresh places. Thus, from the year 1859,
       the workings were gradually transferred along a line running
       parallel to the Alleghany mountains for a distance of more
       than 200 miles, whilst in Baku the industry dates from time
       immemorial (the Persians worked near the village of Ballaghana)
       and up to the present time keeps to one and the same place.
       The amounts of the Pennsylvanian and Baku annual outputs are
       at present equal--namely, about 250 million poods (4 million
       tons). It may be that the Baku beds, as being of more recent
       geological formation, are not so exhausted by nature as those of
       Pennsylvania, and perhaps in the neighbourhood of Baku naphtha
       is still being formed, which is partially indicated by the
       continued activity of the mud volcanoes. As many varieties of
       naphtha contain in solution solid slightly volatile hydrocarbons
       like paraffin and mineral wax, the production of ozocerite,
       or mountain wax, is accounted for in conjunction with the
       formation of naphtha. Ozocerite is found in Galicia, also in
       the neighbourhood of Novorossisk, in the Caucasus, and on the
       islands of the Caspian Sea (particularly in the Chileken and Holy
       Islands); it is met with in large masses, and is used for the
       production of paraffin and _ceresene_, for the manufacture of
       candles, and similar purposes.

       As the naphtha treasures of the Caucasus have hardly been
       exploited (near Baku and near Kouban and Grosnyi), and as naphtha
       finds numerous uses, the subject presents most interesting
       features to chemists and geologists, and is worthy of the close
       attention of practical men.




                               CHAPTER IX

              COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN


[Illustration: FIG. 61.--Dumas and Stas' apparatus for determining
the composition of carbonic anhydride. Carbon, graphite, or a diamond
is placed in the tube E in the furnace, and heated in a stream of
oxygen displaced from the bottle by water flowing from A. The oxygen
is purified from carbonic anhydride and water in the tubes B, C, D.
Carbonic anhydride, together with a certain amount of carbon monoxide,
is formed in E. The latter is converted into carbonic acid by passing
the products of combustion through a tube F, containing cupric oxide
heated in a furnace. The cupric oxide oxidises this CO into CO_{2},
forming metallic copper. The potash bulbs H and tubes I, J, K retain
the carbonic anhydride. Thus, knowing the weight of carbon taken and
the weight of the resultant carbonic anhydride (by weighing H, I, J, K
before and after the experiment), the composition of carbonic anhydride
and the equivalent of carbon may be determined.]

Carbonic anhydride (or carbonic acid or carbon dioxide, CO_{2}) was
the first of all gases distinguished from atmospheric air. Paracelsus
and Van Helmont, in the sixteenth century, knew that on heating
limestone a particular gas separated, which is also formed during the
alcoholic fermentation of saccharine solutions (for instance, in the
manufacture of wine); they knew that it was identical with the gas
which is produced by the combustion of charcoal, and that in some
cases it is found in nature. In course of time it was found that this
gas is absorbed by alkali, forming a salt which, under the action of
acid, again yields this same gas. Priestley found that this gas exists
in air, and Lavoisier determined its formation during respiration,
combustion, putrefaction, and during the reduction of the oxides of
metals by charcoal; he determined its composition, and showed that
it only contains oxygen and carbon. Berzelius, Dumas with Stas, and
Roscoe, determined its composition, showing that it contains twelve
parts of carbon to thirty-two of oxygen. The composition by volume
of this gas is determined from the fact that during the combustion
of charcoal in oxygen, the volume remains unchanged; that is to say,
_carbonic anhydride occupies the same volume as the oxygen which it
contains_--that is, the atoms of the carbon are, so to speak, squeezed
in between the atoms of the oxygen. O_{2} occupies two volumes and is
a molecule of ordinary oxygen; CO_{2} likewise occupies two volumes,
and expresses the composition and molecular weight of the gas. Carbonic
anhydride exists _in nature_, both in a free state and in the most
varied compounds. In a free state it is always contained (Chapter V.) in
the air, and in solution is in all kinds of water. It is evolved from
volcanoes, from mountain fissures, and in some caves. The well-known
Dog grotto, near Agnano on the bay of Baiæ, near Naples, furnishes the
best known example of such an evolution. Similar sources of carbonic
anhydride are also found in other places. In France, for instance, there
is a well-known poisonous fountain in Auvergne. It is a round hole,
surrounded with luxurious vegetation and constantly evolving carbonic
anhydride. In the woods surrounding the Lacher See near the Rhine, in
the neighbourhood of extinct volcanoes, there is a depression constantly
filled with this same gas. The insects which fly to this place perish,
animals being unable to breathe this gas. The birds chasing the insects
also die, and this is turned to profit by the local peasantry. Many
mineral springs carry into the air enormous quantities of this gas.
Vichy in France, Sprüdel in Germany, and Narzan in Russia (in Kislovodsk
near Piatigorsk) are known for their carbonated gaseous waters. Much
of this gas is also evolved in mines, cellars, diggings, and wells.
People descending into such places are suffocated. The combustion,
putrefaction, and fermentation of organic substances give rise to
the formation of carbonic anhydride. It is also introduced into the
atmosphere during the respiration of animals at all times and during
the respiration of plants in darkness and also during their growth.
Very simple experiments prove the formation of carbonic anhydride
under these circumstances; thus, for example, if the air expelled from
the lungs be passed through a glass tube into a transparent solution
of lime (or baryta) in water a white precipitate will soon be formed
consisting of an insoluble compound of lime and carbonic anhydride.
By allowing the seeds of plants to grow under a bell jar, or in a
closed vessel, the formation of carbonic anhydride may be similarly
confirmed. By confining an animal, a mouse, for instance, under a bell
jar, the quantity of carbonic acid which it evolves may be exactly
determined, and it will he found to be many grams per day for a mouse.
Such experiments on the respiration of animals have been also made with
great exactitude with large animals, such as men, bulls, sheep, &c. By
means of enormous hermetically closed bell receivers and the analysis
of the gases evolved during respiration it was found that a man expels
about 900 grams (more than two pounds) of carbonic anhydride per diem,
and absorbs during this time 700 grams of oxygen.[1] It must be remarked
that the carbonic anhydride of the air constitutes the fundamental
food of plants (Chapters III., V., and VIII.) Carbonic anhydride in a
state of combination with a variety of other substances is perhaps even
more widely distributed in nature than in a free state. Some of these
substances are very stable and form a large portion of the earth's
crust. For instance, limestones, calcium carbonate, CaCO_{3}, were
formed as precipitates in the seas existing previously on the earth;
this is proved by their stratified structure and the number of remains
of sea animals which they frequently contain. Chalk, lithographic stone,
limestone, marls (a mixture of limestone and clay), and many other rocks
are examples of such sedimentary formations. Carbonates with various
other bases--such as, for instance, magnesia, ferrous oxide, zinc oxide,
&c.--are often found in nature. The shells of molluscs also have the
composition CaCO_{3} and many limestones were exclusively formed from
the shells of minute organisms. As carbonic anhydride (together with
water) is produced during the combustion of all organic compounds in
a stream of oxygen or by heating them with substances which readily
part with their oxygen--for instance, with copper oxide--this method
is employed for estimating the amount of carbon in organic compounds,
more especially as the CO_{2} can be easily collected and the amount of
carbon calculated from its weight. For this purpose a hard glass tube,
closed at one end, is filled with a mixture of the organic substance
(about 0·2 gram) and copper oxide. The open end of the tube is fitted
with a cork and tube containing calcium chloride for absorbing the water
formed by the oxidation of the substance. This tube is hermetically
connected (by a caoutchouc tube) with potash bulbs or other weighing
apparatus (Chapter V.) containing alkali destined to absorb the carbonic
anhydride. The increase in weight of this apparatus shows the amounts of
carbonic anhydride formed during the combustion of the given substance,
and the quantity of carbon may be determined from this, because three
parts of carbon give eleven parts of carbonic anhydride.

[Illustration: FIG.--62. Apparatus for the combustion of organic
substances by igniting them with oxide of copper.]

  [1] The quantity of carbonic acid gas exhaled by a man during the
      twenty-four hours is not evenly produced; during the night more
      oxygen is taken in than during the day (by night, in twelve hours,
      about 450 grams), and more carbonic anhydride is separated by day
      than during night-time and repose; thus, of the 900 grams produced
      during the twenty-four hours about 375 are given out during the
      night and 525 by day. This depends on the formation of carbonic
      anhydride during the work performed by the man in the day. Every
      movement is the result of some change of matter, for force cannot
      be self-created (in accordance with the law of the conservation of
      energy). Proportionally to the amount of carbon consumed an amount
      of energy is stored up in the organism and is consumed in the
      various movements performed by animals. This is proved by the fact
      that during work a man exhales 525 grams of carbonic anhydride in
      twelve hours instead of 375, absorbing the same amount of oxygen
      as before. After a working day a man exhales by night almost the
      same amount of carbonic anhydride as after a day of rest, so that
      during a total twenty-four hours a man exhales about 900 grams
      of carbonic anhydride and absorbs about 980 grams of oxygen.
      Therefore during work the change of matter increases. The carbon
      expended on the work is obtained from the food; on this account
      the food of animals ought certainly to contain carbonaceous
      substances capable of dissolving under the action of the digestive
      fluids, and of passing into the blood, or, in other words, capable
      of being digested. Such food for man and all other animals is
      formed of vegetable matter, or of parts of other animals. The
      latter in every case obtain their carbonaceous matter from plants,
      in which it is formed by the separation of the carbon from the
      carbonic anhydride taken up during the day by the respiration of
      the plants. The volume of the oxygen exhaled by plants is almost
      equal to the volume of the carbonic anhydride absorbed; that is to
      say, nearly all the oxygen entering into the plant in the form of
      carbonic anhydride is liberated in a free state, whilst the carbon
      from the carbonic anhydride remains in the plant. At the same time
      the plant absorbs moisture by its leaves and roots. By a process
      which is unknown to us, this absorbed moisture and the carbon
      obtained from the carbonic anhydride enter into the composition of
      the plants in the form of so-called carbohydrates, composing the
      greater part of the vegetable tissues, starch and cellulose of the
      composition C_{6}H_{10}O_{5} being representatives of them. They
      may be considered like all carbohydrates as compounds of carbon
      and water, 6C + 5H_{2}O. In this way a _circulation_ of the carbon
      goes on in nature by means of vegetable and animal organisms, in
      which changes the principal factor is the carbonic anhydride of
      the air.

_For the preparation of carbonic anhydride_ in laboratories and often
in manufactories, various kinds of calcium carbonate are used, being
treated with some acid; it is, however, most usual to employ the
so-called muriatic acid--that is, an aqueous solution of hydrochloric
acid, HCl--because, in the first place, the substance formed, calcium
chloride. CaCl_{2}, is soluble in water and does not hinder the further
action of the acid on the calcium carbonate, and secondly because, as
we shall see further on, muriatic acid is a common product of chemical
works and one of the cheapest. For calcium carbonate, either limestone,
chalk, or marble is used.[2]

              CaCO_{3} + 2HCl = CaCl_{2} + H_{2}O + CO_{2}.

The nature of the reaction in this case is the same as in the
decomposition of nitre by sulphuric acid; only in the latter case a
hydrate is formed, and in the former an anhydride of the acid, because
the hydrate, carbonic acid, H_{2}CO_{3}, is unstable and as soon as it
separates decomposes into water and its own anhydride. It is evident
from the explanation of the cause of the action of sulphuric acid
on nitre that not every acid can be employed for obtaining carbonic
anhydride; namely, those will not set it free which chemically are
but slightly energetic, or those which are insoluble in water, or are
themselves as volatile as carbonic anhydride.[3] But as many acids are
soluble in water and are less volatile than carbonic anhydride, the
latter is evolved by the action of most acids on its salts, and this
reaction takes place at ordinary temperatures.[4]

  [2] Other acids may be used instead of hydrochloric; for instance,
      acetic, or even sulphuric, although this latter is not suitable,
      because it forms as a product insoluble calcium sulphate (gypsum)
      which surrounds the untouched calcium carbonate, and thus prevents
      a further evolution of gas. But if porous limestone--for instance,
      chalk--be treated with sulphuric acid diluted with an equal volume
      of water, the liquid is absorbed and acting on the mass of the
      salt, the evolution of carbonic anhydride continues evenly for a
      long time. Instead of calcium carbonate other carbonates may of
      course be used; for instance, washing-soda, Na_{2}CO_{3}, which
      is often chosen when it is required to produce a rapid stream of
      carbonic anhydride (for example, for liquefying it). But natural
      crystalline magnesium carbonate and similar salts are with
      difficulty decomposed by hydrochloric and sulphuric acids. When
      for manufacturing purposes--for instance, in precipitating lime in
      sugar-works--a large quantity of carbonic acid gas is required,
      it is generally obtained by burning charcoal, and the products of
      combustion, rich in carbonic anhydride, are pumped into the liquid
      containing the lime, and the carbonic anhydride is thus absorbed.
      Another method is also practised, which consists in using the
      carbonic anhydride separated during fermentation, or that
      evolved from limekilns. During the fermentation of sweet-wort,
      grape-juice, and other similar saccharine solutions, the glucose
      C_{6}H_{12}O_{6} changes under the influence of the yeast
      organism, forming alcohol (2C_{2}H_{6}O), and carbonic anhydride
      (2CO_{2}) which separates in the form of gas; if the fermentation
      proceeds in closed bottles sparkling wine is obtained. When
      carbonic acid gas is prepared for saturating water and other
      beverages it is necessary to use it in a pure state. Whilst in
      the state in which it is evolved from ordinary limestones by the
      aid of acids it contains, besides a certain quantity of acid,
      the organic matters of the limestone; in order to diminish the
      quantity of these substances the densest kinds of dolomites are
      used, which contain less organic matter, and the gas formed is
      passed through various washing apparatus, and then through a
      solution of potassium permanganate, which absorbs organic matter
      and does not take up carbonic anhydride.

  [3] Hypochlorous acid, HClO, and its anhydride, Cl_{2}O, do not displace
      carbonic acid, and hydrogen sulphide has the same relation to
      carbonic acid as nitric acid to hydrochloric--an excess of either
      one displaces the other.

  [4] Thus, in preparing the ordinary effervescing powders, sodium
      bicarbonate (or acid carbonate of soda) is used, and mixed
      with powdered citric or tartaric acid. In a dry state these
      powders do not evolve carbonic anhydride, but when mixed with
      water the evolution takes place briskly, which is due to the
      substances passing into solution. The salts of carbonic acid
      may be recognised from the fact that they evolve carbonic acid
      with a hissing noise when treated with acids. If vinegar, which
      contains acetic acid, be poured upon limestone, marble, malachite
      (containing copper carbonate), &c., carbonic anhydride is evolved
      with a hissing noise. It is noteworthy that neither hydrochloric
      acid, nor even sulphuric acid nor acetic acid, acts on limestone
      except in presence of water. We shall refer to this later on.

For the preparation of carbonic anhydride in laboratories, marble is
generally used. It is placed in a Woulfe's bottle and treated with
hydrochloric acid in an apparatus similar to the one used for the
production of hydrogen. The gas evolved carries away through the tube
part of the volatile hydrochloric acid, and it is therefore necessary to
wash the gas by passing it through another Woulfe's bottle containing
water. If it be necessary to obtain dry carbonic anhydride, it must be
passed through chloride of calcium.[5]

  [5] The direct observations made (1876) by Messrs. Bogouski and
      Kayander lead to the conclusion that the quantity of carbonic
      anhydride evolved by the action of acids on marble (as homogeneous
      as possible) is directly proportional to the time of action, the
      extent of surface, and the degree of concentration of the acid,
      and inversely proportional to the molecular weight of the acid.
      If the surface of a piece of Carrara marble be equal to one
      decimetre, the time of action one minute, and one cubic decimetre
      or litre contains one gram of hydrochloric acid, then about 0·02
      gram of carbonic anhydride will be evolved. If the litre contains
      _n_ grams of hydrochloric acid, then by experiment the amount
      will be _n_ × 0·02 of carbonic anhydride. Therefore, if the
      litre contains 36·5 (= HCl) grams, about 0·73 gram of carbonic
      anhydride (about half a litre) would he evolved per minute. If
      nitric acid or hydrobromic acid be used instead of hydrochloric,
      then, with a combining proportion of the acid, the same quantity
      of carbonic anhydride will be evolved; thus, if the litre contains
      63 (= HNO_{3}) grams of nitric acid, or 81 (= HBr) grams of
      hydrobromic acid, the quantity of carbonic anhydride evolved will
      still be 0·73 gram. Spring, in 1890, made a series of similar
      determinations.

Carbonic anhydride may also be prepared by heating many of the salts of
carbonic acid; for instance, by heating magnesium carbonate, MgCO_{3}
(_e.g._, in the form of dolomite), the separation is easily effected,
particularly in the presence of the vapours of water. The acid salts
of carbonic acid (for instance, NaHCO_{3}, see further on) readily and
abundantly give carbonic anhydride when heated.

Carbonic anhydride is colourless, has a slight smell and a faint acid
taste; its density in a gaseous state is twenty-two times as great
as that of hydrogen, because its molecular weight is forty-four.[6]
It is an example of those gaseous substances which have been long
ago transformed into all the three states. In order to obtain liquid
carbonic anhydride, the gas must be submitted to a pressure of
thirty-six atmospheres at 0°.[7] Its absolute boiling point = +32°.[8]
Liquid carbonic anhydride is colourless, does not mix with water, but
is soluble in alcohol, ether, and oils; at 0° its specific gravity is
0·83.[8 bis] The boiling point of this liquid lies at -80°--that is to
say, the pressure of carbonic acid gas at that temperature does not
exceed that of the atmosphere. At the ordinary temperature the liquid
remains as such for some time under ordinary pressure, on account of
its requiring a considerable amount of heat for its evaporation. If
the evaporation takes place rapidly, especially if the liquid issues
in a stream, such a decrease of temperature occurs that a part of the
carbonic anhydride is transformed into a solid snowy mass. Water,
mercury, and many other liquids freeze on coming into contact with
snow-like carbonic anhydride.[9] In this form carbonic anhydride may be
preserved for a long time in the open air, because it requires still
more heat to turn it into a gas than when in a liquid state.[9 bis]

  [6] As carbonic anhydride is one and a half times heavier than air, it
      diffuses with difficulty, and therefore does not easily mix with
      air, but sinks in it. This may be shown in various ways; for
      instance, the gas may be carefully poured from one vessel into
      another containing air. If a lighted taper be plunged into the
      vessel containing carbonic anhydride it is extinguished, and then,
      after pouring the gas into the other cylinder, it will burn in the
      former and be extinguished in the latter. If a certain quantity
      of carbonic anhydride be poured into a vessel containing air, and
      soap-bubbles be introduced, they will only sink as far as the
      stratum where the atmosphere of carbonic anhydride commences, as
      this latter is heavier than the soap-bubbles filled with air.
      Naturally, after a certain lapse of time, the carbonic anhydride
      will be diffused throughout the vessel, and form a uniform mixture
      with the air, just as salt in water.

  [7] This liquefaction was first observed by Faraday, who sealed up in
      a tube a mixture of a carbonate and sulphuric acid. Afterwards
      this method was very considerably improved by Thilorier and
      Natterer, whose apparatus is given in Chapter VI. in describing
      N_{2}O. It is, however, necessary to remark that the preparation
      of liquid carbonic anhydride requires good liquefying apparatus,
      constant cooling, and a rapid preparation of large masses of
      carbonic anhydride.

  [8] Carbonic anhydride, having the same molecular weight as nitrous
      oxide, very much resembles it when in a liquid state.

  [8 bis] When poured into a tube, which is then sealed up, liquefied
      carbonic anhydride can be easily preserved, because a thick tube
      easily supports the pressure (about 50 atmospheres) exerted by the
      liquid at the ordinary temperature.

  [9] When a fine stream of liquid carbonic anhydride is discharged into
      a closed metallic vessel, about one-third of its mass solidifies
      and the remainder evaporates. In employing solid carbonic
      anhydride for making experiments at low temperatures, it is best
      to use it mixed with ether, otherwise there will be few points
      of contact. If a stream of air be blown through a mixture of
      liquid carbonic anhydride and ether, the evaporation proceeds
      rapidly, and great cold is obtained. At present in some special
      manufactories (and for making artificial mineral waters)
      carbonic anhydride is liquefied on the large scale, filled into
      wrought-iron cylinders provided with a valve, and in this manner
      it can be transported and preserved safely for a long time. It is
      used, for instance, in breweries.

  [9 bis] Solid carbonic anhydride, notwithstanding its very low
      temperature, can be safely placed on the hand, because it
      continually evolves gas which prevents its coming into actual
      contact with the skin, but if a piece be squeezed between the
      fingers, it produces a severe frost bite similar to a burn. If the
      snow-like solid be mixed with ether, a semi-liquid mass is
      obtained, which is employed for artificial refrigeration. This
      mixture may be used for liquefying many other gases--such as
      chlorine, nitrous oxide, hydrogen sulphide, and others. The
      evaporation of such a mixture proceeds with far greater rapidity
      under the receiver of an air-pump, and consequently the
      refrigeration is more intense. By this means many gases may be
      liquefied which resist other methods--namely, olefiant gas,
      hydrochloric acid gas, and others. Liquid carbonic anhydride in
      this case congeals in the tube into a glassy transparent mass.
      Pictet availed himself of this method for liquefying many permanent
      gases (_see_ Chapter II.)

      Bleekrode, by compressing solid CO_{2} in a cylinder by means of a
      piston, obtained a semi-transparent stick, which contained as much
      as 1·3 and even 1·6 gram of CO_{2} per cubic centimetre. In this
      form the CO_{2} slowly evaporated, and could be kept for a long
      time.

The capacity which carbonic anhydride has of being liquefied stands in
connection with its _considerable solubility in water_, alcohol, and
other liquids. Its solubility in water has been already spoken of in the
first chapter. Carbonic anhydride is still more soluble in alcohol than
in water, namely at 0° one volume of alcohol dissolves 4·3 volumes of
this gas, and at 20° 2·9 volumes.

Aqueous solutions of carbonic anhydride, under a pressure of several
atmospheres, are now prepared artificially, because water saturated
with this gas promotes digestion and quenches thirst. For this purpose
the carbonic anhydride is pumped by means of a force-pump into a closed
vessel containing the liquid, and then bottled off, taking special means
to ensure rapid and air-tight corking. Various effervescing drinks and
artificially effervescing wines are thus prepared. The presence of
carbonic anhydride has an important significance in nature, because by
its means water acquires the property of decomposing and dissolving many
substances which are not acted on by pure water; for instance, calcium
phosphates and carbonates are soluble in water containing carbonic acid.
If the water in the interior of the earth is saturated with carbonic
acid under pressure, the quantity of calcium carbonate in solution may
reach three grams per litre, and on issuing at the surface, as the
carbonic anhydride escapes, the calcium carbonate will be deposited.[10]
Water charged with carbonic anhydride brings about the destruction of
many rocky formations by removing the lime, alkali, &c., from them. This
process has been going on and continues on an enormous scale. Rocks
contain silica and the oxides of various metals; amongst others, the
oxides of aluminium, calcium, and sodium. Water charged with carbonic
acid dissolves both the latter, transforming them into carbonates. The
waters of the ocean ought, as the evolution of the carbonic anhydride
proceeds, to precipitate salts of lime; these are actually found
everywhere on the surface of the ground in those places which previously
formed the bed of the ocean. The presence of carbonic anhydride in
solution in water is essential to the nourishment and growth of water
plants.

  [10] If such water trickles through crevices and enters a cavern, the
       evaporation will be slow, and therefore in those places from
       which the water drips growths of calcium carbonate will be
       formed, just like the icicles formed on the roof-gutters in
       winter-time. Similar conical and cylindrical stony growths form
       the so-called stalactites or pendants hanging from above and
       stalagmites formed on the bottom of caves. Sometimes these two
       kinds meet together, forming entire columns filling the cave.
       Many of these caves are remarkable for their picturesqueness; for
       instance, the cave of Antiparos, in the Grecian Archipelago. This
       same cause also forms spongy masses of calcium carbonate in those
       places where the springs come to the surface of the earth. It is
       therefore very evident that a calcareous solution is sometimes
       capable of penetrating plants and filling the whole of their mass
       with calcium carbonate. This is one of the forms of petrified
       plants. Calcium phosphate in solution in water containing
       carbonic acid plays an important part in the nourishment of
       plants, because all plants contain both lime and phosphoric acid.

Although carbonic anhydride is soluble in water, yet no definite hydrate
is formed;[11] nevertheless an idea of the composition of this hydrate
may be formed from that of the salts of carbonic acid, because a hydrate
is nothing but a salt in which the metal is replaced by hydrogen.
As carbonic anhydride forms salts of the composition K_{2}CO_{3},
Na_{2}CO_{3}, HNaCO_{3}, &c., therefore carbonic acid ought to have
the composition H_{2}CO_{3}--that is, it ought to contain CO_{2} +
H_{2}O. Whenever this substance is formed, it decomposes into its
component parts--that is, into water and carbonic anhydride. _The acid
properties_ of carbonic anhydride[11 bis] are demonstrated by its being
directly absorbed by alkaline solutions and forming salts with them. In
distinction from nitric, HNO_{3}, and similar monobasic acids which with
univalent metals (exchanging one atom for one atom of hydrogen) give
salts such as those of potassium, sodium, and silver containing only one
atom of the metal (NaNO_{3}, AgNO_{3}), and with bivalent[12] metals
(such as calcium, barium, lead) salts containing two acid groups--for
example, Ca(NO_{3})_{2}, Pb(NO_{3})_{2}--carbonic acid, H_{2}CO_{3},
_is bibasic_, that is contains two atoms of hydrogen in the hydrate or
two atoms of univalent metals in their salts: for example, Na_{2}CO_{3}
is washing soda, a normal salt; NaHCO_{3} is the bicarbonate, an
acid salt. Therefore, if M´ be a univalent metal, its carbonates in
general are the normal carbonate M´_{2}CO_{3} and the acid carbonate,
M´HCO_{3}; or if M´´ be a bivalent metal (replacing H_{2}) its normal
carbonate will be M´´CO_{3}; these metals do not usually form acid
salts, as we shall see further on. The bibasic character of carbonic
acid is akin to that of sulphuric acid, H_{2}SO_{4},[13] but the
latter, in distinction from the former, is an example of the energetic
or strong acids (such as nitric or hydrochloric), whilst in carbonic
acid we observe but feeble development of the acid properties; hence
carbonic acid must be considered _a weak acid_. This conception must,
however, be taken as only comparative, as up to this time there is no
definitely established rule for measuring the energy[14] of acids. The
feeble acid properties of carbonic acid may, however, be judged from
the joint evidence of many properties. With such energetic alkalis
as soda and potash, carbonic acid forms normal salts, soluble in
water, but having an alkaline reaction and in many cases themselves
acting as alkalis.[15] The acid salts of these alkalis, NaHCO_{3} and
KHCO_{3}, have a neutral reaction on litmus, although they, like acids,
contain hydrogen, which may be exchanged for metals. The acid salts
of such acids--as, for instance, of sulphuric acid, NaHSO_{4}--have a
clearly defined acid reaction, and therefore carbonic acid is unable
to neutralise the powerful basic properties of such alkalis as potash
or soda. Carbonic acid does not even combine at all with feeble bases,
such as alumina, Al_{2}O_{3}, and therefore if a strong solution of
sodium carbonate, Na_{2}CO_{3}, be added to a strong solution of
aluminium sulphate, Al_{2}(SO_{4})_{3}, although according to double
saline decompositions aluminium carbonate, Al_{2}(CO_{3})_{3}, ought
to be formed, the carbonic acid separates, for this salt splits up in
the presence of water into aluminium hydroxide and carbonic anhydride:
Al_{2}(CO_{3})_{3} + 3H_{2}O = Al_{2}(OH)_{6} + 3CO_{2}. Thus feeble
bases are unable to retain carbonic acid even at ordinary temperatures.
For the same reason, in the case of bases of medium energy, although
they form carbonates, the latter are comparatively easily decomposed by
heating, as is shown by the decomposition of copper carbonate, CuCO_{3}
(_see_ Introduction), and even of calcium carbonate, CaCO_{3}. Only
the normal (not the acid) salts of such powerful bases as potassium
and sodium are capable of standing a red heat without decomposition.
The acid salts--for instance, NaHCO_{3}--decompose even on heating
their solutions (2NaHCO_{3} = Na_{2}CO_{3} + H_{2}O + CO_{2}), evolving
carbonic anhydride. The amount of heat given out by the combination of
carbonic acid with bases also shows its feeble acid properties, being
considerably less than with energetic acids. Thus if a weak solution
of forty grams of sodium hydroxide be saturated (up to the formation
of a normal salt) with sulphuric or nitric acid or another powerful
acid, from thirteen to fifteen thousand calories are given out, but
with carbonic acid only about ten thousand calories.[16] The majority
of carbonates are insoluble in water, and therefore such solutions as
sodium, potassium, or ammonium carbonates form in solutions of most
other salts, MX or M´´X_{2}, insoluble precipitates of carbonates,
M_{2}CO_{3} or M´´CO_{3}. Thus a solution of barium chloride gives with
sodium carbonate a precipitate of barium carbonate, BaCO_{3}. For this
reason rocks, especially those of aqueous origin, very often contain
carbonates; for example, calcium, ferrous, or magnesium carbonates, &c.

  [11] The crystallohydrate, CO_{2},8H_{2}O of Wroblewski (Chapter 1.,
       Note 67), in the first place, is only formed under special
       conditions; in the second place, its existence still requires
       confirmation; and in the third place, it does not correspond with
       that hydrate H_{2}CO_{3} which should occur, judging from the
       composition of the salts.

  [11 bis] It is easy to demonstrate the acid properties of carbonic
       anhydride by taking a long tube, closed at one end, and filling
       it with this gas; a test-tube is then filled with a solution of
       an alkali (for instance, sodium hydroxide), which is then poured
       into the long tube and the open end is corked. The solution is
       then well shaken in the tube, and the corked end plunged into
       water. If the cork be now withdrawn under water, the water will
       fill the tube. The vacuum obtained by the absorption of the
       carbonic anhydride by an alkali is so complete that even an
       electric discharge will not pass through it. This method is often
       applied to produce a vacuum.

  [12] The reasons for distinguishing the uni-, bi-, tri-, and
       quadrivalent metals will be explained hereafter on passing from
       the univalent metals (Na, K, Li) to the bivalent (Mg. Ca, Ba),
       Chapter XIV.

  [13] Up to the year 1840, or thereabout, acids were not distinguished
       by their basicity. Graham, while studying phosphoric acid,
       H_{3}PO_{4}, and Liebig, while studying many organic acids,
       distinguished mono-, bi-, and tribasic acids. Gerhardt and Laurent
       generalised these relations, showing that this distinction extends
       over many reactions (for instance, to the faculty of bibasic acids
       of forming acid salts with alkalis, KHO or NaHO, or with alcohols,
       RHO, &c.); but now, since a definite conception as to atoms and
       molecules has been arrived at, _the basicity of an acid is
       determined by the number of hydrogen atoms,_ contained in a
       molecule of the acid, which can be exchanged for metals. If
       carbonic acid forms acid salts, NaHCO_{3}, and normal salts,
       Na_{2}CO_{3}, it is evident that the hydrate is H_{2}CO_{3}, a
       bibasic acid. Otherwise it is at present impossible to account for
       the composition of these salts. But when C = 6 and O = 8 were
       taken, then the formula CO_{2} expressed the composition, but not
       the molecular weight, of carbonic anhydride; and the composition
       of the normal salt would be Na_{2}C_{2}O_{6} or NaCO_{3},
       therefore carbonic acid might have been considered as a monobasic
       acid. Then the acid salt would have been represented by
       NaCO_{3},HCO_{3}. Such questions were the cause of much argument
       and difference of opinion among chemists about forty years ago. At
       present there cannot be two opinions on the subject if the law of
       Avogadro-Gerhardt and its consequences be strictly adhered to. It
       may, however, be observed here that the monobasic acids R(OH) were
       for a long time considered to be incapable of being decomposed
       into water and anhydride, and this property was ascribed to the
       bibasic acids R(OH)_{2} as containing the elements necessary for
       the separation of the molecule of water, H_{2}O. Thus H_{2}SO_{4}
       or SO_{2}(OH)_{2}, H_{2}CO_{3}, or CO(OH)_{2}, and other bibasic
       acids decompose into an anhydride, RO, and water, H_{2}O. But as
       nitrous, HNO_{2}, iodic, HIO_{3}, hypochlorous, HClO, and other
       monobasic acids easily give their anhydrides N_{2}O_{3},
       I_{2}O_{5}, Cl_{2}O, &c., that method of distinguishing the
       basicity of acids, although it fairly well satisfies the
       requirements of organic chemistry, cannot be considered correct.
       It may also be remarked that up to the present time not one of the
       bibasic acids has been found to have the faculty of being
       distilled without being decomposed into anhydride and water (even
       H_{2}SO_{4}, on being evaporated and distilled, gives SO_{3} +
       H_{2}O), and the decomposition of acids into water and anhydride
       proceeds particularly easily in dealing with feebly energetic
       acids, such as carbonic, nitrous, boric, and hypochlorous. Let us
       add that carbonic acid, as a hydrate corresponding to marsh gas,
       C(HO)_{4} = CO_{2} + 2H_{2}O, ought to be tetrabasic. But in
       general it does not form such salts. Basic salts, however, such as
       CuCO_{3}CuO, may be regarded in this sense, for CCu_{2}O_{4}
       corresponds with CH_{4}O_{4}, as Cu corresponds with H_{2}.
       Amongst the ethereal salts (alcoholic derivatives) of carbonic
       acid corresponding cases are, however, observed; for instance,
       ethylic orthocarbonate, C(C_{2}H_{5}O)_{4} (obtained by the action
       of chloropicrin, C(NO_{2})Cl_{3}, on sodium ethoxide,
       C_{2}H_{5}ONa; boiling point 158°; specific gravity, 0·92). The
       name _orthocarbonic acid_ for CH_{4}O_{4} is taken from
       _orthophosphoric acid_, PH_{3}O_{4}, which corresponds with PH_{3}
       (_see_ Chapter on Phosphorus).

  [14] Long ago endeavours were made to find a _measure of affinity_ of
       acids and bases, because some of the acids, such as sulphuric or
       nitric, form comparatively stable salts, decomposed with
       difficulty by heat and water, whilst others, like carbonic and
       hypochlorous acids, do not combine with feeble bases, and with
       most of the other bases form salts which are easily decomposed.
       The same may be said with regard to bases, among which those of
       potassium, K_{2}O, sodium, Na_{2}O, and barium, BaO, may serve as
       examples of the most powerful, because they combine with the most
       feeble acids and form a mass of salts of great stability, whilst
       as examples of the feeblest bases alumina, Al_{2}O_{3}, or bismuth
       oxide, Bi_{2}O_{3}, may be taken, because they form salts easily
       decomposed by water and by heat if the acid be volatile. Such a
       division of acids and bases into the feeblest and most powerful is
       justified by all evidence concerning them, and is quoted in this
       work. But the teaching of this subject in certain circles has
       acquired quite a new tone, which, in my opinion, cannot be
       accepted without certain reservations and criticisms, although it
       comprises many interesting features. The fact is that Thomsen,
       Ostwald, and others proposed to express the measure of affinity of
       acids to bases by figures drawn from data of the measure of
       displacement of acids in aqueous solutions, judging (1) from the
       amount of heat developed by mixing a solution of the salt with a
       solution of another acid (the avidity of acids, according to
       Thomsen); (2) from the change of the volumes accompanying such a
       mutual action of solutions (Ostwald); (3) from the change of the
       index of refraction of solutions (Ostwald), &c. Besides this there
       are many other methods which allow us to form an opinion about the
       distribution of bases among various acids in aqueous solutions.
       Some of these methods will be described hereafter. It ought,
       however, to be remarked that in making investigations in aqueous
       solutions the affinity to water is generally left out of sight. If
       a base N, combining with acids X and Y in presence of them both,
       divides in such a way that one-third of it combines with X and
       two-thirds with Y, a conclusion is formed that the affinity, or
       power of forming salts, of the acid Y is twice as great as that of
       X. But the presence of the water is not taken into account. If the
       acid X has an affinity for water and for N it will be distributed
       between them; and if X has a greater affinity for water than Y,
       then less of X will combine with N than of Y. If, in addition to
       this, the acid X is capable of forming an acid salt NX_{2}, and Y
       is not, the conclusion of the relative strength of X and Y will be
       still more erroneous, because the X set free will form such a salt
       on the addition of Y to NX. We shall see in Chapter X. that when
       sulphuric and nitric acids in weak aqueous solution act on sodium,
       they are distributed exactly in this way: namely, one-third of the
       sodium combines with the sulphuric and two-thirds with the nitric
       acid; but, in my opinion, this does not show that sulphuric acid,
       compared with nitric acid, possesses but half the degree of
       affinity for bases like soda, and only demonstrates the greater
       affinity of sulphuric acid for water compared with that of nitric
       acid. In this way the methods of studying the distribution in
       aqueous solutions probably only shows the difference of the
       relation of the acid to a base and to water.

       In view of these considerations, although the teaching of the
       distribution of salt-forming elements in _aqueous solutions_ is an
       object of great and independent interest, it can hardly serve to
       determine the measure of affinity between bases and acids. Similar
       considerations ought to be kept in view when determining the
       energy of acids by means of the _electrical conductivity of their
       weak solutions_. This method, proposed by Arrhenius (1884), and
       applied on an extensive scale by Ostwald (who developed it in
       great detail in his _Lehrbuch d. allgemeinen Chemie_, v. ii.,
       1887), is founded on the fact that the relation of the so-called
       molecular electrical-conductivity of weak solutions of various
       acids (I) coincides with the relation in which the same acids
       stand according to the distribution, (II) found by one of the
       above-mentioned methods, and with the relation deduced for them
       from observations upon the velocity of reaction, (III) for
       instance, according to the rate of the splitting up of an ethereal
       salt (into alcohol and acid), or from the rate of the so-called
       inversion of sugar--that is, its transformation into glucose--as
       is seen by comparing the annexed figures, in which the energy of
       hydrochloric acid is taken as equal to 100:--

                                              I    II  III

              Hydrochloric acid, HCl         100  100  100
              Hydrobromic acid, HBr          101   98  105
              Nitric acid, HNO_{3}           100  100   96
              Sulphuric acid, H_{2}SO_{4}     65   49   74
              Formic acid, CH_{2}O_{2}         2    4    1
              Acetic acid, C_{2}H_{4}O_{2}     1    2    1
              Oxalic acid, C_{2}H_{2}O_{4}    20   24   18
              Phosphoric acid, PH_{3}O_{4}     7   --    6

       The coincidence of these figures, obtained by so many various
       methods, presents a most important and instructive relation
       between phenomena of different kinds, but in my opinion it does
       not permit us to assert that the degree of affinity existing
       between bases and various acids is determined by all these various
       methods, because the influence of the water must be taken into
       consideration. On this account, until the theory of solution is
       more thoroughly worked out, this subject (which for the present
       ought to be treated of in special treatises on chemical mechanics)
       must be treated with great caution. But now we may hope to decide
       this question guided by a study of the rate of reaction, the
       influence of acids and bases upon indicators, &c., all of which
       are treated fully in works on physical and theoretical chemistry.

  [15] Thus, for instance, in the washing of fabrics the caustic alkalis,
       such as sodium hydroxide, in weak solutions, act in removing the
       fatty matter just in the same way as carbonate solutions; for
       instance, a solution of soda crystals, Na_{2}CO_{3}. Soap acts in
       the same way, being composed of feeble acids, either fatty or
       resinous, combined with alkali. On this account all such
       substances are applied in manufacturing processes, and answer
       equally well in practice for bleaching and washing fabrics. Soda
       crystals or soap are preferred to caustic alkali, because an
       excess of the latter may have a destructive effect on the fabrics.
       It may be supposed that in aqueous solutions of soap or soda
       crystals, part of the base will form caustic alkali; that is to
       say, the water will compete with the weak acids, and the alkali
       will be distributed between them and the water.

  [16] Although carbonic acid is reckoned among the feeble acids, yet
       there are evidently many others still feebler--for instance,
       prussic acid, hypochlorous acid, many organic acids, &c. Bases
       like alumina, or such feeble acids as silica, when in combination
       with alkalis, are decomposed in aqueous solutions by carbonic
       acid, but on fusion--that is, without the presence of water--they
       displace it, which clearly shows in phenomena of this kind how
       much depends upon the conditions of reaction and the properties of
       the substances formed. These relations, which at first sight
       appear complex, may be best understood if we represent that two
       salts, MX and NY, in general always give more or less of two other
       salts, MY and NX, and then examine the properties of the derived
       substances. Thus, in solution, sodium silicate, Na_{2}SiO_{3},
       with carbonic anhydride will to some extent form sodium carbonate
       and silica, SiO_{2}; but the latter, being colloid, separates, and
       the remaining mass of sodium silicate is again decomposed by
       carbonic anhydride, so that finally silica separates and sodium
       carbonate is formed. In a fused state the case is different;
       sodium carbonate will react with silica to form carbonic anhydride
       and sodium silicate, but the carbonic anhydride will be separated
       as a gas, and therefore in the residue the same reaction will
       again take place, and ultimately the carbonic anhydride is
       entirely eliminated and sodium silicate remains. If, on the other
       hand, nothing is removed from the sphere of the reaction,
       distribution takes place. Therefore, although carbonic anhydride
       is a feeble acid, still not for this reason, but only in virtue of
       its gaseous form, do all soluble acids displace it in saline
       solutions (_see_ Chapter X.)

Carbonic anhydride--which, like water, is formed with the development
of a large amount of heat--is very stable. Only very few substances
are capable of depriving it of its oxygen. However, certain metals,
such as magnesium, potassium and the like, on being heated, burn in
it, depositing carbon and forming oxides. If a mixture of carbonic
anhydride and hydrogen be passed through a heated tube, the formation
of water and carbonic oxide will be observed; CO_{2} + H_{2} = CO
+ H_{2}O. But only a portion of the carbonic acid gas undergoes
this change, and therefore the result will be a mixture of carbonic
anhydride, carbonic oxide, hydrogen, and water, which does not suffer
further change under the action of heat.[17] Although, like water,
carbonic anhydride is exceedingly stable, still on being heated it
partially decomposes into carbonic oxide and oxygen. Deville showed that
such is the case if carbonic anhydride be passed through a long tube
containing pieces of porcelain and heated to 1,300°. If the products
of decomposition--namely, the carbonic oxide and oxygen--be suddenly
cooled, they can be collected separately, although they partly reunite
together. A similar decomposition of carbonic anhydride into carbonic
oxide and oxygen takes place on passing a series of electric sparks
through it (for instance, in the eudiometer). Under these conditions
an increase of volume occurs, because two volumes of CO_{2} give two
volumes of CO and one volume of O. The decomposition reaches a certain
limit (less than one-third) and does not proceed further, so that the
result is a mixture of carbonic anhydride, carbonic oxide, and oxygen,
which is not altered in composition by the continued action of the
sparks. This is readily understood, as it is a reversible reaction.
If the carbonic anhydride be removed, then the mixture explodes when
a spark is passed and forms carbonic anhydride.[17 bis] If from an
identical mixture the oxygen (and not the carbonic anhydride) be
removed, and a series of sparks be again passed, the decomposition is
renewed, and terminates with the complete dissociation of the carbonic
anhydride. Phosphorus is used in order to effect the complete absorption
of the oxygen. In these examples we see that a definite mixture of
changeable substances is capable of arriving at a state of stable
equilibrium, destroyed, however, by the removal of one of the substances
composing the mixture. This is one of the instances of the influence of
mass.

  [17] Hydrogen and carbon are near akin to oxygen as regards affinity,
       but it ought to be considered that the affinity of hydrogen is
       slightly greater than that of carbon, because during the
       combustion of hydrocarbons the hydrogen burns first. Some idea of
       this similarity of affinity may be formed by the quantity of heat
       evolved. Gaseous hydrogen, H_{2}, on combining with an atom of
       oxygen, O = 16, develops 69,000 heat-units if the water formed be
       condensed to a liquid state. If the water remains in the form of a
       gas (steam) the latent heat of evaporation must be subtracted, and
       then 58,000 calories will be developed. Carbon, C, as a solid, on
       combining with O_{2} = 32 develops about 97,000 calories, forming
       gaseous CO_{2}. If it were gaseous like hydrogen, and only
       contained C_{2} in its molecule, much more heat would be
       developed, and judging by other substances, whose molecules on
       passing from the solid to the gaseous state absorb about 10,000 to
       15,000 calories, it must be held that gaseous carbon on forming
       gaseous carbonic anhydride would develop not less than 110,000
       calories--that is, approximately twice as much as is developed in
       the formation of water. And since there is twice as much oxygen in
       a molecule of carbonic anhydride as in a molecule of water, the
       oxygen develops approximately the same quantity of heat on
       combining with hydrogen as with carbon. That is to say, that here
       we find the same close affinity (_see_ Chapter II., Note 7)
       determined by the quantity of heat as between hydrogen, zinc, and
       iron. For this reason here also, as in the case of hydrogen and
       iron, we ought to expect an equal distribution of oxygen between
       hydrogen and carbon, if they are both in excess compared with the
       amount of oxygen; but if there be an excess of carbon, it will
       decompose water, whilst an excess of hydrogen will decompose
       carbonic anhydride. Even if these phenomena and similar ones have
       been explained in isolated cases, a complete theory of the whole
       subject is still wanting in the present condition of chemical
       knowledge.

  [17 bis] The degree or relative magnitude of the dissociation of CO_{2}
       varies with the temperature and pressure--that is, it increases
       with the temperature and as the pressure decreases. Deville found
       that at a pressure of 1 atmosphere in the flame of carbonic oxide
       burning in oxygen, about 40 per cent. of the CO_{2}, is decomposed
       when the temperature is about 3,000°, and at 1,500° less than 1
       per cent. (Krafts); whilst under a pressure of 10 atmospheres
       about 34 per cent. is decomposed at 3,300° (Mallard and Le
       Chatelier). It follows therefore that, under very small pressures,
       the dissociation of CO_{2} will be considerable even at
       comparatively moderate temperatures, but at the temperature of
       ordinary furnaces (about 1,000°) even under the small partial
       pressure of the carbonic acid, there are only small traces of
       decomposition which may be neglected in a practical estimation of
       the combustion of fuels. We may here cite the molecular specific
       heat of CO_{2} (_i.e._ the amount of heat required to raise 44
       units of weight of CO_{2} 1°), according to the determinations and
       calculations of Mallard and Le Chatelier, for a constant volume
       C_{v} = 6·26 + 0·0037_t_; for a constant pressure C_{p} = C_{v} +
       2 (_see_ Chapter XIV., Note 7), _i.e._ the specific heat of CO_{2}
       increases rapidly with a rise of temperature: for example, at 0°
       (per 1 part by weight), it is, at a constant pressure = 0·188, at
       1,000° = 0·272, at 2,000°, about 0·356. A perfectly distinct rise
       of the specific heat (for example, at 2,000°, 0·409), is given by
       a comparison of observations made by the above-mentioned
       investigators and by Berthelot and Vieille (Kournakoff). The cause
       of this must be looked for in dissociation. T. M. Cheltzoff,
       however, considers upon the basis of his researches upon
       explosives that it must be admitted that a maximum is reached at a
       certain temperature (about 2,500°), beyond which the specific heat
       begins to fall.

Although carbonic anhydride is decomposed on heating, yielding
oxygen, it is nevertheless, like water, an unchangeable substance at
ordinary temperatures. Its decomposition, as effected by plants, is on
this account all the more remarkable; in this case the whole of the
oxygen of the carbonic anhydride is separated in the free state. The
mechanism of this change is that the heat and light absorbed by the
plants are expended in the decomposition of the carbonic anhydride.
This accounts for the enormous influence of temperature and light on
the growth of plants. But it is at present not clearly understood how
this takes place, or by what separate intermediate reactions the whole
process of decomposition of carbonic anhydride in plants into oxygen and
the carbohydrates (Note 1) remaining in them, takes place. It is known
that sulphurous anhydride (in many ways resembling carbonic anhydride)
under the action of light (and also of heat) forms sulphur and sulphuric
anhydride, SO_{3}, and in the presence of water, sulphuric acid. But
no similar decomposition has been obtained directly with carbonic
anhydride, although it forms an exceedingly easily decomposable higher
oxide--percarbonic acid;[18] and perhaps that is the reason the oxygen
separates. On the other hand, it is known that plants always form and
contain _organic acids_, and these must be regarded as derivatives of
carbonic acid, as is seen by all their reactions, of which we will
shortly treat. For this reason it might be thought that the carbonic
acid absorbed by the plants first forms (according to Baeyer) formic
aldehyde, CH_{2}O, and from it organic acids, and that these latter
in their final transformation form all the other complex organic
substances of the plants. Many organic acids are found in plants in
considerable quantity; for instance, tartaric acid, C_{4}H_{6}O_{6},
found in grape-juice and in the acid juice of many plants; malic acid,
C_{4}H_{6}O_{5}, found not only in unripe apples but in still larger
quantities in mountain ash berries; citric acid, C_{6}H_{8}O_{7}, found
in the acid juice of lemons, in gooseberries, cranberries, &c.; oxalic
acid, C_{2}H_{2}O_{4}, found in wood-sorrel and many other plants.
Sometimes these acids exist in a free state in the plants, and sometimes
in the form of salts; for instance, tartaric acid is met with in grapes
as the salt known as cream of tartar, but in the impure state called
argol, or tartar, C_{4}H_{5}KO_{6}. In sorrel we find the so-called
salts of sorrel, or acid potassium oxalate, C_{2}HKO_{4}. There is a
very clear connection between carbonic anhydride and the above-mentioned
organic acids--namely, they all, under one condition or another, yield
carbonic anhydride, and can all be formed by means of it from substances
destitute of acid properties. The following examples afford the best
demonstration of this fact: if acetic acid, C_{2}H_{4}O_{2}, the acid
of vinegar, be passed in the form of vapour through a heated tube, it
splits up into carbonic anhydride and marsh gas = CO_{2} + CH_{4}. But
conversely it can also be obtained from those components into which it
decomposes. If one equivalent of hydrogen in marsh gas be replaced (by
indirect means) by sodium, and the compound CH_{3}Na is obtained, this
directly absorbs carbonic anhydride, forming a salt of acetic acid,
CH_{3}Na + CO_{2} = C_{2}H_{3}NaO_{2}; from this acetic acid itself
may be easily obtained. Thus acetic acid decomposes into marsh gas
and carbonic anhydride, and conversely is obtainable from them. The
hydrogen of marsh gas does not, like that in acids, show the property
of being directly replaced by metals; _i.e._ CH_{4} does not show any
acid character whatever, but on combining with the elements of carbonic
anhydride it acquires the properties of an acid. The investigation
of all other organic acids shows similarly that their acid character
depends on their containing the elements of carbonic anhydride. For this
reason there is no organic acid containing less oxygen in its molecule
than there is in carbonic anhydride; every organic acid contains in
its molecule at least two atoms of oxygen. In order to express the
relation between carbonic acid, H_{2}CO_{3}, and organic acids, and in
order to understand the reason of the acidity of these latter, it is
simplest to turn to that law of substitution which shows (Chapter VI.)
the relation between the hydrogen and oxygen compounds of nitrogen,
and permits us (Chapter VIII.) to regard all hydrocarbons as derived
from methane. If we have a given organic compound, A, which has not the
properties of an acid, but contains hydrogen connected to carbon, as in
hydrocarbons, then ACO_{2} will be a monobasic organic acid, A2CO_{2}
a bibasic, A3CO_{2} a tribasic, and so on--that is, each molecule of
CO_{2} transforms one atom of hydrogen into that state in which it may
be replaced by metals, as in acids. This furnishes a direct proof that
in organic acids it is necessary to recognise the group HCO_{2}, or
carboxyl. If the addition of CO_{2} raises the basicity, the removal of
CO_{2} lowers it. Thus from the bibasic oxalic acid, C_{2}H_{2}O_{4}, or
phthalic acid, C_{8}H_{6}O_{4}, by eliminating CO_{2} (easily effected
experimentally) we obtain the monobasic formic acid, CH_{2}O_{2}, or
benzoic acid, C_{7}H_{6}O_{2}, respectively. The nature of carboxyl is
directly explained by the law of substitution. Judging from what has
been stated in Chapters VI. and VIII. concerning this law, it is evident
that CO_{2} is CH_{4} with the exchange of H_{4} for O_{2}, and that
the hydrate of carbonic anhydride, H_{2}CO_{3}, is CO(OH)_{2}, that is,
methane, in which two parts of hydrogen are replaced by two parts of
the water radical (OH, hydroxyl) and the other two by oxygen. Therefore
the group CO(OH), or carboxyl, HCO_{2}, is a part of carbonic acid,
and is equivalent to (OH), and therefore also to H. That is, it is a
univalent residue of carbonic acid capable of replacing one atom of
hydrogen. Carbonic acid itself is a bibasic acid, both hydrogen atoms
in it being replaceable by metals, therefore carboxyl, which contains
one of the hydrogen atoms of carbonic acid, represents a group in which
the hydrogen is exchangeable for metals. And therefore if 1, 2 ... _n_
atoms of non-metallic hydrogen are exchanged 1, 2 ... _n_ times for
carboxyl, we ought to obtain 1, 2 ... _n_-basic acids. _Organic acids
are the products of the carboxyl substitution in hydrocarbons._[18 bis]
If in the saturated hydrocarbons, C_{n}H_{2n + 2}, one part of
hydrogen is replaced by carboxyl, the monobasic saturated (or fatty)
acids, C_{n}H_{2n + 1}(CO_{2}H), will be obtained, as, for instance,
formic acid, HCO_{2}H, acetic acid, CH_{2}CO_{2}H, ... stearic acid,
C_{17}H_{35}CO_{2}H, &c. The double substitution will give bibasic
acids, C_{n}H_{2n}(CO_{2}H)(CO_{2}H); for instance, oxalic acid _n_
= 0, malonic acid _n_ = 1, succinic acid _n_ = 2, &c. To benzene,
C_{6}H_{6} correspond benzoic acid, C_{6}H_{5}(CO_{2}H), phthalic acid
(and its isomerides), C_{6}H_{4}(CO_{2}H)_{2}, up to mellitic acid,
C_{6}(CO_{2}H)_{6}, in all of which the basicity is equal to the number
of carboxyl groups. As many isomerides exist in hydrocarbons, it is
readily understood not only that such can exist also in organic acids,
but that their number and structure may be foreseen. This complex and
most interesting branch of chemistry is treated separately in organic
chemistry.

  [18] Percarbonic acid, H_{2}CO_{4} (= H_{2}CO_{3} + O) is supposed
       by A. Bach (1893) to be formed from carbonic acid in the action of
       light upon plants, (in the same manner as, according to the above
       scheme, sulphuric acid from sulphurous) with the formation of
       carbon, which remains in the form of hydrates of carbon:
       3H_{2}CO_{3} = 2H_{2}CO_{4} + CH_{2}O. This substance CH_{2}O
       expresses the composition of formic aldehyde which, according to
       Baeyer, by polymerisation and further changes, gives other
       hydrates of carbon and forms the first product which is formed in
       plants from CO_{2}. And Berthelot (1872) had already, at the time
       of the discovery of persulphuric (Chapter XX.) and pernitric
       (Chapter VI., Note 26) acids pointed out the formation of the
       unstable percarbonic anhydride, CO_{3}. Thus, notwithstanding the
       hypothetical nature of the above equation, it may be admitted all
       the more as it explains the comparative abundance of peroxide of
       hydrogen (Schöne, Chapter IV.) in the air, and this also at the
       period of the most energetic growth of plants (in July), because
       percarbonic acid should like all peroxides easily give H_{2}O_{2}.
       Besides which Bach (1894) showed that, in the first place, traces
       of formic aldehyde and oxidising agents (CO_{3} or H_{2}O_{2})
       are formed under the simultaneous action of CO_{2} and sunlight
       upon a solution containing a salt of uranium (which is oxidised),
       and diethylaniline (which reacts with CH_{2}O), and secondly, that
       by subjecting BaO_{2}, shaken up in water, to the action of a
       stream of CO_{2} in the cold, extracting (also in the cold) with
       ether, and then adding an alcoholic solution of NaHO, crystalline
       plates of a sodium salt may be obtained, which with water evolve
       oxygen and leave sodium carbonate; they are therefore probably the
       per-salt. All these facts are of great interest and deserve
       further verification and elaboration.

  [18 bis] If CO_{2} is the anhydride of a bibasic acid, and carboxyl
       corresponds with it, replacing the hydrogen of hydrocarbons, and
       giving them the character of comparatively feeble acids, then
       SO_{3} is the anhydride of an energetic bibasic acid, and
       _sulphoxyl_, SO_{2}(OH), corresponds with it, being capable of
       replacing the hydrogen of hydrocarbons, and forming comparatively
       energetic _sulphur oxyacids_ (_sulphonic acids_); for instance,
       C_{6}H_{5}(COOH), benzoic acid, and C_{6}H_{5}(SO_{2}OH),
       benzenesulphonic acid, are derived from C_{6}H_{6}. As the
       exchange of H for methyl, CH_{3}, is equivalent to the addition of
       CH_{2}, the exchange of carboxyl, COOH, is equivalent to the
       addition of CO_{2}; so the exchange of H for sulphoxyl is
       equivalent to the addition of SO_{3}. The latter proceeds
       directly, for instance: C_{6}H_{6} + SO_{3} =
       C_{6}H_{5}(SO_{2}OH).

       As accordingding to the determinations of Thomsen, the heat of
       combustion of the _vapours_ of acids RCO_{2} is known where R is a
       hydrocarbon, and the heat of combustion of the hydrocarbons R
       themselves, it may be seen that the formation of acids, RCO_{2},
       from R + CO_{2}, is always accompanied by a _small_ absorption or
       development of heat. We give the heats of combustion in thousands
       of calories, referred to the molecular weights of the
       substances:--

            R =        H_{2}  CH_{4}  C_{2}H_{6}  C_{6}H_{6}
                       68·4    212       370         777
            RCO_{2} =  69·4    225       387         766

       Thus H_{2}, corresponds with formic acid, CH_{2}O_{2}; benzene,
       C_{6}H_{6}, with benzoic acid, C_{7}H_{6}O_{2}. The data for the
       latter are taken from Stohmann, and refer to the solid condition.
       For formic acid Stohmann gives the heat of combustion as 59,000
       calories in a liquid state, but in a state of vapour, 64·6
       thousand units, which is much less than according to Thomsen.

_Carbonic Oxide._--This gas is formed whenever the combustion of
organic substances takes place in the presence of a large excess of
incandescent charcoal; the air first burns the carbon into carbonic
anhydride, but this in penetrating through the red-hot charcoal is
transformed into carbonic oxide, CO_{2} + C = 2CO. By this reaction
carbonic oxide is prepared by passing carbonic anhydride through
charcoal at a red heat. It may be separated from the excess of carbonic
anhydride by passing it through a solution of alkali, which does not
absorb carbonic oxide. This reduction of carbonic anhydride explains why
carbonic oxide is formed in ordinary clear fires, where the incoming
air passes over a large surface of heated coal. A blue flame is then
observed burning above the coal; this is the burning carbonic oxide.
When charcoal is burnt in stacks, or when a thick layer of coal is
burning in a brazier, and under many similar circumstances, carbonic
oxide is also formed. In metallurgical processes, for instance when
iron is smelted from the ore, very often the same process of conversion
of carbonic anhydride into carbonic oxide occurs, especially if the
combustion of the coal be effected in high, so-called blast, furnaces
and ovens, where the air enters at the lower part and is compelled
to pass through a thick layer of incandescent coal. In this way,
also, combustion with flame may be obtained from those kinds of fuel
which under ordinary conditions burn without flame: for instance,
anthracite, coke, charcoal. Heating by means of a gas-producer--that
is, an apparatus producing combustible carbonic oxide from fuel--is
carried on in the same manner.[19] In transforming one part of charcoal
into carbonic oxide 2,420 heat units are given out, and on burning to
carbonic anhydride 8,080 heat units. It is evident that on transforming
the charcoal first into carbonic oxide we obtain a gas which in burning
is capable of giving out 5,660 heat units for one part of charcoal. This
preparatory transformation of fuel into carbonic oxide, or producer
gas containing a mixture of carbonic oxide (about 1/3 by volume) and
nitrogen (2/3 volume), in many cases presents most important advantages,
as it is easy to completely burn gaseous fuel without an excess of
air, which would lower the temperature.[20] In stoves where solid fuel
is burnt it is impossible to effect the complete combustion of the
various kinds of fuel without admitting an excess of air. Gaseous fuel,
such as carbonic oxide, is easily completely mixed with air and burnt
without excess of it. If, in addition to this, the air and gas required
for the combustion be previously heated by means of the heat which
would otherwise be uselessly carried off in the products of combustion
(smoke)[21] it is easy to reach a high temperature, so high (about
1,800°) that platinum may be melted. Such an arrangement is known as a
_regenerative furnace_.[22] By means of this process not only may the
high temperatures indispensable in many industries be obtained (for
instance, glass-working, steel-melting, &c.), but great advantage
also[23] is gained as regards the quantity of fuel, because the
transmission of heat to the object to be heated, other conditions being
equal, is determined by the difference of temperatures.

  [19] [Illustration: FIG. 63.--Gas-producer for the formation of carbon
       monoxide for heating purposes.]

       In gas-producers all carbonaceous fuels are transformed into
       inflammable gas. In those which (on account of their slight
       density and large amount of water, or incombustible admixtures
       which absorb heat) are not as capable of giving a high temperature
       in ordinary furnaces--for instance, fir cones, peat, the lower
       kinds of coal, &c.--the same gas is obtained as with the best
       kinds of coal, because the water condenses on cooling, and the
       ashes and earthy matter remain in the gas-producer. The
       construction of a gas-producer is seen from the accompanying
       drawing. The fuel lies on the fire-bars O, the air enters through
       them and the ash-hole (drawn by the draught of the chimney of the
       stove where the gas burns, or else forced by a blowing apparatus),
       the quantity of air being exactly regulated by means of valves.
       The gases formed are then led by the tube V, provided with a
       valve, into the gas main U. The addition of fuel ought to proceed
       in such a way as to prevent the generated gas escaping; hence the
       space A is kept filled with the combustible material and covered
       with a lid.

  [20] An excess of air lowers the temperature of combustion, because it
       becomes heated itself, as explained in Chapter III. In ordinary
       furnaces the excess of air is three or four times greater than the
       quantity required for perfect combustion. In the best furnaces
       (with fire-bars, regulated air supply, and corresponding chimney
       draught) it is necessary to introduce twice as much air as is
       necessary, otherwise the smoke contains much carbonic oxide.

  [21] If in manufactories it is necessary, for instance, to maintain the
       temperature in a furnace at 1,000°, the flame passes out at this
       or a higher temperature, and therefore much fuel is lost in the
       smoke. For the draught of the chimney a temperature of 100° to
       150° is sufficient, and therefore the remaining heat ought to be
       utilised. For this purpose the flues are carried under boilers or
       other heating apparatus. The preparatory heating of the air is the
       best means of utilisation when a high temperature is desired
       (_see_ Note 22).

  [22] Regenerative furnaces were introduced by the Brothers Siemens
       about the year 1860 in many industries, and mark a most important
       progress in the use of fuel, especially in obtaining high
       temperatures. The principle is as follows: The products of
       combustion from the furnace are led into a chamber, I, and heat up
       the bricks in it, and then pass into the outlet flue; when the
       bricks are at a red heat the products of combustion are passed (by
       altering the valves) into another adjoining chamber, II, and air
       requisite for the combustion of the generator gases is passed
       through I. In passing round about the incandescent bricks the air
       is heated, and the bricks are cooled--that is, the heat of the
       smoke is returned into the furnace. The air is then passed through
       II, and the smoke through I. The regenerative burners for
       illuminating gas are founded on this same principle, the products
       of combustion heat the incoming air and gas, the temperature is
       higher, the light brighter, and an economy of gas is effected.
       Absolute perfection in these appliances has, of course, not yet
       been attained; further improvement is still possible, but
       dissociation imposes a limit because at a certain high temperature
       combinations do not ensue, possible temperatures being limited by
       reverse reactions. Here, as in a number of other cases, the
       further investigation of the matter must prove of direct value
       from a practical point of view.

  [23] At first sight it appears absurd, useless, and paradoxical to lose
       nearly one-third of the heat which fuel can develop, by turning it
       into gas. Actually the advantage is enormous, especially for
       producing high temperatures, as is already seen from the fact that
       fuels rich in oxygen (for instance, wood) when damp are unable,
       with any kind of hearth whatever, to give the temperature required
       for glass-melting or steel-casting, whilst in the gas-producer
       they furnish exactly the same gas as the driest and most
       carbonaceous fuel. In order to understand the principle which is
       here involved, it is sufficient to remember that a large amount of
       heat, but having a low temperature, is in many cases of no use
       whatever. We are unable here to enter into all the details of the
       complicated matter of the application of fuel, and further
       particulars must be sought for in special technical treatises. The
       following footnotes, however, contain certain fundamental figures
       for calculations concerning combustion.

The transformation of carbonic anhydride, by means of charcoal, into
carbonic oxide (C + CO_{2} = CO + CO) is considered a reversible
reaction, because at a high temperature the carbonic oxide splits up into
carbon and carbonic anhydride, as Sainte-Claire Deville showed by using
the method of the 'cold and hot tube.' Inside a tube heated in a furnace
another thin metallic (silvered copper) tube is fitted, through which a
constant stream of cold water flows. The carbonic oxide coming into
contact with the heated walls of the exterior tube forms charcoal, and
its minute particles settle in the form of lampblack on the lower side of
the cold tube, and, since they are cooled, do not act further on the
oxygen or carbonic anhydride formed.[24] A series of electric sparks
also decomposes carbonic oxide into carbonic anhydride and carbon, and if
the carbonic anhydride be removed by alkali complete decomposition may be
obtained (Deville).[24 bis] Aqueous vapour, which is so similar to
carbonic anhydride in many respects, acts, at a high temperature, on
charcoal in an exactly similar way, C + H_{2}O = H_{2} + CO. From 2
volumes of carbonic anhydride with charcoal 4 volumes of carbonic oxide
(2 molecules) are obtained, and precisely the same from 2 volumes of
water vapour with charcoal 4 volumes of a gas consisting of hydrogen and
carbonic oxide (H_{2} + CO) are formed. This mixture of combustible gases
is called _water gas_.[25] But aqueous vapour (and only when strongly
superheated, otherwise it cools the charcoal) only acts on charcoal to
form a large amount of carbonic oxide at a very high temperature (at
which carbonic anhydride dissociates); it begins to react at about 500°,
forming carbonic anhydride according to the equation C + 2H_{2}O = CO_{2}
+ 2H_{2}. Besides this, carbonic oxide on splitting up forms carbonic
anhydride, and therefore water gas always contains a mixture[26] in which
hydrogen predominates, the volume of carbonic oxide being comparatively
less, whilst the amount of carbonic anhydride increases as the
temperature of the reaction decreases (generally it is more than 3 per
cent.)

  [24] The first product of combustion of charcoal is always carbonic
       anhydride, and not carbonic oxide. This is seen from the fact that
       with a shallow layer of charcoal (less than a decimetre if the
       charcoal be closely packed) carbonic oxide is not formed at all.
       It is not even produced with a deep layer of charcoal if the
       temperature is not above 500°, and the current of air or oxygen is
       very slow. With a rapid current of air the charcoal becomes
       red-hot, and the temperature rises, and then carbonic oxide
       appears (Lang 1888). Ernst (1891) found that below 995° carbonic
       oxide is always accompanied by CO_{2}, and that the formation of
       CO_{2} begins about 400°. Naumann and Pistor determined that the
       reaction of carbonic anhydride with carbon commences at about
       550°, and that between water and carbon at about 500°. At the
       latter temperature carbonic anhydride is formed, and only with a
       rise of temperature is carbonic oxide formed (Lang) from the
       action of the carbonic anhydride on the carbon, and from the
       reaction CO_{2} + H_{2} = CO + H_{2}O. Rathke (1881) showed that
       at no temperature whatever is the reaction as expressed by the
       equation CO_{2} + C = 2CO_{2}, complete; a part of the carbonic
       anhydride remains, and Lang determined that at about 1,000° not
       less than 3 p.c. of the carbonic anhydride remains untransformed
       into carbonic oxide, even after the action has been continued for
       several hours. The endothermal reactions, C + 2H_{2}O = CO_{2}
       +2H_{2}, and CO + H_{2}O = CO_{2} + H_{2}, are just as incomplete.
       This is made clear if we note that on the one hand the
       above-mentioned reactions are all reversible, and therefore
       bounded by a limit; and, on the other hand, that at about 500°
       oxygen begins to combine with hydrogen and carbon, and also that
       the lower limits of dissociation of water, carbonic anhydride, and
       carbonic oxide lie near one another between 500° and 1,200°. For
       water and carbonic oxide the lower limit of the commencement of
       dissociation is unknown, but judging from the published data
       (according to Le Chatelier, 1888) that of carbonic anhydride may
       be taken as about 1,050°. Even at about 200° half the carbonic
       anhydride dissociates if the pressure be small, about 0·001
       atmosphere. At the atmospheric pressure, not more than 0·05 p.c.
       of the carbonic anhydride decomposes. The reason of the influence
       of pressure is here evidently that the splitting up of carbonic
       anhydride into carbonic oxide and oxygen is accompanied by an
       increase in volume (as in the case of the dissociation of nitric
       peroxide. _See_ Chapter VI., Note 46). As in stoves and lamps, and
       also with explosive substances, the temperature is not higher than
       2,000° to 2,500°, it is evident that although the partial pressure
       of carbonic anhydride is small, still its dissociation cannot here
       be considerable, and probably does not exceed 5 p.c.

  [24 bis] Besides which L. Mond (1890) showed that the powder of freshly
       reduced metallic nickel (obtained by heating the oxide to redness
       in a stream of hydrogen) is able, when heated even to 350°, to
       completely decompose carbonic oxide into CO_{2} and carbon, which
       remains with the nickel and is easily removed from it by heating
       in a stream of air. Here 2CO = CO_{2} + C. It should be remarked
       that heat is evolved in this reaction (Note 25), and therefore
       that the influence of 'contact' may here play a part. Indeed, this
       reaction must be classed among the most remarkable instances of
       the influence of contact, especially as metals analogous to Ni (Fe
       and Co) do not effect this reaction (_see_ Chapter II., Note 17).

  [25] A molecular weight of this gas, or 2 volumes CO (28 grams), on
       combustion (forming CO_{2}) gives out 68,000 heat units (Thomsen
       67,960 calories). A molecular weight of hydrogen, H_{2} (or 2
       volumes), develops on burning into _liquid_ water 69,000 heat
       units (according to Thomsen 68,300), but if it forms aqueous
       vapour 58,000 heat units. Charcoal, resolving itself by combustion
       into the molecular quantity of CO_{2} (2 volumes), develops 97,000
       heat units. From the data furnished by these exothermal reactions
       it follows: (1) that the oxidation of charcoal into carbonic oxide
       develops 29,000 heat units; (2) that the reaction C + CO_{2} = 2CO
       _absorbs_ 39,000 heat units; (3) C + H_{2}O = H_{2} + CO _absorbs_
       (if the water be in a state of vapour) 29,000 calories, but if the
       water be liquid 40,000 calories (almost as much as C + CO_{2});
       (4) C + H_{2}O = CO_{2} + 2H_{2} _absorbs_ (if the water be in a
       state of vapour) 19,000 heat units; (5) the reaction CO + H_{2}O =
       CO_{2} + H_{2} _develops_ 10,000 heat units if the water be in the
       state of vapour; and (6) the decomposition expressed by the
       equation 2CO = C + CO_{2} (Note 24 bis) is accompanied by the
       _evolution_ of 39,000 units of heat.

       Hence it follows that 2 volumes of CO or H_{2} burning into CO_{2}
       or H_{2}O develop almost the same amount of heat, just as also the
       heat effects corresponding with the equations

                         C + H_{2}O = CO + H_{2}

                          C + CO_{2} = CO + CO

       are nearly equal.

  [26] _Water gas_, obtained from steam and charcoal at a white heat,
       contains about 50 p.c. of hydrogen, about 40 p.c. of carbonic
       oxide, about 5 p.c. of carbonic anhydride, the remainder being
       nitrogen from the charcoal and air. Compared with producer gas,
       which contains much nitrogen, this is a gas much richer in
       combustible matter, and therefore capable of giving high
       temperatures, and is for this reason of the greatest utility. If
       carbonic anhydride could be as readily obtained in as pure a state
       as water, then CO might be prepared directly from CO_{2} + C, and
       in that case the utilisation of the heat of the carbon would be
       the same as in water gas, because CO evolves as much heat as
       H_{2}, and even more if the temperature of the smoke be over 100°,
       and the water remains in the form of vapour (Note 25). But
       producer gas contains a large proportion of nitrogen, so that its
       effective temperature is below that given by water gas; therefore
       in places where a particularly high temperature is required (for
       instance, for lighting by means of incandescent lime or magnesia,
       or for steel melting, &c.), and where the gas can be easily
       distributed through pipes, water gas is at present held in high
       estimation, but when (in ordinary furnaces, re-heating,
       glass-melting, and other furnaces) a very high temperature is not
       required, and there is no need to convey the gas in pipes,
       producer gas is generally preferred on account of the simplicity
       of its preparation, especially as for water gas such a high
       temperature is required that the plant soon becomes damaged.

       There are numerous systems for making water gas, but the American
       patent of T. Lowe is generally used. The gas is prepared in a
       cylindrical generator, into which hot air is introduced, in order
       to raise the coke in it to a white heat. The products of
       combustion containing carbonic oxide are utilised for superheating
       steam, which is then passed over the white hot coke. Water gas, or
       a mixture of hydrogen and carbonic oxide, is thus obtained.

       Water gas is sometimes called '_the fuel of the future_,' because
       it is applicable to all purposes, develops a high temperature, and
       is therefore available, not only for domestic and industrial uses,
       but also for gas-motors and for lighting. For the latter purpose
       platinum, lime, magnesia, zirconia, and similar substances (as in
       the Drummond light, Chapter III.), are rendered incandescent in
       the flame, or else the gas is _carburetted_--that is, mixed with
       the vapours of volatile hydrocarbons (generally benzene or
       naphtha, naphthalene, or simply naphtha gas), which communicate to
       the pale flame of carbonic oxide and hydrogen a great brilliancy,
       owing to the high temperature developed by the combustion of the
       non-luminous gases. As water gas, possessing these properties, may
       be prepared at central works and conveyed in pipes to the
       consumers, and as it may be produced from any kind of fuel, and
       ought to be much cheaper than ordinary gas, it may as a matter of
       fact be expected that in course of time (when experience shall
       have determined the cheapest and best way to prepare it) it will
       not only supplant ordinary gas, but will with advantage everywhere
       replace the ordinary forms of fuel, which in many respects are
       inconvenient. At present its consumption spreads principally for
       lighting purposes, and for use in gas-engines instead of ordinary
       illuminating gas. In some cases Dowson gas is prepared in
       producers. This is a mixture of water and producer gases obtained
       by passing steam into an ordinary producer (Note 19), when the
       temperature of the carbon has become sufficiently high for the
       reaction C + H_{2}O = CO + H_{2}.

Metals like iron and zinc which at a red heat are capable of decomposing
water with the formation of hydrogen, also decompose carbonic anhydride
with the formation of carbonic oxide; so both the ordinary products of
complete combustion, water and carbonic anhydride, are very similar in
their reactions, and we shall therefore presently compare hydrogen and
carbonic oxide. The metallic oxides of the above-mentioned metals, when
reduced by charcoal, also give carbonic oxide. Priestley obtained it by
heating charcoal with zinc oxide. As free carbonic anhydride may be
transformed into carbonic oxide, so, in precisely the same way, may that
carbonic acid which is in a state of combination; hence, if magnesium or
barium carbonates (MgCO_{3} or BaCO_{3}) be heated to redness with
charcoal, or iron or zinc, carbonic oxide will be produced--for instance,
it is obtained by heating an intimate mixture of 9 parts of chalk and 1
part of charcoal in a clay retort.

Many organic substances[27] on being heated, or under the action of
various agents, yield carbonic oxide; amongst these are many organic
or carboxylic acids. The simplest are formic and oxalic acids. Formic
acid, CH_{2}O_{2}, on being heated to 200°, easily decomposes into
carbonic oxide and water, CH_{2}O_{2} = CO + H_{2}O.[27 bis] Usually,
however, carbonic oxide is prepared in laboratories, not from formic
but from oxalic acid, C_{2}H_{2}O_{4}, the more so as formic acid is
itself prepared from oxalic acid. The latter acid is easily obtained by
the action of nitric acid on starch, sugar, &c.; it is also found in
nature. Oxalic acid is easily decomposed by heat; its crystals first
lose water, then partly volatilise, but the greater part is decomposed.
The decomposition is of the following nature: it splits up into water,
carbonic oxide, and carbonic anhydride,[28] C_{2}H_{2}O_{4} = H_{2}O
+ CO_{2} + CO. This decomposition is generally practically effected
by mixing oxalic acid with strong sulphuric acid, because the latter
assists the decomposition by taking up the water. On heating a mixture
of oxalic and sulphuric acids a mixture of carbonic oxide and carbonic
anhydride is evolved. This mixture is passed through a solution of an
alkali in order to absorb the carbonic anhydride, whilst the carbonic
oxide passes on.[28 bis]

  [27] The so-called yellow prussiate, K_{4}FeC_{6}N_{6}, on being heated
       with ten parts of strong sulphuric acid forms a considerable
       quantity of very pure carbonic oxide quite free from carbonic
       anhydride.

  [27 bis] To perform this reaction, the formic acid is mixed with
       glycerine, because when heated alone it volatilises much below its
       temperature of decomposition. When heated with sulphuric acid the
       salts of formic acid yield carbonic oxide.

  [28] The decomposition of formic and oxalic acids, with the formation
       of carbonic oxide, considering these acids as carboxyl
       derivatives, may be explained as follows:--The first is H(COOH)
       and the second (COOH)_{2}, or H_{2} in which one or both halves of
       the hydrogen are exchanged for carboxyl; therefore they are equal
       to H_{2} + CO_{2} and H_{2} + 2CO_{2}; but H_{2} reacts with
       CO_{2}, as has been stated above, forming CO and H_{2}O. From this
       it is also evident that oxalic acid on losing CO_{2} forms formic
       acid, and also that the latter may proceed from CO + H_{2}O, as we
       shall see further on.

  [28 bis] Greshoff (1888) showed that with a solution of nitrate of
       silver, iodoform, CHI_{3}, forms CO according to the equation
       CHI_{3} + 3AgNO_{3} + H_{2}O = 3AgI + 3HNO_{3} + CO. The reaction
       is immediate and is complete.

In its physical _properties_ carbonic oxide resembles nitrogen; this is
explained by the equality of their molecular weights. The absence of
colour and smell, the low temperature of the absolute boiling point,
-140° (nitrogen, -146°), the property of solidifying at -200°
(nitrogen, -202°), the boiling point of -190° (nitrogen, -203°), and the
slight solubility (Chapter I., Note 30), of carbonic oxide are almost the
same as in those of nitrogen. The chemical properties of both gases are,
however, very different, and in these carbonic oxide resembles hydrogen.
Carbonic oxide burns with a blue flame, giving 2 volumes of carbonic
anhydride from 2 volumes of carbonic oxide, just as 2 volumes of hydrogen
give 2 volumes of aqueous vapour. It explodes with oxygen, in the
eudiometer, like hydrogen.[29] When breathed it acts as a strong poison,
being absorbed by the blood;[30] this explains the action of charcoal
fumes, the products of the incomplete combustion of charcoal and other
carbonaceous fuels. Owing to its faculty of combining with oxygen,
carbonic oxide acts as a powerful reducing agent, taking up the oxygen
from many compounds at a red heat, and being itself transformed into
carbonic anhydride. The reducing action of carbonic oxide, however, is
(like that of hydrogen, Chapter II.) naturally confined to those oxides
which easily part with their oxygen--as, for instance, copper
oxide--whilst the oxides of magnesium or potassium are not reduced.
Metallic iron itself is capable of reducing carbonic anhydride to
carbonic oxide, just as it liberates the hydrogen from water. Copper,
which does not decompose water, does not decompose carbonic oxide. If a
platinum wire heated to 300°, or spongy platinum at the ordinary
temperature, be plunged into a mixture of carbonic oxide and oxygen, or
of hydrogen and oxygen, the mixture explodes. These reactions are very
similar to those peculiar to hydrogen. The following important
distinction, however, exists between them--namely: the molecule of
hydrogen is composed of H_{2}, a group of elements divisible into two
like parts, whilst, as the molecule of carbonic oxide, CO, contains
unlike atoms of carbon and oxygen, in none of its reactions of
combination can it give two molecules of matter containing its elements.
This is particularly noticeable in the action of chlorine on hydrogen and
on carbonic oxide respectively; with the former chlorine forms hydrogen
chloride, and with the latter it produces the so-called carbonyl
chloride, COCl_{2}: that is to say, the molecule of hydrogen, H_{2},
under the action of chlorine divides, forming two molecules of
hydrochloric acid, whilst the molecule of carbonic oxide enters in its
entirety into the molecule of carbonyl chloride. This characterises the
so-called _diatomic_ or _bivalent_ reactions of radicles or _residues_.
H is a monatomic residue or radicle, like K, Cl, and others, whilst
carbonic oxide, CO, is an indivisible (undecomposable) bivalent radicle,
equivalent to H_{2} and not to H, and therefore combining with X_{2} and
interchangeable with H_{2}. This distinction is evident from the annexed
comparison:

        HH, hydrogen.            CO, carbonic oxide.
        HCl, hydrochloric acid.  COCl_{2}, carbonyl chloride.
        HKO, potash.             CO(KO)_{2}, potassium carbonate.
        HNH_{2}, ammonia.        CO(NH_{2})_{2}, urea.
        HCH_{3}, methane.        CO(CH_{3})_{2}, acetone.
        HHO, water.              CO(HO)_{2}, carbonic acid.

  [29] It is remarkable that, according to the investigations of Dixon,
       perfectly dry carbonic oxide does not explode with oxygen when a
       spark of low intensity is used, but an explosion takes place if
       there is the slightest admixture of moisture. L. Meyer, however,
       showed that sparks of an electric discharge of considerable
       intensity produce an explosion. N. N. Beketoff demonstrated that
       combustion proceeds and spreads slowly unless there be perfect
       dryness. I think that this may he explained by the fact that water
       with carbonic oxide gives carbonic anhydride and hydrogen, but
       hydrogen with oxygen gives hydrogen peroxide (Chapter VII.), which
       with carbonic oxide forms carbonic anhydride and water. The water,
       therefore, is renewed, and again serves the same purpose. But it
       may be that here it is necessary to acknowledge a simple contact
       influence. After Dixon had shown the influence of traces of
       moisture upon the reaction CO + O, many researches were made of a
       similar nature. The fullest investigation into the influence of
       moisture upon the course of many chemical reactions was made by
       Baker in 1894. He showed that with perfect dryness, many chemical
       transformations (for example, the formation of ozone from oxygen,
       the decomposition of AgO, KClO_{3} under the action of heat, &c.)
       proceeds in exactly the same manner as in the presence of
       moisture; but that in many cases traces of moisture have an
       evident influence. We may mention the following instances: (1) Dry
       SO_{3} does not act upon dry CaO or CuO; (2) perfectly dry
       sal-ammoniac does not give NH_{3} with dry CaO, but simply
       volatilises; (3) dry NO and O do not react; (4) perfectly dry
       NH_{3} and HCl do not combine; (5) perfectly dry sal-ammoniac does
       not dissociate at 350° (Chapter VII., Note 15 bis); and (6)
       perfectly dry chlorine does not act upon metals, &c.

  [30] Carbonic oxide is very rapid in its action, because it is absorbed
       by the blood in the same way as oxygen. In addition to this, the
       absorption spectrum of the blood changes so that by the help of
       blood it is easy to detect the slightest traces of carbonic oxide
       in the air. M. A. Kapoustin found that linseed oil and therefore
       oil paints, are capable of giving off carbonic oxide while drying
       (absorbing oxygen).

Such monatmic (univalent) residues, X, as H, Cl, Na, NO_{2}, NH_{4},
CH_{3}, CO_{2}H (carboxyl), OH, and others, in accordance with
the law of substitution, combine together, forming compounds, XX';
and with oxygen, or in general with diatomic (bivalent) residues,
Y--for instance, O, CO, CH_{2}, S, Ca, &c. forming compounds XX´Y; but
diatomic residues, Y, sometimes capable of existing separately may
combine together, forming YY´ and with X_{2} or XX´, as we see from
the transition of CO into CO_{2} and COCl_{2}. This combining power of
carbonic oxide appears in many of its reactions. Thus it is very easily
absorbed by cuprous chloride, CuCl, dissolved in fuming hydrochloric
acid, forming a crystalline compound, COCu_{2}Cl_{2},2H_{2}O,
decomposable by water; it combines directly with potassium (at 90°),
forming (KCO)_{_n_}[31] with platinum dichloride, PtCl_{2}, with
chlorine, Cl_{2}, &c.

  [31] The molecule of metallic potassium (Scott, 1887), like that of
       mercury, contains only one atom, and it is probably in virtue of
       this that the molecules CO and K combine together. But as in the
       majority of cases potassium acts as a univalent radicle, the
       polymeride K_{2}C_{2}O_{2} is formed, and probably
       K_{10}C_{10}O_{10}, because products containing C_{10} are formed
       by the action of hydrochloric acid. The black mass formed by the
       combination of carbonic oxide with potassium explodes with great
       ease, and oxidises in the air. Although Brodie, Lerch, and Joannis
       (who obtained it in 1873 in a colourless form by means of NH_{3}K,
       described in Chapter VI., Note 14) have greatly extended our
       knowledge of this compound, much still remains unexplained. It
       probably exists in various polymeric and isomeric forms, having
       the composition (KCO)_{_n_} and (NaCO)_{_n_}.

But the most remarkable compounds are (1) the compound of CO with
metallic nickel, a colourless volatile liquid, Ni(CO)_{4}, obtained by
L. Mond (described in Chapter XXII.) and (2) the compounds of carbonic
oxide with the alkalis, for instance with potassium or barium hydroxide,
&c.--although it is not directly absorbed by them, as it has no acid
properties. Berthelot (1861) showed that potash in the presence of water
is capable of absorbing carbonic oxide, but the absorption takes place
slowly, little by little, and it is only after being heated for many
hours that the whole of the carbonic oxide is absorbed by the potash. The
salt CHKO_{2} is obtained by this absorption; it corresponds with an acid
found in nature--namely, the simplest organic (carboxylic) acid, _formic
acid_, CH_{2}O_{2}. It can be extracted from the potassium salt by means
of distillation with dilute sulphuric acid, just as nitric acid is
prepared from sodium nitrate. The same acid is found in ants and in
nettles (when the stings of the nettles puncture the skin they break, and
the corrosive formic acid enters into the body); it is also obtained
during the action of oxidising agents on many organic substances; it is
formed from oxalic acid, and under many conditions splits up into
carbonic oxide and water. In the formation of formic acid from carbonic
oxide we observe an example of the synthesis of organic compounds, such
as are now very numerous, and are treated of in detail in works on
organic chemistry.

Formic acid, H(CHO_{2}), carbonic acid, HO(CHO_{2}), and oxalic acid,
(CHO_{2})_{2}, are the simple organic or carboxylic acids, R(CHO_{2})
corresponding with HH and HOH. Commencing with carbonic oxide, CO, the
formation of carboxylic acids is clearly seen from the fact that CO is
capable of combining with X_{2}, that is of forming COX_{2}. If, for
instance, one X is an aqueous residue, OH (hydroxyl), and the other X
is hydrogen, then the simplest organic acid--formic acid, H(COOH)--is
obtained. As all hydrocarbons (Chapter VIII.) correspond with the
simplest, CH_{4}, so all organic acids may be considered to proceed from
formic acid.

In a similar way it is easy to explain the relation to other compounds
of carbon of those compounds which contain nitrogen. By way of an
example, we will take one of the carboxyl acids, R(CO_{2}H), where R is
a hydrocarbon radicle (residue). Such an acid, like all others, will
give by combination with NH_{3} an ammoniacal salt, R(CO_{2}NH_{4}).
This salt contains the elements for the formation of two molecules of
water, and under suitable conditions by the action of bodies capable
of taking it up, water may in fact be separated from R(CO_{2}NH_{4}),
forming by the loss of one molecule of water, _amides_, RCONH_{2}, and
by the loss of two molecules of water, _nitriles_, RCN, otherwise known
as _cyanogen compounds_ or _cyanides_.[32] If all the carboxyl acids are
united not only by many common reactions but also by a mutual conversion
into each other (an instance of which we saw above in the conversion
of oxalic acid into formic and carbonic acids) one would expect the
same for all the cyanogen compounds also. The common character of
their reactions, and the reciprocity of their transformation, were
long ago observed by Gay-Lussac, who recognised a common group or
radicle (residue) cyanogen, CN, in all of them. The simplest compounds
are _hydrocyanic_ or _prussic acid_, HCN, cyanic acid, OHCN, and free
cyanogen, (CN)_{2}, which correspond to the three simplest carboxyl
acids: formic, HCO_{2}H, carbonic, OHCO_{2}H, and oxalic, (CO_{2}H)_{2}.
Cyanogen, like carboxyl, is evidently a monatomic residue and acid,
similar to chlorine. As regards the amides RCONH_{2}, corresponding
to the carboxyl acids, they contain the ammoniacal residue NH_{2},
and form a numerous class of organic compounds met with in nature and
obtained in many ways,[33] but not distinguished by such characteristic
peculiarities as the cyanogen compounds.

  [32] The connection of the cyanogen compounds with the rest of the
       hydrocarbons by means of carboxyl was enunciated by me, about the
       year 1860, at the first Annual Meeting of the Russian Naturalists.

  [33] Thus, for instance, _oxamide_, or the amide of oxalic acid,
       (CNH_{2}O)_{2}, is obtained in the form of an insoluble
       precipitate on adding a solution of ammonia to an alcoholic
       solution of ethyl oxalate, (CO_{2}C_{2}H_{5})_{2}, which is formed
       by the action of oxalic acid on alcohol: (CHO_{2})_{2} +
       2(C_{2}H_{5})OH = 2HOH + (CO_{2}C_{2}H_{5})_{2}. As the nearest
       derivatives of ammonia, the amides treated with alkalis yield
       ammonia and form the salt of the acid. The nitriles do not,
       however, give similar reactions so readily. The majority of amides
       corresponding to acids have a composition RNH_{2}, and therefore
       recombine with water with great ease even when simply boiled with
       it, and with still greater facility in presence of acids or
       alkalis. Under the action of alkalis the amides naturally give off
       ammonia, through the combination of water with the amide, when a
       salt of the acid from which the amide was derived is formed:
       RNH_{2} + KHO = RKO + NH_{3}.

       The same reaction takes place with acids, only an ammoniacal salt
       of the acid is of course formed whilst the acid held in the amide
       is liberated: RNH_{2} + HCl + H_{2}O = RHO + NH_{4}Cl.

       Thus in the majority of cases amides easily pass into ammoniacal
       salts, but they differ essentially from them. No ammoniacal salt
       sublimes or volatilises unchanged, and generally when heated it
       gives off water and yields an amide, whilst many amides volatilise
       without alteration and frequently are volatile crystalline
       substances which may be easily sublimed. Such, for instance, are
       the amides of benzoic, formic, and many other organic acids.

The reactions and properties of the amides and nitriles of the organic
acids are described in detail in books on organic chemistry; we will here
only touch upon the simplest of them, and to clearly explain the
derivative compounds will first consider the ammoniacal salts and amides
of carbonic acid.

As carbonic acid is bibasic, its ammonium salts ought to have the
following composition: _acid carbonate of ammonium_, H(NH_{4})CO_{3},
and _normal carbonate_, (NH_{4})_{2}CO_{3}; they represent compounds
of one or two molecules of ammonia with carbonic acid. The acid salt
appears in the form of a non-odoriferous and (when tested with litmus)
neutral substance, soluble at the ordinary temperature in six parts
of water, insoluble in alcohol, and obtainable in a crystalline form
either without water of crystallisation or with various proportions
of it. If an aqueous solution of ammonia be saturated with an excess
of carbonic anhydride, and then evaporated over sulphuric acid in
the bell jar of an air-pump, crystals of this salt are separated.
Solutions of all other ammonium carbonates, when evaporated under
the air-pump, yield crystals of this salt. A solution of this salt,
even at the ordinary temperature, gives off carbonic anhydride, as do
all the acid salts of carbonic acid (for instance, NaHCO_{3}), and
at 38° the separation of carbonic anhydride takes place with great
rapidity. _On losing carbonic anhydride_ and water, the acid salt is
converted into the normal salt, 2(NH_{4})HCO_{3} = H_{2}O + CO_{2} +
(NH_{4})2CO_{3}; the latter, however, decomposes in solution, and can
therefore only be obtained in crystals, (NH_{4})_{2}CO_{3},H_{2}O, at
low temperatures, and from solutions containing _an excess of ammonia_
as the product of dissociation of this salt: (NH_{4})_{2}CO_{3} = NH_{3}
+ (NH_{4})HCO_{3}. But the normal salt,[34] according to the general
type, is capable of decomposing _with separation of water_, and forming
_ammonium carbamate_, NH_{4}O(CONH_{2}) = (NH_{4})_{2}CO_{3}-H_{2}O;
this still further complicates the chemical transformations of the
carbonates of ammonium. It is in fact evident that, by changing the
ratios of water, ammonia, and carbonic acid, various intermediate salts
will be formed containing mixtures or combinations of those mentioned
above. Thus the ordinary commercial _carbonate of ammonia_ is obtained
by heating a mixture of chalk and sulphate of ammonia (Chapter VI.),
or sal-ammoniac, 2NH_{4}Cl + CaCO_{3} = CaCl_{2} + (NH_{4})_{2}CO_{3}.
The normal salt, however, through loss of part of the ammonia, partly
forms the acid salt, and, partly through loss of water, forms carbamate,
and most frequently presents the composition NH_{4}O(CONH_{2}) +
2OH(CO_{2}NH_{4}) = 4NH_{3} + 3CO_{2} + 2H_{2}O. This salt, in parting
under various conditions with ammonia, carbonic anhydride, and water,
does not present a constant composition, and ought rather to be regarded
as a mixture of acid salt and amide salt. The latter must be recognised
as entering into the composition of the ordinary carbonate of ammonia,
because it contains less water than is required for the normal or acid
salt;[35] but on being dissolved in water this salt gives a mixture of
acid and normal salts.

  [34] The acid salt, (NH_{4})HCO_{3}, on losing water ought to form the
       _carbamic acid_, OH(CNH_{2}O); but it is not formed, which is
       accounted for by the instability of the acid salt itself. Carbonic
       anhydride is given off and ammonia is produced, which gives
       ammonium carbamate.

  [35] In the normal salt, 2NH_{3} + CO_{2} + H_{2}O, in the acid salt,
       NH_{3} + CO_{2} + H_{2}O, but in the commercial salt only 2H_{2}O
       to 3CO_{2}.

Each of the two ammoniacal salts of carbonic acid has its corresponding
amide. That of the acid salt should be acid, if the water given off takes
up the hydrogen of the ammonia, as it should according to the common type
of formation of the amides, so that OHCONH_{2}, or _carbamic acid_, is
formed from OHCO_{3}NH_{4}. This acid is not known in a free state, but
its corresponding ammoniacal salt or _ammonium carbamate_ is known. The
latter is easily and immediately formed by mixing 2 volumes of _dry_
ammonia with 1 volume of dry carbonic anhydride, 2NH_{3} + CO_{2} =
NH_{4}O(CONH_{2}); it is a solid substance, smells strongly of ammonia,
attracts moisture from the air, and decomposes completely at 60°. The
fact of this decomposition may be proved[36] by the density of its
vapour, which = 13 (H = 1); this exactly corresponds with the density of
a mixture of 2 volumes of ammonia and 1 volume of carbonic anhydride. It
is easily understood that such a combination will take place with any
ammonium carbonate under the action of salts which take up the water--for
instance, sodium or potassium carbonate[37]--as in an anhydrous state
ammonia and carbonic anhydride only form one compound, CO_{2}2NH_{3}.[38]
As the normal ammonium carbonate contains two ammonias, and as the amides
are formed with the separation of water at the expense of the hydrogen of
the ammonias, so this salt has its symmetrical amide, CO(NH_{2})_{2}.
This must be termed carbamide. It is identical with urea, CN_{2}H_{4}O,
which, contained in the urine (about 2 per cent. in human urine), is for
the higher animals (especially the carnivorous) the ordinary product of
excretion[39] and oxidation of the nitrogenous substances found in the
organism. If ammonium carbamate be heated to 140° (in a sealed tube,
Bazaroff), or if carbonyl chloride, COCl_{2}, be treated with ammonia
(Natanson), urea will be obtained, which shows its direct connection with
carbonic acid--that is, the presence of carbonic acid and ammonia in it.
From this it will be understood how urea during the putrefaction of urine
is converted into ammonium carbonate, CN_{2}H_{4}O + H_{2}O = CO_{2} +
2NH_{3}.

  [36] Naumann determined the following dissociation tensions of
       the vapour of ammonium carbamate (in millimetres of mercury):--

                 -10°  0°  +10°  20°  30°  40°  50°  60°
                   5   12   30   62   124  248  470  770

       Horstmann and Isambert studied the tensions corresponding to
       excess of NH_{3} or CO_{2}, and found, as might have been
       expected, that with such excess the mass of the salt formed (in a
       solid state) increases and the decomposition (transition into
       vapour) decreases.

  [37] Calcium chloride enters into double decomposition with ammonium
       carbamate. Acids (for instance, sulphuric) take up ammonia, and
       set free carbonic anhydride; whilst alkalis (such as potash) take
       up carbonic anhydride and set free ammonia, and therefore, in this
       case for removing water only sodium or potassium carbonate can be
       taken. An aqueous solution of ammonium carbamate does not entirely
       precipitate a solution of CaCl_{2}, probably because calcium
       carbamate is soluble in water, and all the (NH_{3})_{2}CO_{2} is
       not converted by dissolving into the normal salt,
       (NH_{4}O)_{2}CO_{3}.

  [38] It must be imagined that the reaction takes place at first between
       equal volumes (Chapter VII.); but then carbamic acid,
       HO(CNH_{2}O), is produced, which, as an acid, immediately combines
       with the ammonia, forming NH_{4}O(CNH_{2}O).

  [39] Urea is undoubtedly a product of the oxidation of complex
       nitrogenous matters (albumin) of the animal body. It is found in
       the blood. It is absorbed from the blood by the kidneys. A man
       excretes about 30 grams of urea per day. As a derivative of
       carbonic anhydride, into which it is readily converted, urea is in
       a sense a product of oxidation.

Thus urea, both by its origin and decomposition, is an amide of carbonic
acid. Representing as it does ammonia (two molecules) in which hydrogen
(two atoms) is replaced by the bivalent radicle of carbonic acid, urea
retains the property of ammonia of entering into combination, with acids
(thus nitric acid forms CN_{2}H_{4}O,HNO_{3}), with bases (for instance,
with mercury oxide), and with salts (such as sodium chloride, ammonium
chloride), but containing an acid residue it has no alkaline properties.
It is soluble in water without change, but at a red heat loses ammonia
and forms _cyanic acid_, CNHO,[39 bis] which is a nitrile of carbonic
acid--that is to say, is a cyanogen compound, corresponding to the acid
ammonium carbonate, OH(CNH_{4}O_{2}), which on parting with 2H_{2}O ought
to form cyanic acid, CNOH. Liquid cyanic acid, exceedingly unstable at
the ordinary temperatures, gives its stable solid polymer cyanuric acid,
O_{3}H_{3}C_{3}N_{3}. Both have the same composition, and they pass one
into another at different temperatures. If crystals of cyanuric acid be
heated to a temperature, _t_°, then the vapour tension, _p_, in
millimetres of mercury (Troost and Hautefeuille) will be:

             _t._  160°,  170°,  200°,  250°,  300°,   350°
             _p._   56,    68,   130,   220,   430,  1,200

The vapour contains cyanic acid, and, if it be rapidly cooled, it
condenses into a mobile volatile liquid (specific gravity at 0° =
1·14). If the liquid cyanic acid be gradually heated, it passes into
a new amorphous polymeride (cyamelide), which, on being heated, like
cyanuric acid, forms vapours of cyanic acid. If these fumes are heated
above 150° they pass directly into cyanuric acid. Thus at a temperature
of 350°, the pressure does not rise above 1,200 mm. on the addition of
vapours of cyanic acid, because the whole excess is transformed into
cyanuric acid. Hence, the above-mentioned figures give the tension
of dissociation of cyanuric acid, or the greatest pressure which the
vapours of HOCN are able to attain at a given temperature, whilst at
a greater pressure, or by the introduction of a larger mass of the
substance into a given volume, the whole of the excess is converted into
cyanuric acid. The properties of cyanic acid which we have described
were principally observed by Wöhler, and clearly show the _faculty of
polymerisation of cyanogen compounds_. This is observed in many other
cyanogen derivatives, and is to be regarded as the consequence of the
above-mentioned explanation of their nature. All cyanogen compounds are
ammonium salts, R(CNH_{4}O_{2}), deprived of water, 2H_{2}O; therefore
the molecules, RCN, ought to possess the faculty of combining with two
molecules of water or with other molecules in exchange for it (for
instance, with H_{2}S, or HCl, or 2H_{2}, &c.), and are therefore
capable of combining together. The combination of molecules of the same
kind to form more complex ones is what is meant by polymerisation.[40]

  [39 bis] Its polymer, C_{3}N_{3}H_{3}O_{3}, is formed together
       with it. Cyanic acid is a very unstable, easily changeable liquid,
       while cyanuric acid is a crystalline solid which is very stable at
       the ordinary temperature.

  [40] Just as the aldehydes (such as C_{2}H_{4}O) are alcohols (like
       C_{2}H_{6}O) which have lost hydrogen and are also capable of
       entering into combination with many substances, and of
       polymerising, forming slightly volatile polymerides, which
       depolymerise on heating. Although there are also many similar
       phenomena (for instance, the transformation of yellow into red
       phosphorus, the transition of cinnamene into metacinnamene, &c.)
       of polymerisation, in no other case are they so clearly and simply
       expressed as in cyanic acid. The details relating to this must be
       sought for in treatises on organic and theoretical chemistry. If
       we touch on certain sides of this question it is principally with
       the view of showing the phenomenon of polymerisation by typical
       examples, for it is of more frequent occurrence than was formerly
       supposed among compounds of several elements.

Besidea being a substance very prone to form polymerides, cyanic acid
presents many other features of interest, expounded in greater detail in
organic chemistry. However we may mention here the production of the
cyanates by the oxidation of the metallic cyanides. Potassium cyanate,
KCNO, is most often obtained in this way. Solutions of cyanates by the
addition of sulphuric acid yield cyanic acid, which, however, immediately
decomposes: CNHO + H_{2}O = CO_{2} + NH_{3}. A solution of ammonium
cyanate, CN(NH_{4})O, behaves in the same manner, but only in the cold.
On being heated it completely changes because it is transformed into
urea. The composition of both substances is identical, CN_{2}H_{4}O, but
the structure, or disposition of, and connection between, the elements is
different: in the ammonium cyanate one atom of nitrogen exists in the
form of cyanogen, CN--that is, united with carbon--and the other as
ammonium, NH_{4}, but, as cyanic acid contains the hydroxyl radicle of
carbonic acid, OH(CN), the ammonium in this salt is united with oxygen.
The composition of this salt is best expressed by supposing one atom of
the hydrogen in water to be replaced by ammonium and the other by
cyanogen--_i.e._ that its composition is not symmetrical--whilst in urea
both the nitrogen atoms are symmetrically and uniformly disposed as
regards the radicle CO of carbonic acid: CO(NH_{2})_{2}. For this reason,
urea is much more stable than ammonium cyanate, and therefore the latter,
on being slightly heated in solution, is converted into urea. This
remarkable isomeric transformation was discovered by Wöhler in 1828.[41]
Formamide, HCONH_{2}, and _hydrocyanic acid_, HCN, as a nitrile,
correspond with formic acid, HCOOH, and therefore ammonium formate,
HCOONH_{4}, and formamide, when acted on by heat and by substances which
take up water (phosphoric anhydride) form hydrocyanic acid, HCN, whilst,
under many conditions (for instance, on combining with hydrochloric acid
in presence of water), this hydrocyanic acid forms formic acid and
ammonia. Although containing hydrogen in the presence of two acid-forming
elements--namely, carbon and nitrogen[42]--hydrocyanic acid does not
give an acid reaction with litmus (cyanic acid has very marked acid
properties); _but it forms salts_, _MCN_, thus presenting the properties
of a feeble acid, and for this reason is called an _acid_. The small
amount of energy which it has is shown by the fact that the cyanides of
the alkali metals--for instance, potassium cyanide (KHO + HCN = H_{2}0 +
KCN) in solution have a strongly alkaline reaction.[43] If ammonia be
passed over charcoal at a red heat, especially in the presence of an
alkali, or if gaseous nitrogen be passed through a mixture of charcoal
and an alkali (especially potash, KHO), and also if a mixture of
nitrogenous organic substances and alkali be heated to a red heat, in all
these cases the alkali metal combines with the carbon and nitrogen,
forming a metallic cyanide, MCN--for example, KCN.[43 bis] Potassium
cyanide is much used in the arts, and is obtained, as above stated, under
many circumstances--as, for instance, in iron smelting, especially with
the assistance of wood charcoal, the ash of which contains much potash.
The nitrogen of the air, the alkali of the ash, and the charcoal are
brought into contact at a high temperature during iron smelting, and
therefore, under these conditions, a considerable quantity of potassium
cyanide is formed. In practice it is not usual to prepare potassium
cyanide directly, but a peculiar compound of it containing potassium,
iron, and cyanogen. This compound is potassium ferrocyanide, and is also
known as _yellow prussiate of potash_. This saline substance (_see_
Chapter XXII) has the composition K_{4}FeC_{6}N_{6} + 2H_{2}O. The name
of cyanogen ([Greek: kuanos]) is derived from the property which this
yellow prussiate possesses of forming, with a solution of a ferric salt,
FeX_{3}, the familiar pigment Prussian blue. The yellow prussiate is
manufactured on a large scale, and is generally used as the source of
the other cyanogen compounds.

  [41] It has an important historical interest, more especially as at
       that time such an easy preparation of substances occurring in
       organisms without the aid of organic life was quite unexpected,
       for they were supposed to be formed under the influence of the
       forces acting in organisms, and without the latter their formation
       was considered impossible. And in addition to destroying this
       illusion, the easy transition of NH_{4}OCN into CO(NH_{2})_{2} is
       the best example of the passage of one system of equilibrium of
       atoms into another more stable system.

  [42] If ammonia and methane (marsh gas) do not show any acid
       properties, that is in all probability due to the presence of a
       large amount of hydrogen in both; but in hydrocyanic acid one atom
       of hydrogen is under the influence of two acid-forming elements.
       Acetylene, C_{2}H_{2}, which contains but little hydrogen,
       presents acid properties in certain respects, for its hydrogen is
       easily replaced by metals. Hydronitrous acid, HN_{3}, which
       contains little hydrogen, also has the properties of an acid.

  [43] Solutions of cyanides--for instance, those of potassium or
       barium--are decomposed by carbonic acid. Even the carbonic
       anhydride of the air acts in a similar way, and for this reason
       these solutions do not keep, because, in the first place, free
       hydrocyanic acid itself decomposes and polymerises, and, in the
       second place, with alkaline liquids it forms ammonia and formic
       acid. Hydrocyanic acid does not liberate carbonic anhydride from
       solutions of sodium or potassium carbonates. But a mixture of
       solutions of potassium carbonate and hydrocyanic acid yields
       carbonic anhydride on the addition of oxides like zinc oxide,
       mercuric oxide, &c. This is due to the great inclination which the
       cyanides exhibit of forming double salts. For instance,
       ZnK_{2}(CN)_{4} is formed, which is a soluble double salt.

  [43 bis] The conversion of the atmospheric nitrogen into cyanogen
       compounds, although possible, has not yet been carried out on a
       large scale, and one of the problems for future research should be
       the discovery of a practical and economical means of converting
       the atmospheric nitrogen into metallic cyanides, not only because
       potassium cyanide has found a vast and important use for the
       extraction of gold from even the poorest ores, but more especially
       because the cyanides furnish the means for effecting the synthesis
       of many complex carbon compounds, and the nitrogen contained in
       cyanogen easily passes into other forms of combination such as
       ammonia, which is of great importance in agriculture.

If four parts of yellow prussiate be mixed with eight parts of water and
three parts of sulphuric acid, and the mixture be heated, it decomposes,
volatile hydrocyanic acid separating. This was obtained for the first
time by Scheele in 1782, but it was only known to him in solution. In
1809 Ittner prepared anhydrous prussic acid, and in 1815 Gay-Lussac
finally settled its properties and showed that it contains only hydrogen,
carbon, and nitrogen, CNH. If the distillate (a weak solution of HCN) be
redistilled, and the first part collected, the anhydrous acid may be
prepared from this stronger solution. In order to do this, pieces of
calcium chloride are added to the concentrated solution, when the
anhydrous acid floats as a separate layer, because it is not soluble in
an aqueous solution of calcium chloride. If this layer be then distilled
over a new portion of calcium chloride at the lowest temperature
possible, the prussic acid may be obtained completely free from water. It
is, however, necessary to use the greatest caution in work of this kind,
because prussic acid, besides being extremely poisonous, is exceedingly
volatile.[44]

  [44] The mixture of the vapours of water and hydrocyanic acid, evolved
       on heating yellow prussiate with sulphuric acid, may be passed
       directly through vessels or tubes filled with calcium chloride.
       These tubes must be cooled, because, in the first place,
       hydrocyanic acid easily changes on being heated, and, in the
       second place, the calcium chloride when warm would absorb less
       water. The mixture of hydrocyanic acid and aqueous vapour on
       passing over a long layer of calcium chloride gives up water, and
       hydrocyanic acid alone remains in the vapour. It ought to be
       cooled as carefully as possible in order to bring it into a liquid
       condition. The method which Gay-Lussac employed for obtaining pure
       hydrocyanic acid consisted in the action of hydrochloric acid gas
       on mercuric cyanide. The latter may he obtained in a pure state if
       a solution of yellow prussiate be boiled with a solution of
       mercuric nitrate, filtered, and crystallised by cooling; the
       mercuric cyanide is then obtained in the form of colourless
       crystals, Hg(CN)_{2}.

       If a strong solution of hydrochloric acid be poured upon these
       crystals, and the mixture of vapours evolved, consisting of
       aqueous vapour, hydrochloric acid, and hydrocyanic acid, be passed
       through a tube containing, first, marble (for absorbing the
       hydrochloric acid), and then lumps of calcium chloride, on cooling
       the hydrocyanic acid will be condensed. In order to obtain the
       latter in an anhydrous form, the decomposition of heated mercury
       cyanide by hydrogen sulphide may be made use of. Here the sulphur
       and cyanogen change places, and hydrocyanic acid and mercury
       sulphide are formed: Hg(CN)_{2} + H_{2}S = 2HCN + HgS.

Anhydrous prussic acid is a very mobile and volatile liquid; its specific
gravity is 0·697 at 18°; at lower temperatures, especially when mixed
with a small quantity of water, it easily congeals; it boils at 26°, and
therefore very easily evaporates, and at ordinary temperatures may be
regarded as a gas. An insignificant amount, when inhaled or brought into
contact with the skin, causes death. It is soluble in all proportions in
water, alcohol, and ether: weak aqueous solutions are used in
medicine.[45]

  [45] A weak (up to 2 p.c.) aqueous solution of hydrocyanic acid is
       obtained by the distillation of certain vegetable substances. The
       so-called laurel water in particular enjoys considerable notoriety
       from its containing hydrocyanic acid. It is obtained by the
       steeping and distillation of laurel leaves. A similar kind of
       water is formed by the infusion and distillation of bitter
       almonds. It is well known that bitter almonds are poisonous, and
       have a peculiar characteristic taste. This bitter taste is due to
       the presence of a certain substance called amygdalin, which can be
       extracted by alcohol. This amygdalin decomposes in an infusion of
       bruised almonds, forming the so-called bitter almond oil, glucose,
       and hydrocyanic acid:

       C_{10}H_{27}NO_{11} + H_{2}O = C_{7}H_{6}O + CNH + 2C_{6}H_{12}O_{6}
          Amygdalin in        Water      Bitter Hydrocyanic    Glucose
         bitter almonds                  almond    acid
                                           oil

       If after this the infusion of bitter almonds be distilled with
       water, the hydrocyanic acid and the volatile bitter almond oil are
       carried over with the aqueous vapour. The oil is insoluble in
       water, or only sparingly soluble, while the hydrocyanic acid
       remains as an aqueous solution. Bitter almond water is similar to
       laurel water, and is used like the former in medicine, naturally
       only in small quantities because any considerable amount has
       poisonous effects. Perfectly pure anhydrous hydrocyanic acid keeps
       without change, just like the weak solutions, but the strong
       solutions only keep in the presence of other acids. In the
       presence of many admixtures these solutions easily give a brown
       polymeric substance, which is also formed in a solution of
       potassium cyanide.

The salts MCN--for instance, potassium, sodium,
ammonium--as well as the salts M´´(CN)_{2}--for example, barium,
calcium, mercury--are soluble in water, but the cyanides of manganese,
zinc, lead, and many others are insoluble in water. They form double
salts with potassium cyanide and similar metallic cyanides, an example
of which we will consider in a further description of the yellow
prussiate. Not only are some of the double salts remarkable for their
constancy and comparative stability, but so also are the soluble salt
HgC_{2}N_{2}, the insoluble silver cyanide AgCN, and even potassium
cyanide in the absence of water. The last salt,[46] when fused, acts
as a reducing agent with its elements K and C, and oxidises when fused
with lead oxide, forming potassium cyanate, KOCN, which establishes the
connection between HCN and OHCN--that is, between the nitriles of formic
and carbonic acids--and this connection is the same as that between
the acids themselves, since formic acid, on oxidation, yields carbonic
acid. Free cyanogen, (CN)_{2} or CNCN, corresponds to hydrocyanic acid
in the same manner as free chlorine, Cl_{2} or ClCl, corresponds to
hydrochloric acid. This composition, judging from what has been already
stated, exactly expresses that of the nitrile of oxalic acid, and, as
a matter of fact, oxalate of ammonia and the amide corresponding with
it (oxamide, Note 33), on being heated with phosphoric anhydride, which
takes up the water, yield _cyanogen_, (CN)_{2}. This substance is also
produced by simply heating some of the metallic cyanides. Mercuric
cyanide is particularly adapted for this purpose, because it is easily
obtained in a pure state and is then very stable. If mercuric cyanide be
heated, it decomposes, in like manner to mercury oxide, into metallic
mercury and cyanogen: HgC_{2}N_{2} = Hg + C_{2}N_{2}.[47] When cyanogen
is formed, part of it always polymerises into a dark brown insoluble
substance called _paracyanogen_, capable of forming cyanogen when
heated to redness.[48] Cyanogen is a colourless, poisonous gas, with a
peculiar smell and easily condensed by cooling into a colourless liquid,
insoluble in water and having a specific gravity of 0·86. It boils at
about -21°, and therefore cyanogen may be easily condensed into a liquid
by a strong freezing mixture. At -35° liquid cyanogen solidifies. The gas
is soluble in water and in alcohol to a considerable extent--namely,
1 volume of water absorbs as much as 4-1/2 volumes, and alcohol 23
volumes. Cyanogen resists the action of a tolerably high temperature
without decomposing, but under the action of the electric spark the
carbon is separated, leaving a volume of nitrogen equal to the volume
of the gas taken. As it contains carbon it burns, and the colour of the
flame is reddish-violet, which is due to the presence of nitrogen, all
compounds of which impart more or less of this reddish-violet hue to
the flame. During the combustion of cyanogen, carbonic anhydride and
nitrogen are formed. The same products are obtained in the eudiometer
with oxygen or by the action of cyanogen on many oxides at a red heat.

  [46] This salt will be described in Chapter XIII.

  [47] For the preparation it is necessary to take completely dry mercuric
       cyanide, because when heated in the presence of moisture it gives
       ammonia, carbonic anhydride, and hydrocyanic acid. Instead of
       mercuric cyanide, a mixture of perfectly dry yellow prussiate and
       mercuric chloride may be used, then double decomposition and the
       formation of mercuric cyanide take place in the retort. Silver
       cyanide also disengages cyanogen, on being heated.

  [47] _Paracyanogen_ is a brown substance (having the composition of
       cyanogen) which is formed during the preparation of cyanogen by
       all methods, and remains as a residue. Silver cyanide, on being
       slightly heated, fuses, and on being further heated evolves a gas;
       a considerable quantity of paracyanogen remains in the residue.
       Here it is remarkable that exactly half the cyanogen becomes
       gaseous, and the other half is transformed into paracyanogen.
       Metallic silver will be found in the residue with the
       paracyanogen; it may be extracted with mercury or nitric acid,
       which does not act on paracyanogen. If paracyanogen be heated in a
       vacuum it decomposes, forming cyanogen; but here the pressure _p_
       for a given temperature _t_ cannot exceed a certain limit, so that
       the phenomenon presents all the external appearance of a physical
       transformation into vapour; but, nevertheless, it is a complete
       change in the nature of the substance, though limited by the
       _pressure of dissociation_, as we saw before in the transformation
       of cyanuric into hydrocyanic acid, and as would be expected from
       the fundamental principles of dissociation. Troost and
       Hautefeuille (1868) found that for paracyanogen,

                    _t_ = 530°  581°  600°    635°
                    _p_ =  90   143   296   1,089 mm.

       However, even at 550° part of the cyanogen decomposes into carbon
       and nitrogen. The reverse transition of cyanogen into paracyanogen
       commences at 350°, and at 600° proceeds rapidly. And if the
       transition of the first kind is likened to evaporation, then the
       reverse transition, or polymerisation, presents a likeness to the
       transition of vapours into the solid state.

The relation of cyanogen to the metallic cyanides is seen not only in the
fact that it is formed from mercuric cyanide, but also by its forming
cyanide of sodium or potassium on being heated with either of those
metals, the sodium or potassium taking fire in the cyanogen. On heating a
mixture of hydrogen and cyanogen to 500° (Berthelot),[49] or under the
action of the silent discharge (Boilleau), hydrocyanic acid is formed, so
that the reciprocity of the transitions does not leave any doubt in the
matter that all the nitriles of the organic acids contain cyanogen, just
as all the organic acids contain carboxyl and in it the elements of
carbonic anhydride. Besides the amides,[50] the nitriles (or cyanogen
compounds, RCN), and nitro-compounds (containing the radicle of nitric
acid, RNO_{2}), there are a great number of other substances containing
at the same time carbon and nitrogen, particulars of which must be sought
for in special works on organic chemistry.

  [49] Cyanogen (like chlorine) is absorbed by a solution of sodium
       hydroxide, sodium cyanide and cyanate being produced: C_{2}N_{2} +
       2NaHO = NaCN + CNNaO + H_{2}O. But the latter salt decomposes
       relatively easily, and moreover part of the cyanogen liberated by
       heat from its compounds undergoes a more complex transformation.

  [50] If, in general, compounds containing the radicle NH_{2} are called
       amides, some of the _amines_ ought to be ranked with them; namely,
       the hydrocarbons C_{_n_}H_{2_m_}, in which part of the hydrogen is
       replaced by NH_{2}; for instance, methylamine, CH_{3}NH_{2},
       aniline, C_{6}H_{5}NH_{2}, &c. In general the amines may be
       represented as ammonia in which part or all of the hydrogen is
       replaced by hydrocarbon radicles--as, for example, trimethylamine,
       N(CH_{3})_{3}. They, like ammonia, combine with acids and form
       crystalline salts. Analogous substances are sometimes met with in
       nature, and bear the general name of _alkaloids_; such are, for
       instance, quinine in cinchona bark, nicotine in tobacco, &c.




                                CHAPTER X

          SODIUM CHLORIDE--BERTHOLLET'S LAWS--HYDROCHLORIC ACID


In the preceding chapters we have become acquainted with the most
important properties of the four elements, hydrogen, oxygen, nitrogen,
and carbon. They are sometimes termed the _organogens_, because
they enter into the composition of organic substances. Their mutual
combinations may serve as types for all other chemical compounds--that
is, they present the same atomic relations (types, forms, or grades of
combinations) as those in which the other elements also combine together.

               Hydrogen,  HH,     or, in general, HR.
               Water,     H_{2}O,      "     "    H_{2}R.
               Ammonia,   H_{3}N,      "     "    H_{3}R.
               Marsh gas, H_{4}C,      "     "    H_{4}R.

One, two, three, and four atoms of hydrogen enter into these molecules
for one atom of another element. No compounds of one atom of oxygen with
three or four atoms of hydrogen are known; hence the atom of oxygen does
not possess certain properties which are found in the atoms of carbon
and nitrogen.

The faculty of an element to form a compound of definite composition
with hydrogen (or an element analogous to it) gives the possibility
of foretelling the composition of many other of its compounds. Thus,
if we know that an element, M, combines with hydrogen, forming, by
preference, a gaseous substance such as HM, but not forming H_{2}M,
H_{3}M, H_{n}M_{m}, then we must conclude, on the basis of the law of
substitution, that this element will give compounds M_{2}O, M_{3}N,
MHO, MH_{3}C, &c. Chlorine is an example of this kind. If we know
that another element, R, like oxygen, gives with hydrogen a molecule
H_{2}R, then we may expect that it will form compounds similar to
hydrogen peroxide, the metallic oxides, carbonic anhydride, or carbonic
oxide, and others. Sulphur is an instance of this kind. Hence the
elements may be classified according to their resemblance to hydrogen,
oxygen, nitrogen, and carbon, and in conformity with this analogy it
is possible to foretell, if not the properties (for example, the
acidity or basicity), at any rate the composition,[1] of some of their
compounds. This forms the substance of _the conception of the valency
or atomicity of the elements_. Hydrogen is taken as the representative
of the univalent elements, giving compounds, RH, R(OH), R_{2}O, RCl,
R_{3}N, R_{4}C, &c. Oxygen, in that form in which it gives water, is the
representative of the bivalent elements, forming RH_{2}, RO, RCl_{2},
RHCl, R(OH)Cl, R(OH)_{2}, R_{2}C, RCN, &c. Nitrogen in ammonia is the
representative of the trivalent elements, giving compounds RH_{3},
R_{2}O_{3}, R(OH)_{3}, RCl_{3}, RN, RHC, &c. In carbon are exemplified
the properties of the quadrivalent elements, forming RH_{4}, RO_{2},
RO(OH)_{2}, R(OH)_{4}, RHN, RCl_{4}, RHCl_{3}, &c. We meet with these
_forms of combination_, or degrees of union of atoms, in all other
elements, some being analogous to hydrogen, others to oxygen, and others
to nitrogen or to carbon. But besides these quantitative analogies or
resemblances, which are foretold by the law of substitution (Chapter
VI.), there exist among the elements qualitative analogies and relations
which are not fully seen in the compounds of the elements which have
been considered, but are most distinctly exhibited in the formation of
bases, acids, and salts of different types and properties. Therefore,
for a complete study of the nature of the elements and their compounds
it is especially important to become acquainted with the salts, as
substances of a peculiar character, and with the corresponding acids and
bases. Common table salt, or sodium chloride, NaCl, may in every respect
be taken as a type of salts in general, and we will therefore pass to
the consideration of this substance, and of hydrochloric acid, and of
the base sodium hydroxide, formed by the non-metal chlorine and the
metal sodium, which correspond with it.

  [1] But it is impossible to foretell all the compounds formed by an
      element from its atomicity or valency, because the atomicity of the
      elements is variable, and furthermore this variability is not
      identical for different elements. In CO_{2}, COX_{2}, CH_{4}, and
      the multitude of carbon compounds corresponding with them, the C is
      quadrivalent, but in CO either the carbon must be taken as bivalent
      or the atomicity of oxygen be accounted as variable. Moreover,
      carbon is an example of an element which preserves its atomicity to
      a greater degree than most of the other elements. Nitrogen in
      NH_{3}, NH_{2}(OH), N_{2}O_{3}, and even in CNH, must be considered
      as trivalent, but in NH_{4}Cl, NO_{2}(OH), and in all their
      corresponding compounds it is necessarily pentavalent. In N_{2}O,
      if the atomicity of oxygen = 2, nitrogen has an uneven atomicity
      (1, 3, 5), whilst in NO it is bivalent. If sulphur be bivalent,
      like oxygen, in many of its compounds (for example, H_{2}S,
      SCl_{2}, KHS, &c.), then it could not be foreseen from this that it
      would form SO_{2}, SO_{3}, SCl_{4}, SOCl_{2}, and a series of
      similar compounds in which its atomicity must be acknowledged as
      greater than 2. Thus SO_{2}, sulphurous anhydride, has many points
      in common with CO_{2}, and if carbon be quadrivalent then the S in
      SO_{2} is quadrivalent. Therefore the principle of atomicity
      (valency) of the elements cannot be considered established as the
      basis for the study of the elements, although it gives an easy
      method of grasping many analogies. I consider the four following as
      the chief obstacles to acknowledging the atomicity of the elements
      as a primary conception for the consideration of the properties of
      the elements: 1. Such univalent elements as H, Cl, &c., appear in a
      free state as molecules H_{2}, Cl_{2}, &c., and are consequently
      like the univalent radicles CH_{3}, OH, CO_{2}H, &c., which, as
      might be expected, appear as C_{2}H_{6}, O_{2}H_{2},
      C_{2}O_{4}H_{2} (ethane, hydrogen peroxide, oxalic acid), whilst on
      the other hand, potassium and sodium (perhaps also iodine at a high
      temperature) contain only one atom, K, Na, in the molecule in a
      free state. Hence it follows that _free affinities_ may exist.
      Granting this, nothing prevents the assumption that free affinities
      exist in all unsaturated compounds; for example, two free
      affinities in NH_{3}. If such instances of free affinities be
      admitted, then all the possible advantages to be gained by the
      application of the doctrine of atomicity (valency) are lost. 2.
      There are instances--for example, Na_{2}H--where univalent elements
      are combined in molecules which are more complex than R_{2}, and
      form molecules, R_{3}, R_{4}, &c.; this may again be either taken
      as evidence of the existence of free affinities, or else
      necessitates such primary univalent elements as sodium and hydrogen
      being considered as variable in their atomicity. 3. The periodic
      system of the elements, with which we shall afterwards become
      acquainted, shows that there is a law or rule for the variation of
      the forms of oxygen and hydrogen compounds; chlorine is univalent
      with respect to hydrogen, and septavalent with respect to oxygen;
      sulphur is bivalent to hydrogen, and sexavalent to oxygen;
      phosphorus is trivalent to hydrogen and pentavalent in respect to
      oxygen--the sum is in every case equal to 8. Only carbon and its
      analogues (for example, silicon) are quadrivalent to both hydrogen
      and oxygen. Hence the power of the elements to change their
      atomicity is an essential part of their nature, and therefore
      constant valency cannot he considered as a fundamental property. 4.
      Crystallo-hydrates (for instance, NaCl,2H_{2}O, or NaBr,2H_{2}O),
      double salts (such as PtCl_{4},2KCl,H_{2}SiF_{6}, &c.), and similar
      complex compounds (and, according to Chap. I., solutions also)
      demonstrate the capacity not only of the elements themselves, but
      also of their saturated and limiting compounds, of entering into
      further combination. Therefore the admission of a definite limited
      atomicity of the elements includes in itself an admission of
      limitation which is not in accordance with the nature of chemical
      reactions.

_Sodium chloride_, NaCl, the familiar table salt, occurs, although in
very small quantities, in all the primary formations of the earth's
crust,[2] from which it is washed away by the atmospheric waters; it is
contained in small quantities in all waters flowing through these
formations, and is in this manner conveyed to the oceans and seas. The
immense mass of salt in the oceans has been accumulated by this process
from the remote ages of the earth's creation, because the water has
evaporated from them while the salt has remained in solution. The salt of
sea water serves as the source not only for its direct extraction, but
also for the formation of other masses of workable salt, such as rock
salt, and of saline springs and lakes.[2 bis]

  [2] The primary formations are those which do not bear any distinct
      traces of having been deposited from water (have not a stratified
      formation and contain no remains of animal or vegetable life),
      occur under the sedimentary formations of the earth, and are
      everywhere uniform in composition and structure, the latter being
      generally distinctly crystalline. If it be assumed that the earth
      was originally in a molten condition, the first primary formations
      are those which formed the first solid crust of the earth. But even
      with this hypothesis of the earth's origin, it is necessary to
      admit that the first aqueous deposits must have caused a change in
      the original crust of the earth, and therefore under the head of
      primary formations must be understood the most ancient of the
      products of decomposition (mostly by atmospheric, aqueous, and
      organic agency, &c.), from which all the rocks and substances of
      the earth's surface have arisen. In speaking of the origin of one
      or another substance, we can only, on the basis of facts, descend
      to the primary formations, of which granite, gneiss, and trachyte
      may be taken as examples.

  [2 bis] Chloride of sodium has been found to occur in the atmosphere in
      the form of a fine dust; in the lower strata it is present in
      larger quantities than in the upper, so that the rain water falling
      on mountains contains less NaCl than that falling in valleys. Müntz
      (1891) found that a litre of rain water collected on the summit of
      the Pic du Midi (2,877 metres above the sea level) contained 0·34
      milligram of chloride of sodium, while a litre of rain collected
      from the valley contained 2·5-7·6 milligrams.

The extraction of salt _from sea water_ is carried on in several ways. In
southern climes, especially on the shores of the Atlantic Ocean and the
Mediterranean and Black Seas, the summer heats are taken advantage of. A
convenient low-lying sea shore is chosen, and a whole series of basins,
communicating with each other, are constructed along it. The upper of
these basins are filled with sea water by pumping, or else advantage is
taken of high tides. These basins are sometimes separated from the sea by
natural sand-banks (limans) or by artificial means, and in spring the
water already begins to evaporate considerably. As the solution becomes
more concentrated, it is run into the succeeding basins, and the upper
ones are supplied with a fresh quantity of sea water, or else an
arrangement is made enabling the salt water to flow by degrees through
the series of basins. It is evident that the beds of the basins should be
as far as possible impervious to water, and for this purpose they are
made of beaten clay. The crystals of salt begin to separate out when the
concentration attains 28 p.c. of salt (which corresponds to 28° of
Baumé's hydrometer). They are raked off, and employed for all those
purposes to which table salt is applicable. In the majority of cases only
the first half of the sodium chloride which can be separated from the sea
water is extracted, because the second half has a bitter taste from the
presence of magnesium salts which separate out together with the sodium
salt. But in certain localities--as, for instance, in the estuary of the
Rhone, on the island of Camarga[3]--the evaporation is carried on to the
very end, in order to obtain those magnesium and potassium salts which
separate out at the end of the evaporation of sea water. Various salts
are separated from sea water in its evaporation. From 100 parts of sea
water there separates out, by natural and artificial evaporation, about
one part of tolerably pure table salt at the very commencement of the
operation; the total amount held in solution being about 2-1/2 p.c. The
remaining portion separates out intermixed with the bitter salts of
magnesium which, owing to their solubility and the small amount in which
they are present (less than 1 p.c.), only separate out, in the first
crystallisations, in traces. Gypsum, or calcium sulphate,
CaSO_{4},2H_{2}O, because of its sparing solubility, separates together
with or even before the table salt. When about half of the latter has
separated, then a mixture of table salt and magnesium sulphate separates
out, and on still further evaporation the chlorides of potassium and
magnesium begin to separate in a state of combination, forming the double
salt KMgCl_{3},6H_{2}O, which occurs in nature as _carnallite_.[4] After
the separation of this salt from sea water, there remains a mother liquor
containing a large amount of magnesium chloride in admixture with various
other salts.[5] The extraction of sea salt is usually carried on for the
purpose of procuring table salt, and therefore directly it begins to
separate mixed with a considerable proportion[6] of magnesium salts (when
it acquires a bitter taste) the remaining liquor is run back into the
sea.

  [3] The extraction of the potassium salts (or so-called summer salts)
      was carried on at the Isle of Camarga about 1870, when I had
      occasion to visit that spot. At the present time the deposits of
      Stassfurt provide a much cheaper salt, owing to the evaporation and
      separation of the salt being carried on there by natural means and
      only requiring a treatment and refining, which is also necessary in
      addition for the 'summer salt' obtained from sea-water.

  [4] The double salt KCl,MgCl_{2} is a crystallohydrate of KCl and
      MgCl_{2}, and is only formed from solutions containing an excess of
      magnesium chloride, because water decomposes this double salt,
      extracting the more soluble magnesium chloride from it.

  [5] Owing to the fundamental property of salts of interchanging their
      metals, it cannot be said that sea water contains this or that
      salt, but only that it contains certain amounts of certain metals M
      (univalent like Na and K, and bivalent like Mg and Ca), and haloids
      X (univalent like Cl, Br, and bivalent like SO_{4}, CO_{3}), which
      are disposed in every possible kind of grouping; for instance, K as
      KCl, KBr, K_{2}SO_{4}, Mg as MgCl_{2}, MgBr_{2}, MgSO_{4}, and so
      on for all the other metals. In evaporation different salts
      separate out consecutively only because they reach saturation. A
      proof of this may be seen in the fact that a solution of a mixture
      of sodium chloride and magnesium sulphate (both of which salts are
      obtained from sea water, as was mentioned above), when evaporated,
      deposits crystals of these salts, but when refrigerated (if the
      solution be sufficiently saturated) the salt Na_{2}SO_{4},10H_{2}O
      is first deposited because it is the first to arrive at saturation
      at low temperatures. Consequently this solution contains MgCl_{2}
      and Na_{2}SO_{4}, besides MgSO_{4} and NaCl. So it is with sea
      water.

  [6] The salt extracted from water is piled up in heaps and left exposed
      to the action of rain water, which purifies it, owing to the water
      becoming saturated with sodium chloride and then no longer
      dissolving it, but washing out the impurities.

The same process which is employed for artificially obtaining salt in a
crystalline form from sea water has been repeatedly accomplished during
the geological evolution of the earth on a gigantic scale; upheavals of
the earth have cut off portions of the sea from the remainder (as the
Dead Sea was formerly a part of the Mediterranean, and the Sea of Aral of
the Caspian), and their water has evaporated and formed (if the mass of
the inflowing fresh water were less than that of the mass evaporated)
deposits of _rock salt_. It is always accompanied by gypsum, because the
latter is separated from sea water with or before the sodium chloride.
For this reason rock salt may always be looked for in those localities
where there are deposits of gypsum. But inasmuch as the gypsum remains on
the spot where it has been deposited (as it is a sparingly soluble salt),
whilst the rock salt (as one which is very soluble) may be washed away by
rain or fresh running water, it may sometimes happen that although gypsum
is still found there may be no salt; but, on the other hand, where there
is rock salt there will always be gypsum. As the geological changes of
the earth's surface are still proceeding at the present day, so in the
midst of the dry land salt lakes are met with, which are sometimes
scattered over vast districts formerly covered by seas now dried up. Such
is the origin of many of the salt lakes about the lower portions of the
Volga and in the Kirghiz steppes, where at a geological epoch preceding
the present the Aralo-Caspian Sea extended. Such are the Baskunchaksky
(in the Government of Astrakhan, 112 square kilometres superficial area),
the Eltonsky (140 versts from the left bank of the Volga, and 200 square
kilometres in superficial area), and upward of 700 other salt lakes lying
about the lower portions of the Volga. In those in which the inflow of
fresh water is less than that yearly evaporated, and in which the
concentration of the solution has reached saturation, the
_self-deposited_ salt is found already deposited on their beds, or is
being yearly deposited during the summer months. Certain limans, or
sea-side lakes, of the Azoff Sea are essentially of the same
character--as, for instance, those in the neighbourhood of Henichesk and
Berdiansk. The saline soils of certain Central Asian steppes, which
suffer from a want of atmospheric fresh water, are of the same origin.
Their salt originally proceeded from the salt of seas which previously
covered these localities, and has not yet been washed away by fresh
water. The main result of the above-described process of nature is the
formation of masses of rock salt, which are, however, being gradually
washed away by the subsoil waters flowing in their neighbourhood, and
afterwards rising to the surface in certain places as _saline springs_,
which indicate the presence of masses of deposited rock salt in the
depths of the earth. If the subsoil water flows along a stratum of salt
for a sufficient length of time it becomes saturated; but in flowing in
its further course along an impervious stratum (clay) it becomes diluted
by the fresh water leaking through the upper soil, and therefore the
greater the distance of a saline spring from the deposit of rock salt,
the poorer will it be in salt. A perfectly saturated brine, however, may
be procured from the depths of the earth by means of bore-holes. The
deposits of rock salt themselves, which are sometimes hidden at great
depths below the earth's strata, may be discovered by the guidance of
bore-holes and the direction of the strata of the district. Deposits of
rock salt, about 35 metres thick and 20 metres below the surface, were
discovered in this manner in the neighbourhood of Brianstcheffky and
Dekonoffky, in the Bakhmut district of the Government of Ekaterinoslav.
Large quantities of most excellent rock salt are now (since 1880)
obtained from these deposits, whose presence was indicated by the
neighbouring salt springs (near Slaviansk and Bakhmut) and by bore-holes
which had been sunk in these localities for procuring strong (saturated)
brines. But the Stassfurt deposits of rock salt near Magdeburg in Germany
are celebrated as being the first discovered in this manner, and for
their many remarkable peculiarities.[7] The plentiful distribution of
saline springs in this and the neighbouring districts suggested the
presence of deposits of rock salt in the vicinity. Deep bore-holes sunk
in this locality did in fact give a richer brine--even quite saturated
with salt. On sinking to a still greater depth, the deposits of salt
themselves were at last arrived at. But the first deposit which was met
with consisted of a bitter salt unfit for consumption, and was therefore
called refuse salt (_Abraumsalz_). On sinking still deeper vast beds of
real rock salt were struck. In this instance the presence of these upper
strata containing salts of potassium, magnesium, and sodium is an
excellent proof of the formation of rock salt from sea water. It is very
evident that not only a case of evaporation to the end--as far, for
instance, as the separation of carnallite--but also the preservation of
such soluble salts as separate out from sea water after the sodium
chloride, must be a very exceptional phenomenon, which is not repeated in
all deposits of rock salt. The Stassfurt deposits therefore are of
particular interest, not only from a scientific point of view, but also
because they form a rich source of potassium salts which have many
practical uses.[7 bis]

  [7] When the German savants pointed out the exact locality of the
      Stassfurt salt-beds and their depth below the surface, on the basis
      of information collected from various quarters respecting
      bore-holes and the direction of the strata, and when the borings,
      conducted by the Government, struck a salt-bed which was bitter and
      unfit for use, there was a great outcry against science, and the
      doubtful result even caused the cessation of the further work of
      deepening the shafts. It required a great effort to persuade the
      Government to continue the work. Now, when the pure salt
      encountered below forms one of the important riches of Germany, and
      when those 'refuse salts' have proved to be most valuable (as a
      source of potassium and magnesium), we should see in the
      utilisation of the Stassfurt deposits one of the conquests of
      science for the common welfare.

  [7 bis] In Western Europe, deposits of rock salt have long been known
      at Wieliczka, near Cracow, and at Cardona in Spain. In Russia the
      following deposits are known: (_a_) the vast masses of rock salt (3
      square kilometres area and up to 140 metres thick) lying directly
      on the surface of the earth at Iletzky Zastchit, on the left bank
      of the river Ural, in the Government of Orenburg; (_b_) the
      Chingaksky deposit, 90 versts from the river Volga, in the
      Enotaeffsky district of the Government of Astrakhan; (_c_) the
      Kulepinsky (and other) deposits (whose thickness attains 150
      metres), on the Araks, in the Government of Erivan in the Caucasus;
      (_d_) the Katchiezmansky deposit in the province of Kars; (_e_) the
      Krasnovodsky deposit in the Trans-Caspian province; and (_f_) the
      Bardymkulsky salt mines in Kokhand.

A saturated brine, formed by the continued contact of subsoil water with
rock salt, is extracted by means of bore-holes, as, for instance, in the
Governments of Perm, Kharkoff, and Ekaterinoslav. Sometimes, as at
Berchtesgaden (and at Hallein) in Austria, spring water is run on to
underground beds of rock salt containing much clay.

[Illustration: FIG. 64.--Graduator for the evaporation of the water of
saline springs.]

If a saline spring, or the salt water pumped from bore-holes, contains
but little salt, then the first concentration of the natural solution
is not carried on by the costly consumption of fuel, but by the cheaper
method of evaporation by means of the wind. For this purpose so-called
graduators are constructed: they consist of long and lofty sheds,
which are sometimes several versts long, and generally extend in a
direction at right angles to that of the usual course of the wind in
the district. These sheds are open at the sides, and are filled with
brushwood as shown in fig. 64. Troughs, A B, C D, into which the salt
water is pumped, run along the top. On flowing from these troughs,
through the openings, _a_, the water spreads over the brushwood and
distributes itself in a thin layer over it, so that it presents a very
large surface for evaporation, in consequence of which it rapidly
becomes concentrated in warm or windy weather. After trickling over the
brushwood, the solution collects in a reservoir under the graduator,
whence it is usually pumped up by the pumps P P´, and again run a
second and third time through the graduator, until the solution reaches
a degree of concentration at which it becomes profitable to extract
the salt by direct heating. Generally the evaporation in the graduator
is not carried beyond a concentration of 12 to 15 parts of salt in
100 parts of solution. Strong natural solutions of salt, and also the
graduated solutions, are evaporated in large shallow metallic vessels,
which are either heated by the direct action of the flame from below
or from above. These vessels are made of boiler plate, and are called
salt-pans. Various means are employed for accelerating the evaporation
and for economising fuel, which are mainly based on an artificial
draught to carry off the steam as it is formed, and on subjecting the
saline solution to a preliminary heating by the waste heat of the
steam and furnace gases. Furthermore, the first portions of the salt
which crystallise out in the salt-pans are invariably contaminated
with gypsum, since the waters of saline springs always contain this
substance. It is only the portions of the salt which separate later that
are distinguished by their great purity. The salt is ladled out as it is
deposited, left to drain on inclined tables and then dried, and in this
manner the so-called bay salt is obtained. Since it has become possible
to discover the saline deposits themselves, the extraction of table salt
from the water of saline springs by evaporation, which previously was
in general use, has begun to be disused, and is only able to hold its
ground in places where fuel is cheap.

In order to understand the full importance of the extraction of salt,
it need only be mentioned that on the average 20 lbs. of table salt are
consumed yearly per head of population, directly in food or for cattle.
In those countries where common salt is employed in technical processes,
and especially in England, almost an equal quantity is consumed in the
production of substances containing chlorine and sodium, and especially
in the manufacture of washing soda, &c., and of chlorine compounds
(bleaching powder and hydrochloric acid). The yearly production of salt
in Europe amounts to as much as 7-1/2 million tons.

Although certain lumps of rock salt and crystals of bay salt sometimes
consist of almost pure sodium chloride, still the ordinary commercial
salt contains various impurities, the most common of which are
magnesium salts. If the salt be pure, its solution gives no precipitate
with sodium carbonate, Na_{2}CO_{3}, showing the absence of magnesium
salts, because magnesium carbonate, MgCO_{3}, is insoluble in water.
Rock salt, which is ground for use, generally contains also a
considerable admixture of clay and other insoluble impurities.[8] For
ordinary use the bulk of the salt obtained can be employed directly
without further purification; but some salts are purified by solution
and crystallisation of the solution after standing, in which case the
evaporation is not carried on to dryness, and the impurities remain
in the _mother liquor_ or in the sediment. When perfectly pure salt
is required for chemical purposes it is best obtained as follows: a
saturated solution of table salt is prepared, and hydrochloric acid
gas is passed through it; this precipitates the sodium chloride (which
is not soluble in a strong solution of hydrochloric acid), while the
impurities remain in solution. By repeating the operation and fusing
the salt (when adhering hydrochloric acid is volatilised) a pure
salt is obtained, which is again crystallised from its solution by
evaporation.[9]

  [8] The fracture of rock salt generally shows the presence of
      interlayers of impurities which are sometimes very small in weight,
      but visible owing to their refraction. In the excellently laid out
      salt mines of Briansk I counted (1888), if my memory does not
      deceive me, on an average ten interlayers per metre of thickness,
      between which the salt was in general very pure, and in places
      quite transparent. If this be the case, then there would be 350
      interlayers for the whole thickness (about 35 metres) of the bed.
      They probably correspond with the yearly deposition of the salt. In
      this case the deposition would have extended over more than 300
      years. This should be observable at the present day in lakes where
      the salt is saturated and in course of deposition.

  [9] My own investigations have shown that not only the sulphates, but
      also the potassium salts, are entirely removed by this method.

Pure sodium chloride, in well-defined crystals (slowly deposited at the
bottom of the liquid) or in compact masses (in which form rock salt is
sometimes met with), is a colourless and transparent substance
resembling, but more brittle and less hard than, glass.[10] Common salt
always crystallises in the cubic system, most frequently in _cubes_, and
more rarely in octahedra. Large transparent cubes of common salt, having
edges up to 10 centimetres long, are sometimes found in masses of rock
salt.[11] When evaporated in the open the salt often separates out on
the surface[12] as cubes, which grow on to each other in the form of
pyramidal square funnels. In still weather, these clusters are able to
support themselves on the surface of the water for a long time, and
sometimes go on increasing to a considerable extent, but they sink
directly the water penetrates inside them. Salt fuses to a colourless
liquid (sp. gr. 1·602, according to Quincke) at 851° (V. Meyer); if pure
it solidifies to a non-crystalline mass, and if impure to an opaque mass
whose surface is not smooth. In fusing, sodium chloride commences to
volatilise (its weight decreases) and at a white heat it volatilises with
great ease and completely; but at the ordinary temperature it may, like
all ordinary salts, be considered as non-volatile, although as yet no
exact experiments have been made in this direction.

  [10] According to the determinations of Klodt, the Briansk rock salt
       withstands a pressure of 340 kilograms per square centimetre,
       whilst glass withstands 1,700 kilos. In this respect salt is twice
       as secure as bricks, and therefore immense masses may be extracted
       from underground workings with perfect safety, without having
       recourse to brickwork supports, merely taking advantage of the
       properties of the salt itself.

  [11] To obtain well-formed crystals, a saturated solution is mixed with
       ferric chloride, several small crystals of sodium chloride are
       placed at the bottom, and the solution is allowed to evaporate
       slowly in a vessel with a loose-fitting cover. Octahedral crystals
       are obtained by the addition of borax, urea, &c., to the solution.
       Very fine crystals are formed in a mass of gelatinous silica.

  [12] If a solution of sodium chloride be slowly heated from above,
       where the evaporation takes place, then the upper layer will
       become saturated before the lower and cooler layers, and therefore
       crystallisation will begin on the surface, and the crystals first
       formed will float, having also dried from above, on the surface
       until they become quite soaked. Being heavier than the solution
       the crystals are partially immersed under it, and the following
       crystallisation, also proceeding on the surface, will only form
       crystals along the side of the original crystals. A funnel is
       formed in this manner. It will be borne on the surface like a boat
       (if the liquid be quiescent), because it will grow more from the
       upper edges. We can thus understand this at first sight strange
       funnel form of crystallisation of salt. In explanation why the
       crystallisation under the above conditions begins at the surface
       and not at the lower layers, it must be mentioned that the
       specific gravity of a crystal of sodium chloride = 2·16, and that
       of a solution saturated at 25° contains 26·7 p.c. of salt and has
       a specific gravity at 25°/4° of 1·2004; at 15° a saturated
       solution contains 26·5 p.c. of salt and has a sp. gr. 1·203 at
       15°/4°. Hence a solution saturated at a higher temperature is
       specifically lighter, notwithstanding the greater amount of salt
       it contains. With many substances _surface crystallisation_ cannot
       take place because their solubility increases more rapidly with
       the temperature than their specific gravity decreases. In this
       case the saturated solution will always be in the lower layers,
       where also the crystallisation will take place. Besides which it
       may be added that as a consequence of the properties of water and
       solutions, when they are heated from above (for instance, by the
       sun's rays), the warmer layers being the lightest remain above,
       whilst when heated from below they rise to the top. For this
       reason the water at great depths below the surface is always cold,
       which has long been known. These circumstances, as well as those
       observed by Soret (Chapter I., Note 19), explain the great
       differences of density and temperature, and in the amount of salts
       held in the oceans at different latitudes (in polar and tropical
       climes) and at various depths.

A saturated[13] solution of table salt (containing 26·4 p.c.) has at the
ordinary temperature a specific gravity of about 1·2. The specific
gravity of the crystals is 2·167 (17°). The salt which separates out at
the ordinary and higher temperatures contains no water of
crystallisation;[14] but if the crystals are formed at a low
temperature, especially from a saturated solution cooled to -12°, then
they present a prismatic form, and contain two equivalents of water,
NaCl,2H_{2}O. At the ordinary temperature these crystals split up into
sodium chloride and its solution.[15] Unsaturated solutions of table salt
when cooled below 0° give[16] crystals of ice, but when the solution has
a composition NaCl,10H_{2}O it solidifies completely at a temperature of
-23°. A solution of table salt saturated at its boiling point boils at
about 109°, and contains about 42 parts of salt per 100 parts of water.

  [13] By combining the results of Poggiale, Müller, and Karsten (they
       are evidently more accurate than those of Gay-Lussac and others) I
       found that a saturated solution at _t_°, from 0° to 108°, contains
       35·7 + 0·024_t_ + 0·0002_t_^2 grams of salt per 100 grams of
       water. This formula gives a solubility at 0° = 35·7 grams (= 26·3
       p.c.), whilst according to Karsten it is 36·09, Poggiale 35·5, and
       Müller 35·6 grams.

  [14] Perfectly pure _fused_ salt is not hygroscopic, according to
       Karsten, whilst the crystallised salt, even when quite pure,
       attracts as much as 0·6 p.c. of water from moist air, according to
       Stas. (In the Briansk mines, where the temperature throughout the
       whole year is about +10°, it may be observed, as Baron Klodt
       informed me, that in the summer during damp weather the walls
       become moist, while in winter they are dry).

       If the salt contain impurities--such as magnesium sulphate, &c.--it
       is more hygroscopic. If it contain any magnesium chloride, it
       partially deliquesces in a damp atmosphere. The crystallised and
       not perfectly pure salt decrepitates when heated, owing to its
       containing water. The pure salt, and also the transparent rock
       salt, or that which has been once fused, does not decrepitate.
       Fused sodium chloride shows a faint alkaline reaction to litmus,
       which has been noticed by many observers, and is due to the
       presence of sodium oxide (probably by the action of the oxygen of
       the atmosphere). According to A. Stcherbakoff very sensitive
       litmus (washed in alcohol and neutralised with oxalic acid) shows
       an alkaline reaction even with the crystallised salt.

       It may be observed that rock salt sometimes contains cavities
       filled with a colourless liquid. Certain kinds of rock salt emit
       an odour like that of hydrocarbons. These phenomena have as yet
       received very little attention.

  [15] By cooling a solution of table salt saturated at the ordinary
       temperature to -15°, I obtained first of all well-formed tabular
       (six-sided) crystals, which when warmed to the ordinary
       temperature disintegrated (with the separation of anhydrous sodium
       chloride), and then prismatic needles up to 20 mm. long were
       formed from the same solution. I have not yet investigated the
       reason of the difference in crystalline form. It is known
       (Mitscherlich) that NaI,2H_{2}O also crystallises either in plates
       or prisms. Sodium bromide also crystallises with 2H_{2}O at the
       ordinary temperature.

  [16] Notwithstanding the great simplicity (Chapter I., Note 49) of the
       observations on the formation of ice from solution, still even for
       sodium chloride they cannot yet be considered as sufficiently
       harmonious. According to Blagden and Raoult, the temperature of
       the formation of ice from a solution containing _c_ grams of salt
       per 100 grams of water =-0·6_c_ to _c_ = 10, according to Rosetti
       =-0·649_c_ to _c_ = 8·7, according to De Coppet (to _c_ = 10)
       =-0·55_c_-0·006_c_^2, according to Karsten (to _c_ = 10)-0·762_c_
       + 0·0084_c_^2, and according to Guthrie a much lower figure. By
       taking Rosetti's figure and applying the rule given in Chapter I.,
       Note 49 we obtain--

                     _i_ - 0·649 × 58·5/18·5 = 2·05.

       Pickering (1893) gives for _c_ = 1-0·603, for _c_ = 2-1·220; that
       is (_c_ up to 2·7) about -(0·600 + 0·005_c_)_c_.

       The data for strong solutions are not less contradictory. Thus
       with 20 p.c. of salt, ice is formed at -14·4° according to
       Karsten,-17° according to Guthrie,-17·6° according to De Coppet.
       Rüdorff states that for strong solutions the temperature of the
       formation of ice descends in proportion to the contents of the
       compound, NaCl,2H_{2}O (per 100 grams of water) by 0°·342 per 1
       gram of salt, and De Coppet shows that there is no
       proportionality, in a strict sense, for either a percentage of
       NaCl or of NaCl,2H_{2}O.

Of all its physical properties the specific gravity of solutions of
sodium chloride is the one which has been the most fully investigated. A
comparison of all the existing determinations of the specific gravity of
solutions of NaCl[17] at 15° (in vacuo, taking water at 4° as 10,000),
with regard to _p_ (the percentage amount of the salt in solution), show
that it is expressed by the equation S_{15} = 9991·6 + 71·17_p_ +
0·2140_p_^2. For instance, for a solution 200H_{2}O + NaCl, in which case
_p_ = 1·6, S_{15} = 1·0106. It is seen from the formula that the addition
of water produces a contraction.[18] The specific gravity[19] at certain
temperatures and concentrations in vacuo referred to water at 4° =
10,000[20] is here given for

                           0°    15°    30°    100°
               _p_ =  5  10372  10353  10307   9922
                     10  10768  10728  10669  10278
                     15  11164  11107  11043  10652
                     20  11568  11501  11429  11043

It should be remarked that Baumé's hydrometer is graduated by taking a
10 p.c. solution of sodium chloride as 10° on the scale, and therefore
it gives approximately the percentage amount of the salt in a solution.
Common salt is somewhat soluble in alcohol,[21] but it is insoluble in
ether and in oils.

  [17] A collection of observations on the specific gravity of solutions
       of sodium chloride is given in my work cited in Chapter I.,
       Note 50.

       Solutions of common salt have also been frequently investigated as
       regards rate of _diffusion_ (Chapter I.), but as yet there are no
       complete data in this respect. It may be mentioned that Graham and
       De Vries demonstrated that diffusion in gelatinous masses (for
       instance, gelatin jelly, or gelatinous silica) proceeds in the
       same manner as in water, which may probably lead to a convenient
       and accurate method for the investigation of the phenomena of
       diffusion. N. Umoff (Odessa, 1888) investigated the diffusion of
       common salt by means of glass globules of definite density. Having
       poured water into a cylinder over a layer of a solution of sodium
       chloride, he observed during a period of several months the
       position (height) of the globules, which floated up higher and
       higher as the salt permeated upwards. Umoff found that at a
       constant temperature the distances of the globules (that is, the
       length of a column limited by layers of definite concentration)
       remain constant; that at a given moment of time the concentration,
       _q_, of different layers situated at a depth _z_ is expressed by
       the equation B-K_z_ = log.(A-_q_), where A, B, and K are
       constants; that at a given moment the rate of diffusion of the
       different layers is proportional to their depth, &_c._

  [18] If _S__{0} be the specific gravity of water, and _S_ the specific
       gravity of a solution containing _p_ p.c. of salt, then by mixing
       equal weights of water and the solution, we shall obtain a
       solution containing 1/2_p_ of the salt, and if it be formed
       without contraction, then its specific gravity _x_ will be
       determined by the equation 2/_x_ = 1/S_{0} + 1/S, because the
       volume is equal to the weight divided by the density. In reality,
       the specific gravity is always found to be greater than that
       calculated on the supposition of an absence of contraction.

  [19] Generally the specific gravity is observed by weighing in air and
       dividing the weight in grams by the volume in cubic centimetres,
       the latter being found from the weight of water displaced, divided
       by its density at the temperature at which the experiment is
       carried out. If we call this specific gravity S_{1}, then as a
       cubic centimetre of air under the usual conditions weighs about
       0·0012 gram, the sp. gr. in a vacuum S = S_{1} + 0·0012 (S_{1}
       - 1), if the density of water = 1.

  [20] If the sp. gr. S_{2} be found directly by dividing the weight of a
       solution by the weight of water at the same temperature and in the
       same volume, then the true sp. gr. _S_ referred to water at 4° is
       found by multiplying S_{2} by the sp. gr. of water at the
       temperature of observation.

  [21] According to Schiff 100 grams of alcohol, containing _p_ p.c. by
       weight of C_{2}H_{6}O, dissolves at 15°--

                    _p_ =  10    20    40    60   80
                          28·5  22·6  13·2  5·9  1·2 grams NaCl.

Common salt gives very few compounds[22] (double salts) and these are
very readily decomposed: it is also decomposed with great difficulty and
its dissociation is unknown.[23] But it is easily decomposed, both when
fused and in solution, by the action of a galvanic current. If the dry
salt be fused in a crucible and an electric current be passed through it
by immersing carbon or platinum electrodes in it (the positive electrode
is made of carbon and the negative of platinum or mercury), it is
_decomposed_: the suffocating gas, chlorine, is liberated at the positive
pole and metallic sodium at the negative pole. Both of them act on the
excess of water at the moment of their evolution; the sodium evolves
hydrogen and forms caustic soda, and the chlorine evolves oxygen and
forms hydrochloric acid, and therefore on passing a current through a
solution of common salt metallic sodium will not be obtained--but oxygen,
chlorine, and hydrochloric acid will appear at the positive pole, and
hydrogen and caustic soda at the negative pole.[23 bis] Thus salt, like
other salts, is decomposed by the action of an electric current into a
metal and a haloid (Chapter III.) Naturally, like all other salts, it may
be formed from the corresponding base and acid with the separation of
water. In fact if we mix caustic soda (base) with hydrochloric acid
(acid), table salt is formed, NaHO + HCl = NaCl + H_{2}O.

  [22] Amongst the double salts formed by sodium chloride that obtained
       by Ditte (1870) by the evaporation of the solution remaining after
       heating sodium iodate with hydrochloric acid until chlorine ceases
       to be liberated, is a remarkable one. Its composition is
       NaIO_{3},NaCl,14H_{2}O. Rammelsberg obtained a similar (perhaps
       the same) salt in well-formed crystals by the direct reaction of
       both salts.

  [23] But it gives sodium in the flame of a Bunsen's burner (see
       Spectrum Analysis), doubtless under the reducing action of the
       elements carbon and hydrogen. In the presence of an excess of
       hydrochloric acid in the flame (when the sodium would form sodium
       chloride), no sodium is formed in the flame and the salt does not
       communicate its usual coloration.

  [23 bis] There is no doubt, however, but that chloride of sodium is
       also decomposed in its aqueous solutions with the separation of
       sodium, and that it does not simply enter into double
       decomposition with the water (NaCl + H_{2}O = NaHO + HCl). This is
       seen from the fact that when a saturated solution of NaCl is
       rapidly decomposed by an electric current, a large amount of
       chlorine appears at the anode and a sodium amalgam forms at the
       mercury cathode, which acts but slowly upon the strong solution of
       salt. Castner's process for the electrolysis of brine into
       chlorine and caustic soda is an application of this method which
       has been already worked in England on an industrial scale.

With resspect to the double decompositions of sodium chloride it should
be observed that they are most varied, and serve as means of obtaining
nearly all the other compounds of sodium and chlorine.

_The double decompositions of sodium chloride_ are almost exclusively
based on the possibility of the metal sodium being exchanged for
hydrogen and other metals. But neither hydrogen nor any other metal can
directly displace the sodium from sodium chloride. This would result
in the separation of metallic sodium, which itself displaces hydrogen
and the majority of other metals from their compounds, and is not,
so far as is known, ever separated by them. The replacement of the
sodium in sodium chloride by hydrogen and various metals can only take
place by the transference of the sodium into some other combination.
If hydrogen or a metal, M, be combined with an element X, then the
double decomposition NaCl + MX = NaX + MCl takes place. Such double
decompositions take place under special conditions, sometimes completely
and sometimes only partially, as we shall endeavour to explain. In order
to acquaint ourselves with the double decompositions of sodium chloride,
we will follow the methods actually employed in practice to procure
compounds of sodium and of chlorine from common salt. For this purpose
we will first describe the treatment of sodium chloride with sulphuric
acid for the preparation of hydrochloric acid and sodium sulphate.
We will then describe the substances obtained from hydrochloric acid
and sodium sulphate. Chlorine itself, and nearly all the compounds of
this element, may be procured from hydrochloric acid, whilst sodium
carbonate, caustic soda, metallic sodium itself and all its compounds,
may be obtained from sodium sulphate.

Even in the animal organism salt undergoes similar changes, furnishing
the sodium, alkali, and hydrochloric acid which take part in the
processes of animal life.

Its necessity as a constituent in the food both of human beings and of
animals becomes evident when we consider that both hydrochloric acid and
salts of sodium are found in the substances which are separated out from
the blood into the stomach and intestines. Sodium salts are found in the
blood and in the bile which is elaborated in the liver and acts on the
food in the alimentary canal, whilst hydrochloric acid is found in the
acid juices of the stomach. Chlorides of the metals are always found in
considerable quantities in the urine, and if they are excreted they must
be replenished in the organism; and for the replenishment of the loss,
substances containing chlorine compounds must be taken in food. Not only
do animals consume those small amounts of sodium chloride which are
found in drinking water or in plants or other animals, but experience
has shown that many wild animals travel long distances in search of salt
springs, and that domestic animals which in their natural condition do
not require salt, willingly take it, and that the functions of their
organisms become much more regular from their doing so.

_The action of sulphuric acid on sodium chloride._--If sulphuric acid
be poured over common salt, then even at the ordinary temperature, as
Glauber observed, an odorous gas, hydrochloric acid, is evolved. The
reaction which takes place consists in the sodium of the salt and the
hydrogen of the sulphuric acid changing places.

           NaCl    +  H_{2}SO_{4}  =      HCl       +   NaHSO_{4}
          Sodium      Sulphuric       Hydrochloric     Acid sodium
         chloride       acid              acid          sulphate

At the ordinary temperature this reaction is not complete, but soon
ceases. When the mixture is heated, the decomposition proceeds until,
if there be sufficient salt present, all the sulphuric acid taken is
converted into acid sodium sulphate. Any excess of acid will remain
unaltered. If 2 molecules of sodium chloride (117 parts) be taken per
molecule of sulphuric acid (98 parts), then on heating the mixture to
a moderate temperature only one-half (58·5) of the salt will suffer
change. Complete decomposition, after which neither hydrogen nor
chlorine is left in the residue, proceeds (when 117 parts of table salt
are taken per 98 parts of sulphuric acid) _at a red heat only_. Then--

           2NaCl  +  H_{2}SO_{4}  =      2HCl      +  Na_{2}SO_{4}
           Table      Sulphuric      Hydrochloric        Sodium
           salt         acid             acid           sulphate

This double decomposition is the result of the action of the acid
salt, NaHSO_{4}, first formed, on sodium chloride, for the acid salt,
since it contains hydrogen, itself acts like an acid, NaCl + NaHSO_{4}
= HCl + Na_{2}SO_{4}. By adding this equation to the first we obtain
the second, which expresses the ultimate reaction. Hence in the above
reaction, non-volatile or sparingly volatile table salt and sparingly
volatile sulphuric acid are taken, and as the result of their reaction,
after the hydrogen and sodium have exchanged places, there is obtained
non-volatile sodium sulphate and gaseous hydrochloric acid. The fact
of the latter being a gaseous substance forms the main reason for the
reaction proceeding to the very end. The mechanism of this kind of
double decomposition, and the cause of the course of the reaction, are
exactly the same as those we saw in the decomposition of nitre (Chapter
VI.) by the action of sulphuric acid. The sulphuric acid in each case
displaces the other, volatile, acid.

Not only in these two instances, but in every instance, if a volatile
acid can be formed by the substitution of the hydrogen of sulphuric
acid for a metal, then this volatile acid will be formed. From this it
may be concluded that the volatility of the acid should be considered
as the cause of the progress of the reaction; and indeed if the acid be
soluble but not volatile, or if the reaction take place in an enclosed
space where the resulting acid cannot volatilise, or at the ordinary
temperature when it does not pass into the state of elastic vapour--then
the decomposition does not proceed to the end, but only up to a certain
limit. In this respect the explanations given at the beginning of this
century by the French chemist Berthollet in his work 'Essai de Statique
Chimique' are very important. _The doctrine of Berthollet_ starts from
the supposition that the chemical reaction of substances is determined
not only by the degrees of affinity between the different parts, but
also by the relative masses of the reacting substances and by those
physical conditions under which the reaction takes place. Two substances
containing the elements MX and NY, being brought into contact with
each other, form by double decomposition the compounds MY and NX; but
the formation of these two new compounds will not proceed to the end
unless one of them is removed from the sphere of action. But it can
only be removed if it possesses different physical properties from
those of the other substances which are present with it. Either it must
be a gas while the others are liquid or solid, or an insoluble solid
while the others are liquid or soluble. The relative amounts of the
resultant substances, if nothing separates out from their intermixture,
depend only on the relative quantities of the substances MX and NY, and
upon the degrees of attraction existing between the elements M, N, X,
and Y; but however great their mass may be, and however considerable
the attractions, still in any case if nothing separates out from the
sphere of action the decomposition will presently cease, a state of
equilibrium will be established, and instead of two there will remain
four substances in the mass: namely, a portion of the original bodies
MX and NY, and a certain quantity of the newly formed substances MY and
NX, if it be assumed that neither MN or XY nor any other substances are
produced, and this may for the present[24] be admitted in the case of
the double decomposition of salts in which M and X are metals and X and
Y haloids. As the ordinary double decomposition here consists merely
in the exchange of metals, the above simplification is applicable. The
sum total of existing data concerning the double decomposition of salts
leads to the conclusion that from salts MX + NY there always arises
a certain quantity of NX and MY, as should be the case according to
Berthollet's doctrine. A portion of the historical data concerning
this subject will be afterwards mentioned, but we will at once proceed
to point out the observations made by Spring (1888) which show that
_even in a solid state_ salts are subject to a similar interchange of
metals if in a condition of sufficiently close contact (it requires
time, a finely divided state, and intimate mixture). Spring took two
non-hygroscopic salts, potassium nitrate, KNO_{3}, and well-dried sodium
acetate, C_{2}H_{3}NaO_{2}, and left a mixture of their powders for
several months in a desiccator. An interchange of metals took place, as
was seen from the fact that the resultant mass rapidly attracted the
moisture of the air, owing to the formation of sodium nitrate, NaNO_{3},
and potassium acetate, C_{2}H_{3}KO_{2}, both of which are highly
hygroscopic.[24 bis]

  [24] If MX and NY represent the molecules of two salts, and if there be
       _no third substance_ present (such as water in a solution), the
       formation of XY would also be possible; for instance, cyanogen,
       iodine, &c. are capable of combining with simple haloids, as well
       as with the complex groups which in certain salts play the part of
       haloids. Besides which the salts MX and NY or MY with NX may form
       double salts. If the number of molecules be unequal, or if the
       valency of the elements or groups contained in them be different,
       as in NaCl + H_{2}SO_{4}, where Cl is a univalent haloid and
       SO_{4} is bivalent, then the matter may be complicated by the
       formation of other compounds besides MY and NX, and when a solvent
       participates in the action, and especially if present in large
       proportion, the phenomena must evidently become still more
       complex; and this is actually the case in nature. Hence while
       placing before the reader a certain portion of the existing store
       of knowledge concerning the phenomena of double saline
       decompositions, I cannot consider the theory of the subject as
       complete, and have therefore limited myself to a few data, the
       completion of which must be sought in more detailed works on the
       subject of theoretical chemistry, without losing sight of what has
       been said above.

  [24 bis] When the mixture of potassium nitrate and sodium acetate was
       heated by Spring to 100°, it was completely fused into one mass,
       although potassium nitrate fuses at about 340° and sodium nitrate
       at about 320°.

When Berthollet enunciated his doctrine the present views of atoms and
molecules had yet to be developed, and it is now necessary to submit the
matter to examination in the light of these conceptions; we will
therefore consider the reaction of salts, taking M and N, X and Y as
equivalent to each other--that is, as capable of replacing each other 'in
toto,' as Na or K,, 1/2Ca or 1/2Mg (bivalent elements) replace hydrogen.

And since, according to Berthollet's doctrine, when _m_MX of one salt
comes into contact with _n_NY of another salt, a certain quantity _x_MY
and _x_NX is formed, there remains _m_-_x_ of the salt MX, and _n_-_x_
of the salt NY. If _m_ be greater than _n_, then the maximum interchange
could lead to _x_ = _n_, whilst from the salts taken there would be
formed _n_MY + _n_NX + (_m_-_n_)MX--that is, a portion of one only of
the salts taken would remain unchanged because the reaction could only
proceed between _n_MX and _n_NY. If _x_ were actually equal to _n_, the
mass of the salt MX would not have any influence on the _modus operandi_
of the reaction, which is equally in accordance with the teaching of
Bergmann, who supposed double reactions to be independent of the mass
and determined by affinity only. If M had more affinity for X than for
Y, and N more affinity for Y than for X, then according to Bergmann
there would be no decomposition whatever, and _x_ would equal 0. If
the affinity of M for Y and of N for X were greater than those in the
original grouping, then the affinity of M for X and of N for Y would be
overcome, and, according to Bergmann's doctrine, complete interchange
would take place--_i.e._ _x_ would equal _n_. According to Berthollet's
teaching, a distribution of M and N between X and Y will take place
in every case, not only in proportion to the degrees of affinity, but
also in proportion to the masses, so that with a small affinity and a
large mass the same action can be produced as with a large affinity
and a small mass. Therefore, (1) _x_ will always be less than _n_ and
their ratio _x_/_n_ less than unity--that is, the decomposition will be
expressed by the equation, _m_MX + _n_NY = (_m_-_x_)MX + (_n_-_x_)NY
+ _x_MY + _x_NX; (2) by increasing the mass _m_ we increase the
decomposition--that is, we increase _x_ and the ratio _x_/(_n_-_x_),
until with an infinitely large quantity m the fraction _x_/_n_ will
equal 1, and the decomposition will be complete, however small the
affinities uniting MY and NX may be; and (3) if _m_ = _n_, by taking MX
+ NY or MY + NX we arrive at one and the same system _in either case_:
(_n_-_x_)MX + (_n_-_x_)NY + _x_MY + _x_NX. These direct consequences of
Berthollet's teaching are verified by experience. Thus, for example,
a mixture of solutions of sodium nitrate and potassium chloride in
all cases has entirely the same properties as a mixture of solutions
of potassium nitrate and sodium chloride, of course on condition that
the mixed solutions are of identical elementary composition. But this
identity of properties might either proceed from one system of salts
passing entirely into the other (Bergmann's hypothesis) in conformity
with the predominating affinities (for instance, from KCl + NaNO_{3}
there might arise KNO_{3} + NaCl, if it be admitted that the affinities
of the elements as combined in the latter system are greater than in
the former); or, on the other hand, it might be because both systems by
the interchange of a portion of their elements give one and the same
state of equilibrium, as according to Berthollet's teaching. Experiment
proves the latter hypothesis to be the true one. But before citing
the most historically important experiments verifying Berthollet's
doctrine, we must stop to consider the conception _of the mass_ of the
reacting substances. Berthollet understood by mass the actual relative
quantity of a substance; but now it is impossible to understand this
term otherwise than as the number of molecules, for they act as chemical
units, and in the special case of double saline decompositions it is
better to take it as the number of equivalents. Thus in the reaction
NaCl + H_{2}SO_{4} the salt is taken in one equivalent and the acid in
two. If 2NaCl + H_{2}SO_{4} act, then the number of equivalents are
equal, and so on. The _influence of mass_ on the amount of decomposition
_x_/_n_ forms the root of Berthollet's doctrine, and therefore we will
first of all turn our attention to the establishment of this principle
in relation to the double decomposition of salts.

About 1840 H. Rose[25] showed that water decomposes metallic sulphides
like calcium sulphide, CaS, forming hydrogen sulphide, H_{2}S,
notwithstanding the fact that the affinity of hydrogen sulphide, as an
acid, for lime, CaH_{2}O_{2}, as a base, causes them to react on each
other, forming calcium sulphide and water, CaS + 2H_{2}O. Furthermore,
Rose showed that the greater the amount of water acting on the calcium
sulphide, the more complete is the decomposition. The results of this
reaction are evident from the fact that the hydrogen sulphide formed
may be expelled from the solution by heating, and that the resulting
lime is sparingly soluble in water. Rose clearly saw from this that such
feeble agents, in a chemical sense, as carbonic anhydride and water, by
acting in a mass and for long periods of time in nature on the durable
rocks, which resist the action of the most powerful acids, are able
to bring about chemical change--to extract, for example, from rocks
the bases, lime, soda, potash. The influence of the mass of water on
antimonious chloride, bismuth nitrate, &c., is essentially of the same
character. These substances give up to the water a quantity of acid
which is greater in proportion as the mass of the water acting on them
is greater.[25 bis]

  [25] H. Rose is more especially known for his having carefully studied
       and perfected several methods for the exact chemical analysis of
       many mineral substances. His predecessor in this branch of
       research was Berzelius, and his successor Fresenius.

  [25 bis] Historically the influence of the mass of water was the first
       well-observed phenomenon in support of Berthollet's teaching, and
       it should not now be forgotten. In double decompositions taking
       place in dilute solutions where the mass of water is large, its
       influence, notwithstanding the weakness of affinities, must he
       great, according to the very essence of Berthollet's doctrine.

       As explaining the action of the mass of water, the experiments of
       Pattison Muir (1879) are very instructive. These experiments
       demonstrate that the decomposition of bismuth chloride is the more
       complete the greater the relative quantity of water, and the less
       the mass of hydrochloric acid forming one of the products of the
       reaction.

Barium sulphate, BaSO_{4}, which is insoluble in water, when fused with
sodium carbonate, Na_{2}CO_{3}, gives, but not completely, barium
carbonate, BaCO_{3}, (also insoluble), and sodium sulphate, Na_{2}SO_{4}.
If a solution of sodium carbonate acts on precipitated barium sulphate,
the same decomposition is also effected (Dulong, Rose), but it is
restricted by a limit and requires time. A mixture of sodium carbonate
and sulphate is obtained in the solution and a mixture of barium
carbonate and sulphate in the precipitate. If the solution be decanted
off and a fresh solution of sodium carbonate be poured over the
precipitate, then a fresh portion of the barium sulphate passes into
barium carbonate, and so by increasing the mass of sodium carbonate it is
possible to entirely convert the barium sulphate into barium carbonate.
If a definite quantity of sodium sulphate be added to the solution of
sodium carbonate, then the latter will have no action whatever on the
barium sulphate, because then a system in equilibrium determined by the
reverse action of the sodium sulphate on the barium carbonate and by the
presence of both sodium carbonate and sulphate in the solution, is at
once arrived at. On the other hand, if the mass of the sodium sulphate in
the solution be great, then the barium carbonate is reconverted into
sulphate until a definite state of equilibrium is attained between the
two opposite reactions, producing barium carbonate by the action of the
sodium carbonate and barium sulphate by the action of the sodium
sulphate.

Another most important principle of Berthollet's teaching is the
existence of _a limit of exchange decomposition_, or _the attainment
of a state of equilibrium_. In this respect the determinations of
Malaguti (1857) are historically the most important. He took a mixture
of solutions of equivalent quantities of two salts, MX and NY, and
judged the amount of the resulting exchange from the composition of the
precipitate produced by the addition of alcohol. When, for example, zinc
sulphate and sodium chloride (ZnSO_{4} and 2NaCl) were taken, there
were produced by exchange sodium sulphate and zinc chloride. A mixture
of zinc sulphate and sodium sulphate was precipitated by an excess of
alcohol, and it appeared from the composition of the precipitate that
72 per cent. of the salts taken had been decomposed. When, however, a
mixture of solutions of sodium sulphate and zinc chloride was taken, the
precipitate presented the same composition as before--that is, about 28
per cent. of the salts taken had been subjected to decomposition. In
a similar experiment with a mixture of sodium chloride and magnesium
sulphate, 2NaCl + MgSO_{4} or MgCl_{2} + Na_{2}SO_{4}, about half of
the metals underwent the decomposition, which may be expressed by
the equation 4NaCl + 2MgSO_{4} = 2NaCl + MgSO_{4} + Na_{2}SO_{4} +
MgCl_{2} = 2Na_{2}SO_{4} + 2MgCl_{2}. A no less clear limit expressed
itself in another of Malaguti's researches when he investigated the
above-mentioned reversible reactions of the insoluble salts of barium.
When, for example, barium carbonate and sodium sulphate (BaCO_{3} +
Na_{2}SO_{4}) were taken, then about 72 per cent. of the salts were
decomposed, that is, were converted into barium sulphate and sodium
carbonate. But when the two latter salts were taken, then about 19
per cent. of them passed into barium carbonate and sodium sulphate.
Probably the end of the reaction was not reached in either case, because
this would require a considerable time and a uniformity of conditions
attainable with difficulty.

Gladstone (1855) took advantage of the colour of solutions of different
ferric salts for determining the measure of exchange between metals.
Thus a solution of ferric thiocyanate has a most intense red colour, and
by making a comparison between the colour of the resulting solutions and
the colour of solutions of known strength it was possible to judge to a
certain degree the quantity of the thiocyanate formed. This colorimetric
method of determination has an important significance as being the
first in which a method was applied for determining the composition
of a solution without the removal of any of its component parts. When
Gladstone took equivalent quantities of ferric nitrate and potassium
thiocyanate--Fe(NO_{3})_{3} + 3KCNS--only 13 per cent. of the salts
underwent decomposition. On increasing the mass of the latter salt the
quantity of ferric thiocyanate formed increased, but even when more than
300 equivalents of potassium thiocyanate were taken a portion of the
iron still remained as nitrate. It is evident that the affinity acting
between Fe and NO_{3} and between K and CNS on the one hand, is greater
than the affinity acting between Fe and CNS, together with the affinity
of K for NO_{3}, on the other hand. The investigation of the variation
of the fluorescence of quinine sulphate, as well as the variation of the
rotation of the plane of polarisation of nicotine, gave in the hands
of Gladstone many proofs of the entire applicability of Berthollet's
doctrine, and in particular demonstrated the influence of mass which
forms the chief distinctive feature of the teaching of Berthollet,
teaching little appreciated in his own time.

At the beginning of the year 1860, the doctrine of the limit of reaction
and of the influence of mass on the process of chemical transformations
received a very important support in the researches of Berthelot and
P. de Saint-Gilles on the formation of the ethereal salts RX from the
alcohols ROH and acids HX, when water is also formed. This conversion is
essentially very similar to the formation of salts, but differs in that
it proceeds slowly at the ordinary temperature, extending over whole
years, and is not complete--that is, it has a distinct limit determined
by a reverse reaction; thus an ethereal salt RX with water gives an
alcohol ROH and an acid HX--up to that limit generally corresponding
with two-thirds of the alcohol taken, if the action proceed between
molecular quantities of alcohol and acid. Thus common alcohol,
C_{2}H_{5}OH, with acetic acid, HC_{2}H_{3}O_{2}, gives the following
system rapidly when heated, or slowly at the ordinary temperature, ROH
+ HX + 2RX + 2H_{2}O, whether we start from 3RHO + 3HX or from 3RX +
3H_{2}O. The process and completion of the reaction in this instance
are very easily observed, because the quantity of free acid is easily
determined from the amount of alkali requisite for its saturation, as
neither alcohol nor ethereal salt acts on litmus or other reagent for
acids. Under the influence of an increased mass of alcohol the reaction
proceeds further. If two molecules of alcohol, RHO, be taken for every
one molecule of acetic acid, HX, then instead of 66 p.c., 83 p.c. of
the acid passes into ethereal salt, and with fifty molecules of RHO
nearly all the acid is etherised. The researches of Menschutkin in their
details touched on many important aspects of the same subject, such as
the influence of the composition of the alcohol and acid on the limit
and rate of exchange--but these, as well as other details, must be
looked for in special treatises on organic and theoretical chemistry.
In any case the study of etherification has supplied chemical mechanics
with clear and valuable data, which directly confirm the two fundamental
propositions of Berthollet; the influence of mass, and the limit of
reaction--that is, the equilibrium between opposite reactions. The study
of numerous instances of dissociation which we have already touched on,
and shall again meet with on several occasions, gave the same results.
With respect to double saline decompositions, it is also necessary to
mention the researches of Wiedemann on the decomposing action of a mass
of water on the ferric salts, which could be determined by measuring the
magnetism of the solutions, because the ferric oxide (soluble colloid)
set free by the water is less magnetic than the ferric salts.

A very important epoch in the history of Berthollet's doctrine was
attained when, in 1867, the Norwegian chemists, Guldberg and Waage,
expressed it as an algebraical formula. They defined the active mass
as the number of molecules contained in a given volume, and assumed,
as follows from the spirit of Berthollet's teaching, that the action
between the substances was equal to the product of the masses of the
reacting substances. Hence if the salts MX and NY be taken in equivalent
quantities (_m_ = 1 and _n_ = 1) and the salts MY and NX are not
added to the mixture but proceed from it, then if _k_ represent the
coefficient of the rate of the action of MX on NY and if _k_´ represent
the same coefficient for the pair MY and NX, then we shall have at the
moment when the decomposition equals x a measure of action for the first
pair: _k_(1-_x_)(1-_x_) and for the second pair _k´xx_, and a state of
equilibrium or limit will be reached when _k_(1-_x_)^2 = _k_´_x_^2,
whence the ratio _k_/_k_´ = [_x_/(1-_x_)]^2. Therefore in the case of
the action of alcohol on an acid, when _x_ = 2/3, the magnitude _k_/_k_´
= 4, that is, the reaction of the alcohol on the acid is four times as
fast as that of the ethereal salt on water. If the ratio _k_/_k_´ be
known, then the influence of mass may be easily determined from it. Thus
if instead of one molecule of alcohol two be taken, then the equation
will be _k_(2-_x_)(1-_x_) = _k´xx_, whence _x_ = 0·85 or 85 percent.,
which is close to the result of experiment. If 300 molecules of alcohol
be taken, then x proves to be approximately 100 per cent., which is also
found to be the case by experiment.[26]

  [26] From the above it follows that an excess of acid should influence
       the reaction like an excess of alcohol. It is in fact shown by
       experiment that if two molecules of acetic acid be taken to one
       molecule of alcohol, 84 p.c. of alcohol is etherified. If with a
       large preponderance of acid or of alcohol certain discrepancies
       are observed, their cause must be looked for in the incomplete
       correspondence of the conditions and external influences.

But it is impossible to subject the formation of salts to any process
directly analogous to that which is so conveniently effected in
etherification. Many efforts have, however, been made to solve the
problem of the measure of reaction in this case also. Thus, for example,
Khichinsky (1866), Petrieff (1885), and many others investigated the
distribution of metals and haloid groups in the case of one metal and
several haloids taken in excess, as acids; or conversely with an excess
of bases, the distribution of these bases with relation to an acid; in
cases where a portion of the substances forms a precipitate and a portion
remains in solution. But such complex cases, although they in general
confirm Berthollet's teaching (for instance, a solution of silver nitrate
gives some silver oxide with lead oxide, and a solution of nitrate of
lead precipitates some lead oxide under the action of silver oxide, as
Petrieff demonstrated), still, owing to the complexity of the phenomena
(for instance, the formation of basic and double salts), they cannot give
simple results. But much more instructive and complete are researches
like those made by Pattison Muir (1876), who took the simple case of the
precipitation of calcium carbonate, CaCO_{3}, from the mixture of
solutions of calcium chloride and sodium or potassium carbonate, and
found in this case that not only was the rate of action (for example, in
the case of CaCl_{2} + Na_{2}CO_{3}, 75 per cent. of CaCO_{3} was
precipitated in five minutes, 85 per cent. in thirty minutes, and 94 per
cent. in two days) determined by the temperature, relative mass, and
amount of water (a large mass of water decreases the rate), but that the
limit of decomposition was also dependent on these influences. However,
even in researches of this kind the conditions of reaction are
complicated by the non-uniformity of the media, inasmuch as a portion of
the substance is obtained or remains in the form of a precipitate, so
that the system is heterogeneous. The investigation of double saline
decompositions offers many difficulties which cannot be considered as yet
entirely overcome. Although many efforts have long since been made, the
majority of the researches were carried on in aqueous solutions, and as
water is itself a saline compound and able to combine with salts and
enter into double decomposition with them, such reactions taking place in
solutions in reality present very complex cases.[27] In this sense the
reaction between alcohols and acids is much more simple, and therefore
its significance in confirmation of Berthollet's doctrine is of
particular importance. The only cases which can be compared with these
reactions for simplicity are those exchange decompositions investigated
by G. G. Gustavson, which take place between CCl_{4} and RBr_{n} on the
one hand, and CBr_{4} and RCl_{n} on the other. This case is convenient
for investigation inasmuch as the RCl_{n} and RBr_{n} taken (such as
BCl_{3}, SiCl_{4}, TiCl_{4}, POCl_{3}, and SnCl_{4}) belong to those
substances which are decomposed by water, whilst CCl_{4} and CBr_{4} are
not decomposed by water; and therefore, by heating, for instance, a
mixture of CCl_{4} + SiBr_{4} it is possible to arrive at a conclusion as
to the amount of interchange by treating the product with water, which
decomposes the SiBr_{4} left unchanged and the SiCl_{4} formed by the
exchange, and therefore by determining the composition of the product
acted on by the water it is possible to form a conclusion as to the
amount of decomposition. The mixture was always formed with equivalent
quantities--for instance, 4BCl_{3} + 3CBr_{4}. It appeared that there was
no exchange whatever on simple intermixture, but that it proceeded
slowly, when the mixture was heated (for example, with the mixture above
mentioned at 123° 4·86 per cent. of Cl was replaced by Br after 14 days'
heating, and 6·83 per cent. after 28 days, and 10·12 per cent. when
heated at 150° for 60 days). A limit was always reached which
corresponded with that of the complemental system; in the given instance
the system 4BBr_{3} + 3CCl_{4}. In this last 89·97 per cent. of bromine
in the BBr_{3} was replaced by chlorine; that is, there were obtained
89·97 molecules of BCl_{3} and there remained 10·02 molecules of BBr_{3},
and therefore the same state of equilibrium was reached as that given by
the system 4BCl_{3} + 3CBr_{4}. Both systems gave one and the same state
of equilibrium at the limit, which is in agreement with Berthollet's
doctrine.[28]

  [27] As an example two methods may be mentioned, Thomsen's and
       Ostwald's. Thomsen (1869) applied a thermochemical method to
       exceedingly dilute solutions without taking the water into further
       consideration. He took solutions of caustic soda containing
       100H_{2}O per NaHO, and sulphuric acid containing 1/2H_{2}SO_{4} +
       100H_{2}O. In order that these solutions may be mixed in such
       quantities that atomic proportions of acid and alkali would act,
       for forty grams of caustic soda (which answers to its equivalent)
       there should be employed 49 grams of sulphuric acid, and then
       +15,689 heat units would be evolved. If the normal sodium sulphate
       so formed be mixed with _n_ equivalents of sulphuric acid, a
       certain amount of heat is absorbed, namely a quantity equal to
       (_n_.1650)/(_n_ + 0·8) heat units. An equivalent of caustic soda,
       in combining with an equivalent of nitric acid, evolves +13,617
       units of heat, and the augmentation of the amount of nitric acid
       entails an absorption of heat for each equivalent equal to -27
       units; so also in combining with hydrochloric acids +13,740 heat
       units are absorbed, and for each equivalent of hydrochloric acid
       beyond this amount there are absorbed -32 heat units. Thomsen mixed
       each one of three neutral salts, sodium sulphate, sodium chloride
       and sodium nitrate, with an acid which is not contained in it; for
       instance, he mixed a solution of sodium sulphate with a solution
       of nitric acid and determined the number of heat units then
       absorbed. An absorption of heat ensued because a normal salt was
       taken in the first instance, and the mixture of all the above
       normal salts with acid produces an absorption of heat. The amount
       of heat absorbed enabled him to obtain an insight into the process
       taking place in this mixture, for sulphuric acid added to sodium
       sulphate absorbs a considerable quantity of heat, whilst
       hydrochloric and nitric acids absorb a very small amount of heat
       in this case. By mixing an equivalent of sodium sulphate with
       various numbers of equivalents of nitric acid, Thomsen observed
       that the amount of heat absorbed increased more and more as the
       amount of nitric acid was increased; thus when HNO_{3} was taken
       per 1/2Na_{2}SO_{4}, 1,752 heat units were absorbed per equivalent
       of soda contained in the sodium sulphate. When twice as much
       nitric acid was taken, 2,026 heat units, and when three times as
       much, 2,050 heat units were absorbed. Had the double decomposition
       been complete in the case where one equivalent of nitric acid was
       taken per equivalent of Na_{2}SO_{4} then according to calculation
       from similar data there should have been absorbed -2,989 units of
       heat, while in reality only -1,752 units were absorbed. Hence
       Thomsen concluded that a displacement of only about two-thirds of
       the sulphuric acid had taken place--that is, the ratio _k_ : _k_´
       for the reaction 1/2Na_{2}SO_{4} + HNO_{3} and NaNO_{3} +
       1/2H_{2}SO_{4} is equal, as for ethereal salts, to 4. By taking
       this figure and admitting the above supposition, Thomsen found
       that for all mixtures of soda with nitric acid, and of sodium
       nitrate with sulphuric acid, the amounts of heat followed Guldberg
       and Waage's law; that is, the limit of decomposition reached was
       greater the greater the mass of acid added. The relation of
       hydrochloric to sulphuric acid gave the same results. Therefore
       the researches of Thomsen fully confirm the hypotheses of Guldberg
       and Waage and the doctrine of Berthollet.

       Thomson concludes his investigation with the words: (_a_) 'When
       equivalent quantities of NaHO, HNO_{3} (or HCl) and 1/2H_{2}SO_{4}
       react on one another in an aqueous solution, then two-thirds of
       the soda combines with the nitric and one-third with the sulphuric
       acid; (_b_) this subdivision repeats itself, whether the soda be
       taken combined with nitric or with sulphuric acid; (_c_) and
       therefore nitric acid has double the tendency to combine with the
       base that sulphuric acid has, and hence in an aqueous solution it
       is a stronger acid than the latter.'

       'It is therefore necessary,' Thomsen afterwards remarks, 'to have
       an expression indicating the tendency of an acid for the
       saturation of bases. This idea cannot be expressed by the word
       _affinity_, because by this term is most often understood that
       force which it is necessary to overcome in order to decompose a
       substance into its component parts. This force should therefore be
       measured by the amount of work or heat employed for the
       decomposition of the substance. The above-mentioned phenomenon is
       of an entirely different nature,' and Thomsen introduces the term
       _avidity_, by which he designates the tendency of acids for
       neutralisation. 'Therefore the avidity of nitric acid with respect
       to soda is twice as great as the avidity of sulphuric acid. An
       exactly similar result is obtained with hydrochloric acid, so that
       its avidity with respect to soda is also double the avidity of
       sulphuric acid. Experiments conducted with other acids showed that
       not one of the acids investigated had so great an avidity as
       nitric acid; some had a greater avidity than sulphuric acid,
       others less, and in some instances the avidity = 0.' The reader
       will naturally see clearly that the path chosen by Thomsen
       deserves to be worked out, for his results concern important
       questions of chemistry, but great faith cannot be placed in the
       deductions he has already arrived at, because great complexity of
       relations is to be seen in the very method of his investigation.
       It is especially important to turn attention to the fact that all
       the reactions investigated are reactions of double decomposition.
       In them A and B do not combine with C and distribute themselves
       according to their affinity or avidity for combination, but
       reversible reactions are induced. MX and NY give MY and NX, and
       conversely; therefore the affinity or avidity for combination is
       not here directly determined, but only the difference or relation
       of the affinities or avidities. The affinity of nitric acid not
       only for the water of constitution, but also for that serving for
       solution, is much less than that of sulphuric acid. This is seen
       from thermal data. The reaction N_{2}O_{5} + H_{2}O gives +3,600
       heat units, and the solution of the resultant hydrate, 2NHO_{3},
       in a large excess of water evolves +14,986 heat units. The
       formation of SO_{3} + H_{2}O evolves +21,308 heat units, and the
       solution of H_{2}SO_{4} in an excess of water 17,860--that is,
       sulphuric acid gives more heat in both cases. The interchange
       between Na_{2}SO_{4} and 2HNO_{3} is not only accomplished at the
       expense of the production of NaNO_{3}, but also at the expense of
       the formation of H_{2}SO_{4}, hence the affinity of sulphuric acid
       for water plays its part in the phenomena of displacement.
       Therefore in determinations like those made by Thomsen the water
       does not form a medium which is present without participating in
       the process; it also takes part in the reaction. (Compare Chapter
       IX., Note 14.)

       Whilst retaining essentially the methods of Thomsen, Ostwald
       (1876) determined the variation of the sp. gr. (and afterwards of
       volume), proceeding in the same dilute solutions, on the
       saturation of acids by bases, and in the decomposition of the
       salts of one acid by the other, and arrived at conclusions of just
       the same nature as Thomsen's. Ostwald's method will be clearly
       understood from an example. A solution of caustic soda containing
       an almost molecular (40 grams) weight per litre had a specific
       gravity of 1·04051. The specific gravities of solutions of equal
       volume and equivalent composition of sulphuric and nitric acids
       were 1·02970 and 1·03084 respectively. On mixing the solutions of
       NaHO and H_{2}SO_{4} there was formed a solution of Na_{2}SO_{4}
       of sp. gr. 1·02959; hence there ensued a decrease of specific
       gravity which we will term Q, equal to 1·04051 +
       1·02970-2(1·02959) = 0·01103. So also the specific gravity after
       mixture of the solutions of NaHO and HNO_{3} was 1·02633, and
       therefore Q = 0·01869. When one volume of the solution of nitric
       acid was added to two volumes of the solution of sodium sulphate,
       a solution of sp. gr. 1·02781 was obtained, and therefore the
       resultant decrease of sp. gr.

           Q_{1} = 2(1·02959) + 1·03084-3(1·02781) = 0·00659.

       Had there been no chemical reaction between the salts, then
       according to Ostwald's reasoning the specific gravity of the
       solutions would not have changed, and if the nitric acid had
       entirely displaced the sulphuric acid Q_{2} would be =
       0·01869-0·01103 = 0·00766. It is evident that a portion of the
       sulphuric acid was displaced by the nitric acid. But the measure
       of displacement is not equal to the ratio between Q_{1} and Q_{2},
       because a decrease of sp. gr. also occurs on mixing the solution
       of sodium sulphate with sulphuric acid, whilst the mixing of the
       solutions of sodium nitrate and nitric acid only produces a slight
       variation of sp. gr. which falls within the limits of experimental
       error. Ostwald deduces from similar data the same conclusions as
       Thomsen, and thus reconfirms the formula deduced by Guldberg and
       Waage, and the teaching of Berthollet.

       The participation of water is seen still more clearly in the
       methods adopted by Ostwald than in those of Thomsen, because in
       the saturation of solutions of acids by alkalis (which Kremers,
       Reinhold, and others had previously studied) there is observed,
       not a contraction, as might have been expected from the quantity
       of heat which is then evolved, but an expansion, of volume (a
       decrease of specific gravity, if we calculate as Ostwald did in
       his first investigations). Thus by mixing 1,880 grams of a
       solution of sulphuric acid of the composition SO_{3} + 100H_{2}O,
       occupying a volume of 1,815 c.c., with a corresponding quantity of
       a solution 2(NaHO + 5H_{2}O), whose volume = 1,793 c.c., we obtain
       not 3,608 but 3,633 c.c., an expansion of 25 c.c. per gram
       molecule of the resulting salt, Na_{2}SO_{4}. It is the same in
       other cases. Nitric and hydrochloric acids give a still greater
       expansion than sulphuric acid, and potassium hydroxide than sodium
       hydroxide, whilst a solution of ammonia gives a contraction. The
       relation to water must be considered as the cause of these
       phenomena. When sodium hydroxide and sulphuric acid dissolve in
       water they develop heat and give a vigorous contraction; the water
       is separated from such solutions with great difficulty. After
       mutual saturation they form the salt Na_{2}SO_{4}, which retains
       the water but feebly and evolves but little heat with it, i.e., in
       other words, has little affinity for water. In the saturation of
       sulphuric acid by soda the water is, so to say, displaced from a
       stable combination and passes into an unstable combination; hence
       an expansion (decrease of sp. gr.) takes place. It is not the
       reaction of the acid on the alkali, but the reaction of water,
       that produces the phenomenon by which Ostwald desires to measure
       the degree of salt formation. The water, which escaped attention,
       itself has affinity, and influences those phenomena which are
       being investigated. Furthermore, in the given instance its
       influence is very great because its mass is large. When it is not
       present, or only present in small quantities, the attraction of
       the base to the acid leads to contraction, and not expansion.
       Na_{2}O has a sp. gr. 2·8, hence its molecular volume = 22; the
       sp. gr. of SO_{3} is 1·9 and volume 41, hence the sum of their
       volumes is 63; for Na_{2}SO_{4} the sp. gr. is 2·65 and volume
       53·6, consequently there is a contraction of 10 c.c. per
       gram-molecule of salt. The volume of H_{2}SO_{4} = 53·3, that of
       2NaHO = 37·4; there is produced 2H_{2}O, volume = 36, +
       Na_{2}SO_{4}, volume = 53·6. There react 90·7 c.c., and on
       saturation there result 89·6 c.c.; consequently contraction again
       ensues, although less, and although this reaction is one of
       substitution and not of combination. Consequently the phenomena
       studied by Ostwald depend but little on the measure of the
       reaction of the salts, and more on the relations of the dissolved
       substances to water. In substitutions, for instance 2NaNO_{3} +
       H_{2}SO_{4} = 2HNO_{3} + Na_{2}SO_{4}, the volumes vary but
       slightly: in the above example they are 2(38·8) + 53·3 and 2(41·2)
       + 53·6; hence 131 volumes act, and 136 volumes are produced. It
       may be concluded, therefore, on the basis of what has been said,
       that on taking water into consideration the phenomena studied by
       Thomsen and Ostwald are much more complex than they at first
       appear, and that this method can scarcely lead to a correct
       interpretation as to the distribution of acids between bases. We
       may add that P. D. Chroustcheff (1890) introduced a new method for
       this class of research, by investigating the electro-conductivity
       of solutions and their mixtures, and obtained remarkable results
       (for example, that hydrochloric acid almost entirely displaces
       formic acid and only 2/3 of sulphuric acid), but details of these
       methods must be looked for in text-books of theoretical
       chemistry.

  [28] G. G. Gustavson's researches, which were conducted in the
       laboratory of the St. Petersburg University in 1871-72, are among
       the first in which the measure of the affinity of the elements for
       the halogens is recognised with perfect clearness in the limit of
       substitution and in the rate of reaction. The researches conducted
       by A. L. Potilitzin (of which mention will be made in Chapter XI.,
       Note 66) in the same laboratory touch on another aspect of the
       same problem which has not yet made much progress, notwithstanding
       its importance and the fact that the theoretical side of the
       subject (thanks especially to Guldberg and Van't Hoff) has since
       been rapidly pushed forward. If the researches of Gustavson took
       account of the influence of mass, and were more fully supplied
       with data concerning velocities and temperatures, they would be
       very important, because of the great significance which the case
       considered has for the understanding of double saline
       decompositions in the absence of water.

       Furthermore Gustavson showed that the greater the atomic weight of
       the element (B, Si, Ti, As, Sn) combined _with chlorine_ the
       greater the amount of chlorine replaced by bromine by the action
       of CBr_{4}, and consequently the less the amount of bromine
       replaced by chlorine by the action of CCl_{4} on bromine
       compounds. For instance, for chlorine compounds the percentage of
       substitution (at the limit) is--

             BCl_{3}   SiCl_{4}   TiCl_{4}   AsCl_{3}   SnCl_{4}
              10·1       12·5       43·6       71·8       77·5

       It should he observed, however, that Thorpe, on the basis of his
       experiments, denies the universality of this conclusion. I may
       mention one conclusion which it appears to me may be drawn from
       the above-cited figures of Gustavson, if they are subsequently
       verified even within narrow limits. If CBr_{4} be heated with
       RCl_{4}, then an exchange of the bromine for chlorine takes place.
       But what would be the result if it were mixed with CCl_{4}?
       Judging by the magnitude of the atomic weights, B = 11, C = 12, Si
       = 28, about 11 p.c. of the chlorine would be replaced by bromine.
       But to what does this point? I think that this shows the existence
       of a motion of the atoms in the molecule. The mixture of CCl_{4}
       and CBr_{4} does not remain in a condition of static equilibrium;
       not only are the molecules contained in it in a state of motion,
       but also the atoms in the molecules, and the above figures show
       the measure of their translation under these conditions. The
       bromine in the CBr_{4} is, _within the limit_, substituted by the
       chlorine of the CCl_{4} in a quantity of about 11 out of 100: that
       is, a portion of the atoms of bromine previously to this moment in
       combination with one atom of carbon pass over to the other atom of
       carbon, and the chlorine passes over from this second atom of
       carbon to replace it. Therefore, also, in the homogeneous mass
       CCl_{4} all the atoms of Cl do not remain constantly combined with
       the same atoms of carbon, and _there is on exchange of atoms
       between different molecules in a homogeneous medium also_. This
       hypothesis may in my opinion explain certain phenomena of
       dissociation, but though mentioning it I do not consider it worth
       while to dwell upon it. I will only observe that a similar
       hypothesis suggested itself to me in my researches on solutions,
       and that Pfaundler enunciated an essentially similar hypothesis,
       and in recent times a like view is beginning to find favour with
       respect to the electrolysis of saline solutions.

Thus we now find ample confirmation from various quarters for the
following rules of Berthollet, applying them to double saline
decompositions: 1. From two salts MX and NY containing different haloids
and metals there result from their reaction two others, MY and NX, but
such a substitution will not proceed to the end unless one product passes
from the sphere of action. 2. This reaction is limited by the existence
of an equilibrium between MX, NY, MY, and NX, because a reverse reaction
is quite as possible as the direct reaction. 3. This limit is determined
both by the measure of the active affinities and by the relative masses
of the substances as measured by the number of the reacting molecules. 4.
Other conditions being constant, the chemical action is proportional to
the product of the chemical masses in action.[29]

  [29] Berthollet's doctrine is hardly at all affected in principle by
       showing that there are cases in which there is no decomposition
       between salts, because the affinity may be so small that even a
       large mass would still give no observable displacements. The
       fundamental condition for the application of Berthollet's
       doctrine, as well as Deville's doctrine of dissociation, lies in
       the reversibility of reactions. There are practically irreversible
       reactions (for instance, CCl_{4} + 2H_{2}O = CO_{2} + 4HCl), just
       as there are non-volatile substances. But while accepting the
       doctrine of reversible reactions and retaining the theory of the
       evaporation of liquids, it is possible to admit the existence of
       non-volatile substances, and in just the same way of reactions,
       without any visible conformity to Berthollet's doctrine. This
       doctrine evidently comes nearer than the opposite doctrine of
       Bergmann to solving the complex problems of chemical mechanics for
       the successful solution of which at the present time the most
       valuable help is to be expected from the working out of data
       concerning dissociation, the influence of mass, and the
       equilibrium and velocity of reactions. But it is evident that from
       this point of view we must not regard a solvent as a
       non-participant space, but must take into consideration the
       chemical reactions accompanying solution, or else bring about
       reactions without solution.

Thus if the salts MX and NY after reaction partly formed salts MY and NX,
then a state of equilibrium is reached and the reaction ceases; but if
one of the resultant compounds, in virtue of its physical properties,
passes from the sphere of action of the remaining substances, then the
reaction will continue. This exit from the sphere of action depends on
the physical properties of the substance and on the conditions under
which the reaction takes place. Thus, for instance, the salt NX may, in
the case of reaction between solutions, separate as a precipitate, an
insoluble substance, while the other three substances remain in solution,
or it may pass into vapour, and in this manner also pass away from the
sphere of action of the remaining substances. Let us now suppose that it
passes away in some form or other from the sphere of action of the
remaining substances--for instance, that it is transformed into a
precipitate or vapour--then a fresh reaction will set in and a
re-formation of the salt NX. If this be removed, then, although the
quantity of the elements N and X in the mass will be diminished, still,
according to Berthollet's law, a certain amount of NX should be again
formed. When this substance is again formed, then, owing to its physical
properties, it will again pass away; hence the reaction, in consequence
of the physical properties of the resultant substances, is able to
proceed to completion notwithstanding the possible weakness of the
attraction existing between the elements entering into the composition of
the resultant substance NX. Naturally, if the resultant substance is
formed of elements having a considerable degree of affinity, then the
complete decomposition is considerably facilitated.

Such a representation of the _modus operandi_ of chemical
transformations is applicable with great clearness to a number of
reactions studied in chemistry, and, what is especially important, the
application of this aspect of Berthollet's teaching does not in any way
require the determination of the measure of affinity acting between the
substances present. For instance, the action of ammonia on solutions of
salts; the displacement, by its means, of basic hydrates insoluble in
water; the separation of volatile nitric acid by the aid of non-volatile
sulphuric acid, as well as the decomposition of common salt by means of
sulphuric acid, when gaseous hydrochloric acid is formed--may be taken
as examples of reactions which proceed to the end, inasmuch as one of
the resultant substances is entirely removed from the sphere of action,
but they in no way indicate the measure of affinity.[30]

  [30] Common salt not only enters into double decomposition with acids
       but also _with every salt_. However, as clearly follows from
       Berthollet's doctrine, this form of decomposition will only in a
       few cases render it possible for new metallic chlorides to be
       obtained, because the decomposition will not be carried on to the
       end unless the metallic chloride formed separates from the mass of
       the active substances. Thus, for example, if a solution of common
       salt be mixed with a solution of magnesium sulphate, double
       decomposition ensues, but not completely, because all the
       substances remain in the solution. In this case the decomposition
       must result in the formation of sodium sulphate and magnesium
       chloride, substances which are soluble in water; nothing is
       disengaged, and therefore the decomposition 2NaCl + MgSO_{4} =
       MgCl_{2} + Na_{2}SO_{4} cannot proceed to the end. However, the
       sodium sulphate formed in this manner may be separated by freezing
       the mixture. The complete separation of the sodium sulphate will
       naturally not take place, owing to a portion of the salt remaining
       in the solution. Nevertheless, this kind of decomposition is made
       use of for the preparation of sodium sulphate from the residues
       left after the evaporation of sea-water, which contain a mixture
       of magnesium sulphate and common salt. Such a mixture is found at
       Stassfurt in a natural form. It might be said that this form of
       double decomposition is only accomplished with a change of
       temperature; but this would not be true, as may be concluded from
       other analogous cases. Thus, for instance, a solution of copper
       sulphate is of a blue colour, while a solution of copper chloride
       is green. If we mix the two salts together the green tint is
       distinctly visible, so that by this means the presence of the
       copper chloride in the solution of copper sulphate is clearly
       seen. If now we add a solution of common salt to a solution of
       copper sulphate, a green coloration is obtained, which indicates
       the formation of copper chloride. In this instance it is not
       separated, but it is immediately formed on the addition of common
       salt, as it should be according to Berthollet's doctrine.

       The complete formation of a metallic chloride from common salt can
       only occur, judging from the above, when it separates from the
       sphere of action. The salts of silver are instances in point,
       because the silver chloride is insoluble in water; and therefore
       if we add a solution of sodium chloride to a solution of a silver
       salt, silver chloride and the sodium salt of that acid which was
       in the silver salt are formed.

As a proof that double decompositions like the above are actually
accomplished in the sense of Berthollet's doctrine, the fact may be cited
that common salt may be entirely decomposed by nitric acid, and nitre may
be completely decomposed by hydrochloric acid, just as they are
decomposed by sulphuric acid; but this only takes place when, in the
first instance, an excess of nitric acid is taken, and in the second
instance, an excess of hydrochloric acid, for a given quantity of the
sodium salt, and when the resultant acid passes off. If sodium chloride
be put into a porcelain evaporating basin, nitric acid added to it, and
the mixture heated, then both hydrochloric and nitric acids are expelled
by the heat. Thus the nitric acid partially acts on the sodium chloride,
but on heating, as both acids are volatile, they are both converted into
vapour; and therefore the residue will contain a mixture of a certain
quantity of the sodium chloride taken and of the sodium nitrate formed.
If a fresh quantity of nitric acid be then added, reaction will again set
in, a certain portion of hydrochloric acid is again evolved, and on
heating is expelled together with nitric acid. If this be repeated
several times, it is possible to expel all the hydrochloric acid, and to
obtain sodium nitrate only in the residue. If, on the contrary, we take
sodium nitrate and add hydrochloric acid to it in an aqueous solution, a
certain quantity of the hydrochloric acid displaces a portion of the
nitric acid, and on heating the excess of hydrochloric acid passes away
with the nitric acid formed. On repeating this process, it is possible to
displace the nitric acid with an excess of hydrochloric acid, just as it
was possible to displace the hydrochloric acid by an excess of nitric
acid. The influence of the mass of the substance in action and the
influence of volatility are here very distinctly seen. Hence it may be
affirmed that sulphuric acid does not displace hydrochloric acid because
of an especially high degree of affinity, but that this reaction is only
carried on to the end because the sulphuric acid is not volatile, whilst
the hydrochloric acid which is formed is volatile.

The preparation of hydrochloric acid in the laboratory and on a large
scale is based upon these data. In the first instance, an excess of
sulphuric acid is employed in order that the reaction may proceed easily
at a low temperature, whilst on a large scale, when it is necessary to
economise every material, equivalent quantities are taken in order to
obtain the normal salt Na_{2}SO_{4} and not the acid salt, which would
require twice as much acid. The hydrochloric acid evolved is a gas which
is very soluble in water. It is most frequently used in practice in this
state of solution under the name of _muriatic acid_.[31]

  [31] The apparatus shown in fig. 46 (Chapter VI., Note 12) is generally
       employed for the preparation of small quantities of hydrochloric
       acid. Common salt is placed in the retort; the salt is generally
       previously fused, as it otherwise froths and boils over in the
       apparatus. When the apparatus is placed in order sulphuric acid
       mixed with water is poured down the thistle funnel into the
       retort. Strong sulphuric acid (about half as much again as the
       weight of the salt) is usually taken, and it is diluted with a
       small quantity of water (half) if it be desired to <DW44> the
       action, as in using strong sulphuric acid the action immediately
       begins with great vigour. The mixture, at first without the aid of
       heat and then at a moderate temperature (in a water-bath), evolves
       hydrochloric acid. Commercial hydrochloric acid contains many
       impurities; it is usually purified by distillation, the middle
       portions being collected. It is purified from arsenic by adding
       FeCl_{2}, distilling, and rejecting the first third of the
       distillate. If free hydrochloric acid gas be required, it is
       passed through a vessel containing strong sulphuric acid to dry
       it, and is collected over a mercury hath.

       Phosphoric anhydride absorbs hydrogen chloride (Bailey and Fowler,
       1888; 2P_{2}O_{3} + 3HCl = POCl_{3} + 3HPO_{3}) at the ordinary
       temperature, and therefore the gas cannot he dried by this
       substance.

[Illustration: FIG. 65.--Section of a salt-cake furnace. B, pan in which
the sodium chloride and sulphuric acid are first mixed and heated. C,
muffle for the ultimate decomposition.]

In chemical works the decomposition of sodium chloride by means of
sulphuric acid is carried on on a very large scale, chiefly with a
view to the preparation of normal sodium sulphate, the hydrochloric
acid being a bye-product.[31 bis] The furnace employed is termed a
_salt cake furnace_. It is represented in fig. 65, and consists of the
following two parts: the pan B and the roaster C, or enclosed space
built up of large bricks _a_ and enveloped on all sides by the smoke
and flames from the fire grate, F. The ultimate decomposition of the
salt by the sulphuric acid is accomplished in the roaster. But the first
decomposition of sodium chloride by sulphuric acid does not require
so high a temperature as the ultimate decomposition, and is therefore
carried on in the front and cooler portion, B, whose bottom is heated by
gas flues. When the reaction in this portion ceases and the evolution of
hydrochloric acid stops, then the mass, which contains about half of the
sodium chloride still undecomposed, and the sulphuric acid in the form
of acid sodium sulphate, is removed from B and thrown into the roaster
C, where the action is completed. Normal sodium sulphate, which we
shall afterwards describe, remains in the roaster. It is employed both
directly in the manufacture of glass, and in the preparation of other
sodium compounds--for instance, in the preparation of soda ash, as will
afterwards be described. For the present we will only turn our attention
to the hydrochloric acid evolved in B and C.

  [31 bis] In chemical works where sulphuric acid of 60° Baumé (22 p.c.
       of water) is employed, 117 parts of sodium chloride are taken to
       about 125 parts of sulphuric acid.

The hydrochloric acid gas evolved is subjected to condensation by
dissolving it in water.[32] If the apparatus in which the decomposition
is accomplished were hermetically closed, and only presented one outlet,
then the escape of the hydrochloric acid would only proceed through the
escape pipe intended for this purpose. But as it is impossible to
construct a perfectly hermetically closed furnace of this kind, it is
necessary to increase the draught by artificial means, or to oblige the
hydrochloric acid gas to pass through those arrangements in which it is
to be condensed. This is done by connecting the ends of the tubes through
which the hydrochloric acid gas escapes from the furnace with high
chimneys, where a strong draught is set up from the combustion of the
fuel. This causes a current of hydrochloric acid gas to pass through the
absorbing apparatus in a definite direction. Here it encounters a current
of water flowing in the opposite direction, by which it is absorbed. It
is not customary to cause the acid to pass through the water, but only to
bring it into contact with the surface of the water. The absorption
apparatus consists of large earthenware vessels having four orifices, two
above and two lateral ones in the wide central portion of each vessel.
The upper orifices serve for connecting the vessels together, and the
hydrochloric acid gas escaping from the furnace passes through these
tubes. The water for absorbing the acid enters at the upper, and flows
out from the lower, vessel, passing through the lateral orifices in the
vessels. The water flows from the chimney towards the furnace and it is
therefore evident that the outflowing water will be the most saturated
with acid, of which it actually contains about 20 per cent. The
absorption in these vessels is not complete. The ultimate absorption of
the hydrochloric acid is carried on in the so-called _coke towers_, which
usually consist of two adjacent chimneys. A lattice-work of bricks is
laid on the bottom of these towers, on which coke is piled up to the top
of the tower. Water, distributing itself over the coke, trickles down to
the bottom of the tower, and in so doing absorbs the hydrochloric acid
gas rising upwards.

  [32] As in works which treat common salt in order to obtain sodium
       sulphate, the hydrochloric acid is sometimes held to be of no
       value, it might be allowed to escape with the waste furnace gases
       into the atmosphere, which would greatly injure the air of the
       neighbourhood and destroy all vegetation. In all countries,
       therefore, there are laws forbidding the factories to proceed in
       this manner, and requiring the absorption of the hydrochloric acid
       by water at the works themselves, and not permitting the solution
       to be run into rivers and streams, whose waters it would spoil. It
       may be remarked that the absorption of hydrochloric acid presents
       no particular difficulties (the absorption of sulphurous acid is
       much more difficult) because hydrochloric acid has a great
       affinity for water and gives a hydrate which boils above 100°.
       Hence, even steam and hot water, as well as weaker solutions, can
       be used for absorbing the acid. However, Warder (1888) showed that
       weak solutions of composition H_{2}O + _n_HCl when boiled (the
       residue will be almost HCl,8H_{2}O) evolve (not water but) a
       solution of the composition H_{2}O + 445_n_^{4}HCl; for example,
       on distilling HCl,10H_{2}O, HCl,23H_{2}O is first obtained in the
       distillate. As the strength of the residue becomes greater, so
       also does that of the distillate, and therefore in order to
       completely absorb hydrochloric acid it is necessary in the end to
       have recourse to water.

       As in Russia the manufacture of sodium sulphate from sodium
       chloride has not yet been sufficiently developed, and as
       hydrochloric acid is required for many technical purposes (for
       instance, for the preparation of zinc chloride, which is employed
       for soaking railway sleepers), therefore salt is often treated
       mainly for the manufacture of hydrochloric acid.

It will be readily understood that hydrochloric acid may be obtained
from all other metallic chlorides.[33] It is frequently formed in other
reactions, many of which we shall meet with in the further course of
this work. It is, for instance, formed by the action of water on sulphur
chloride, phosphorus chloride, antimony chloride, &c.

  [33] Thus the metallic chlorides, which are decomposed to a greater or
       less degree by water, correspond with feeble bases. Such are, for
       example, MgCl_{2}, AlCl_{3}, SbCl_{3}, BiCl_{3}. The decomposition
       of magnesium chloride (and also carnallite) by sulphuric acid
       proceeds at the ordinary temperature; water decomposes MgCl_{2} to
       the extent of 50 p.c. when aided by heat, and _may be employed_ as
       a convenient _method for the production of hydrochloric acid_.
       Hydrochloric acid is also produced by the ignition of certain
       metallic chlorides in a stream of hydrogen, especially of those
       metals which are easily reduced and difficultly oxidised--for
       instance, silver chloride. Lead chloride, when heated to redness
       in a current of steam, gives hydrochloric acid and lead oxide. The
       multitude of the cases of formation of hydrochloric acid are
       understood from the fact that it is a substance which is
       comparatively very stable, resembling water in this respect, and
       even most probably more stable than water, because, at a high
       temperature and even under the action of light, chlorine
       decomposes water, with the formation of hydrochloric acid. The
       combination of chlorine and hydrogen also proceeds by their direct
       action, as we shall afterwards describe.

_Hydrochloric acid_ is a colourless gas having a pungent suffocating
odour and an acid taste. This gas fumes in air and attracts moisture,
because it forms vapour containing a compound of hydrochloric acid and
water. Hydrochloric acid is liquefied by cold, and under a pressure of
40 atmospheres, into a colourless liquid of sp. gr. 0·908 at 0°,[34]
boiling point -35° and absolute boiling point +52°. We have already
seen (Chapter I.) that hydrochloric acid combines very energetically
_with water_, and in so doing evolves a considerable amount of heat.
The solution saturated in the cold attains a density 1·23. On heating
such a solution containing about 45 parts of acid per 100 parts, the
hydrochloric acid gas is expelled with only a slight admixture of
aqueous vapour. But it is impossible to entirely separate the whole of
the hydrochloric acid from the water by this means, as could be done
in the case of an ammoniacal solution. The temperature required for
the evolution of the gas rises and reaches 110°-111°, and after this
remains constant--that is, a solution having a constant boiling point
is obtained (as with HNO_{3}), which, however, does not (Roscoe and
Dittmar) present a constant composition under different pressures,
because the hydrate is decomposed in distillation, as is seen from the
determinations of its vapour density (Bineau). Judging from the facts
(1) that with decrease of the pressure under which the distillation
proceeds the solution of constant boiling point approaches to a
composition of 25 p.c. of hydrochloric acid,[35] (2) that by passing
a stream of dry air through a solution of hydrochloric acid there is
obtained in the residue a solution which also approaches to 25 p.c. of
acid, and more nearly as the temperature falls,[36] (3) that many of the
properties of solutions of hydrochloric acid vary distinctly according
as they contain more or less than 25 p.c. of hydrochloric acid (for
instance, antimonious sulphide gives hydrogen sulphide with a stronger
acid, but is not acted on by a weaker solution, also a stronger solution
fumes in the air, &c.), and (4) that the composition HCl,6H_{2}O
corresponds with 25·26 p.c. HCl--judging from all these data, and also
from the loss of tension which occurs in the combination of hydrochloric
acid with water, it may be said that they form a _definite hydrate_ of
the composition HCl,6H_{2}O. Besides this hydrate there exists also a
crystallo-hydrate, HCl,2H_{2}O,[37] which is formed by the absorption of
hydrochloric acid by a saturated solution at a temperature of -23°. It
crystallises and melts at -18°.[38]

  [34] According to Ansdell (1880) the sp. gr. of liquid hydrochloric
       acid at 0° = 0·908, at 11·67° = 0·854, at 22·7° = 0·808, at 33° =
       0·748. Hence it is seen that the expansion of this liquid is
       greater than that of gases (Chapter II., Note 34).

  [35] According to Roscoe and Dittmar at a pressure of three atmospheres
       the solution of constant boiling point contains 18 p.c. of
       hydrogen chloride, and at a pressure of one-tenth atmosphere 23
       p.c. The percentage is intermediate at medium pressures.

  [36] At 0° 25 p.c., at 100° 20·7 p.c.; Roscoe and Dittmar.

  [37] This crystallo-hydrate (obtained by Pierre and Puchot, and
       investigated by Roozeboom) is analogous to NaCl,2H_{2}O. The
       crystals HCl,2H_{2}O at -22° have a specific gravity 1·46; the
       vapour tension (under dissociation) of the solution having a
       composition HCl,2H_{2}O at -24° = 760, at -19° = 1,010, at -18° =
       1,057, at -17° = 1,112 mm. of mercury. In a solid state the
       crystallo-hydrate at -17·7° has the same tension, whilst at lower
       temperatures it is much less: at -24° about 150, at -19° about 580
       mm. A mixture of fuming hydrochloric acid with snow reduces the
       temperature to -38°. If another equivalent of water be added to the
       hydrate HCl,2H_{2}O at -18°, the temperature of solidification
       falls to -25°, and the hydrate HCl,3H_{2}O is formed (Pickering,
       1893).

  [38] According to Roscoe at 0° one _hundred_ grams of water at a
       pressure _p_ (in millimetres of mercury) dissolves--

                   _p_ = 100   200   300   500   700   1,000
             Grams HCl   65·7  70·7  73·8  78·2  81·7  85·6

       At a pressure of 760 millimetres and temperature _t_, one
       _hundred_ grams of water dissolves

                   _t_ =  0     8°   16°   24°   40°   60°
             Grams HCl   82·5  78·3  74·2  70·0  63·3  56·1

       Roozeboom (1886) showed that at _t_° solutions containing _c_
       grams of hydrogen chloride per 100 grams of water may (with the
       variation of the pressure _p_) be formed together with the
       crystallo-hydrate HCl,2H_{2}O:

             _t_ = -28°·8  -21°   -19°   -18°   -17°·7
             _c_ =  84·2    86·8   92·6   98·4   101·4
             _p_ =   --     334    580    900    1,073  mm.

       The last combination answers to the melted crystallo-hydrate
       HCl,2H_{2}O, which splits up at temperatures above -17°·7, and at
       a constant atmospheric pressure when there are no crystals--

                   _t_ =  -24°  -21°  -18°  -10°    0°
                   _c_ =  101·2  98·3  95·7  89·8  84·2

       From these data it is seen that the hydrate HCl,2H_{2}O can exist
       in a liquid state, which is not the case for the hydrates of
       carbonic and sulphurous anhydrides, chlorine, &c.

       According to Marignac, the specific heat _c_ of a solution HCl +
       _m_H_{2}O (at about 30°, taking the specific heat of water = 1) is
       given by the expression--

          C(36·5 + _m_18) = 18_m_ - 28·39 + 140/_m_ - 268/_m_^2

       if _m_ be not less than 6·25. For example, for HCl + 25H_{2}O,
       C = 0·877.

       According to Thomsen's data, the amount of heat _Q_, expressed in
       thousands of calories, evolved in the solution of 36·5 grams of
       gaseous hydrochloric acid in _m_H_{2}O or 18_m_ grams of water is
       equal to--

                   _m_ =  2     4    10    50    400
                   _Q_ = 11·4  14·3  16·2  17·1   17·3

       In these quantities the latent heat of liquefaction is included,
       which must be taken as 5-9 thousand calories per molecular
       quantity of hydrogen chloride.

       The researches of Scheffer (1888) on the rate of diffusion (in
       water) of solutions of hydrochloric acid show that the coefficient
       of diffusion _k_ decreases with the amount of water _n_, if the
       composition of the solution is HCl,_n_H_{2}O at 0°:--

               _n_ = 5     6·9   9·8   14    27·1  129·5
               _k_ = 2·31  2·08  1·86   1·67  1·52   1·39

       It also appears that strong solutions diffuse more rapidly into
       dilute solutions than into water.

The mean specific gravities at 15°, taking water at its maximum density
(4°) as 10,000, for solutions containing _p_ per cent. of hydrogen
chloride are--

                     _p_    _S_         _p_    _S_
                      5   10,242        25   11,266
                     10   10,490        30   11,522
                     15   10,744        35   11,773
                     20   11,001        40   11,997

The formula _S_ = 9,991·6 + 49·43_p_ + 0·0571_p_^2, up to _p_ =
25·26, which answers to the hydrate HCl,6H_{2}O mentioned above,
gives the specific gravity. Above this percentage _S_ = 9,785·1 +
65·10_p_-0·240_p_^2. The rise of specific gravity with an increase
of percentage (or the differential _ds/dp_) reaches a maximum at
about 25 p.c.[39] The intermediate solution, HCl,6H_{2}O, is further
distinguished by the fact that the variation of the specific gravity
with the variation of temperature is a constant quantity, so that the
specific gravity of this solution is equal to 11,352·7(1-0·000447_t_),
where 0·000447 is the coefficient of expansion of the solution.[40]
In the case of more dilute solutions, as with water, the specific
gravity per 1° (or the differential _ds_/_dt_) rises with a rise of
temperature.[41]

                             _p_ =  0      5   10   15   20
              _S__{0} - _S__{15} =  7·2   23   38   52   64
             _S__{15} - _S__{30} = 34·1   42   50   59   67

Whilst for solutions which contain a greater proportion of hydrogen
chloride than HCl,6H_{2}O, these coefficients _decrease_ with a
rise of temperature; for instance, for 30 p.c. of hydrogen chloride
_S__{0}-_S__{15} = 88 and _S__{15}-_S__{30} = 87 (according to
Marignac's data). In the case of HCl,6H_{2}O these differences are
constant, and equal 76.

  [39] If it be admitted that the maximum of the differential corresponds
       with HCl,6H_{2}O, then it might be thought that the specific
       gravity is expressed by a parabola of the third order; but such an
       admission does not give expressions in accordance with fact. This
       is all more fully considered in my work mentioned in Chapter I.,
       Note 19.

  [40] As in water, the coefficient of expansion (or the quantity _k_ in
       the expression S_{_t_} = S_{_0_}-_k_S_{_0_}_t_, or V_t_ =
       1/(1-_kt_)) attains a magnitude 0·000447 at about 48°, it might be
       thought that at 48° all solutions of hydrochloric acid would have
       the same coefficient of expansion, but in reality this is not the
       case. At low and at the ordinary temperatures the coefficient of
       expansion of aqueous solutions is greater than that of water, and
       increases with the amount of substance dissolved.

  [41] The figures cited above may serve for the direct determination of
       that variation of the specific gravity of solutions of
       hydrochloric acid with the temperature. Thus, knowing that at 15°
       the specific gravity of a 10 p.c. solution of hydrochloric acid =
       10,492, we find that at _t_° it = 10,530-_t_(2·13 + 0·027_t_).
       Whence also may be found the coefficient of expansion (Note 40).

Thus the formation of two definite hydrates, HCl,2H_{2}O and HCl,6H_{2}O,
between hydrochloric acid and water may be accepted upon the basis of
many facts. But both of them, if they occur in a liquid state, dissociate
with great facility into hydrogen chloride and water, and are completely
decomposed when distilled.

All solutions of hydrochloric acid present the properties of an
energetic acid. They not only transform blue vegetable colouring matter
into red, and disengage carbonic acid gas from carbonates, &c., but
they also entirely saturate bases, even such energetic ones as potash,
lime, &c. In a dry state, however, hydrochloric acid does not alter
vegetable dyes, and does not effect many double decompositions which
easily take place in the presence of water. This is explained by the
fact that the gaso-elastic state of the hydrochloric acid prevents its
entering into reaction. However, incandescent iron, zinc, sodium, &c.,
act on gaseous hydrochloric acid, displacing the hydrogen and leaving
half a volume of hydrogen for each volume of hydrochloric acid gas; this
reaction may serve for determining the composition of hydrochloric acid.
Combined with water hydrochloric acid acts as an acid much resembling
nitric acid[42] in its energy and in many of its reactions; however, the
latter contains oxygen, which is disengaged with great ease, and so very
frequently acts as an oxidiser, which hydrochloric acid is not capable
of doing. The majority of metals (even those which do not displace the
H from H_{2}SO_{4}, but which, like copper, decompose it to the limit
of SO_{2}) displace the hydrogen from hydrochloric acid. Thus hydrogen
is disengaged by the action of zinc, and even of copper and tin.[42 bis]
Only a few metals withstand its action; for example, gold and
platinum. Lead in compact masses is only acted on feebly, because the
lead chloride formed is insoluble and prevents the further action of the
acid on the metal. The same is to be remarked with respect to the feeble
action of hydrochloric acid on mercury and silver, because the compounds
of these metals, AgCl and HgCl, are insoluble in water. Metallic
chlorides are not only formed by the action of hydrochloric acid on the
metals, but also by many other methods; for instance, by the action of
hydrochloric acid on the carbonates, oxides, and hydroxides, and also
by the action of chlorine on metals and certain of their compounds.
Metallic chlorides have a composition MCl; for example, NaCl, KCl,
AgCl, HgCl, if the metal replaces hydrogen equivalent for equivalent,
or, as it is said, if it be monatomic or univalent. In the case of
bivalent metals, they have a composition MCl_{2}; for example, CaCl_{2},
CuCl_{2}, PbCl_{2}, HgCl_{2}, FeCl_{2}, MnCl_{2}. The composition of
the haloid salts of other metals presents a further variation; for
example, AlCl_{3}, PtCl_{4}, &c. Many metals, for instance Fe, give
several degrees of combination with chlorine (FeCl_{2}, FeCl_{3}) as
with hydrogen. In their composition the metallic chlorides differ
from the corresponding oxides, in that the O is replaced by Cl_{2},
as should follow from the law of substitution, because oxygen gives
OH_{2}, and is consequently bivalent, whilst chlorine forms HCl, and is
therefore univalent. So, for instance, ferrous oxide, FeO, corresponds
with ferrous chloride, FeCl_{2}, and the oxide Fe_{2}O_{3} with ferric
chloride, which is also seen from the origin of these compounds, for
FeCl_{2} is obtained by the action of hydrochloric acid on ferrous
oxide or carbonate and FeCl_{3} by its action on ferric oxide. In a
word, all the typical properties of acids are shown by hydrochloric
acid, and all the typical properties of salts in the metallic chlorides
derived from it. Acids and salts composed like HCl and M_{n}Cl_{2m}
without any oxygen bear the name of haloid salts; for instance, HCl is
a haloid acid, NaCl a haloid salt, chlorine a halogen. The capacity
of hydrochloric acid to give, by its action on bases, MO, a metallic
chloride, MCl_{2}, and water, is limited at high temperatures by the
reverse reaction MCl_{2} + H_{2}O = MO + 2HCl, and the more pronounced
are the basic properties of MO the feebler is the reverse action, while
for feebler bases such as Al_{2}O_{3}, MgO, &c., this reverse reaction
proceeds with ease. Metallic chlorides corresponding with the peroxides
either do not exist, or are easily decomposed with the disengagement
of chlorine. Thus there is no compound BaCl_{4} corresponding with
the peroxide BaO_{2}. Metallic chlorides having the general aspect of
salts, like their representative sodium chloride, are, as a rule, easily
fusible, more so than the oxides (for instance, CaO is infusible at a
furnace heat, whilst CaCl_{2} is easily fused) and many other salts.
Under the action of heat many chlorides are more stable than the oxides,
some can even be converted into vapour; thus corrosive sublimate,
HgCl_{2}, is particularly volatile, whilst the oxide HgO decomposes at
a red heat. Silver chloride, AgCl, is fusible and is decomposed with
difficulty, whilst Ag_{2}O is easily decomposed. The majority of the
metallic chlorides are soluble in water, but silver chloride, cuprous
chloride, mercurous chloride, and lead chloride are sparingly soluble in
water, and are therefore easily obtained as precipitates when a solution
of the salts of these metals is mixed with a solution of any chloride
or even with hydrochloric acid. The metal contained in a haloid salt
may often be replaced by another metal, or even by hydrogen, just as is
the case with a metal in an oxide. Thus copper displaces mercury from
a solution of mercuric chloride, HgCl_{2} + Cu = CuCl_{2} + Hg, and
hydrogen at a red heat displaces silver from silver chloride, 2AgCl +
H_{2} = Ag_{2} + 2HCl. These, and a whole series of similar reactions,
form the typical methods of double saline decompositions. The measure of
decomposition and the conditions under which reactions of double saline
decompositions proceed in one or in the other direction are determined
by the properties of the compounds which take part in the reaction, and
of those capable of formation at the temperature, &c., as was shown in
the preceding portions of this chapter, and as will be frequently found
hereafter.

  [42] Thus, for instance, with feeble bases they evolve in dilute
       solutions (Chapter III., Note 53) almost equal amounts of heat;
       their relation to sulphuric acid is quite identical. They both
       form fuming solutions as well as hydrates; they both form
       solutions of constant boiling point.

  [42 bis] Pybalkin (1891) found that copper begins to disengage hydrogen
       at 100°, and that chloride of copper begins to give up its
       chlorine to hydrogen gas at 230°; for silver these temperatures
       are 117° and 260°--that is, there is less difference between them.

If hydrochloric acid enters into double decomposition with basic oxides
and their hydrates, this is only due to its acid properties; and for the
same reason it rarely enters into double decomposition with acids and
acid anhydrides. Sometimes, however, it combines with the latter, as, for
instance, with the anhydride of sulphuric acid, forming the compound
SO_{3}HCl; and in other cases it acts on acids, giving up its hydrogen to
their oxygen and forming chlorine, as will be seen in the following
chapter.

Hydrochloric acid, as may already be concluded from the composition of
its molecule, belongs to the monobasic acids, and does not, therefore,
give true acid salts (like HNaSO_{4} or HNaCO_{3}); nevertheless
many metallic chlorides, formed from powerful bases, are capable of
_combining with hydrochloric acid_, just as they combine with water,
or with ammonia, or as they give double salts. Compounds have long
been known of hydrochloric acid with auric, platinic, and antimonious
chlorides, and other similar metallic chlorides corresponding with
very feeble bases. But Berthelot, Engel, and others have shown that
the capacity of HCl for combining with M_{_n_}Cl_{_m_} is much more
frequently encountered than was previously supposed. Thus, for
instance, dry hydrochloric acid when passed into a solution of zinc
chloride (containing an excess of the salt) gives in the cold (0°)
a compound HCl,ZnCl_{2},2H_{2}O, and at the ordinary temperature
HCl,2ZnCl_{2},2H_{2}O, just as it is able at low temperatures to
form the crystallo-hydrate ZnCl_{2},3H_{2}O (Engel, 1886). Similar
compounds are obtained with CdCl_{2},CuCl_{2}, HgCl_{2},Fe_{2}Cl_{6},
&c. (Berthelot, Ditte, Cheltzoff, Lachinoff, and others). These
compounds with hydrochloric acid are generally more soluble in water
than the metallic chlorides themselves, so that whilst hydrochloric
acid decreases the solubility of M_{_n_}Cl_{_m_}, corresponding with
energetic bases (for instance, sodium or barium chlorides), it increases
the solubility of the metallic chlorides corresponding with feeble bases
(cadmium chloride, ferric chloride, &c.) Silver chloride, which is
insoluble in water, is soluble in hydrochloric acid. Hydrochloric acid
also combines with certain unsaturated hydrocarbons (for instance, with
turpentine, C_{10}H_{16},2HCl) and their derivatives. _Sal-ammoniac_,
or ammonia hydrochloride, NH_{4}Cl = NH_{3},HCl, also belongs to
this class of compounds.[43] If hydrogen chloride gas be mixed with
ammonia gas a solid compound consisting of equal volumes of each is
immediately formed. The same compound is obtained on mixing solutions of
the two gases. It is also produced by the action of hydrochloric acid
on ammonium carbonate. Sal-ammoniac is usually prepared, in practice,
by the last method.[44] The specific gravity of sal-ammoniac is 1·55.
We have already seen (Chapter VI.) that sal-ammoniac, like all other
ammonium salts, easily decomposes; for instance, by volatilisation with
alkalis, and even partially when its solution is boiled. The other
properties and reactions of sal-ammoniac, especially in solution, fully
recall those already mentioned in speaking of sodium chloride. Thus, for
instance, with silver nitrate it gives a precipitate of silver chloride;
with sulphuric acid it gives hydrochloric acid and ammonium sulphate,
and it forms double salts with certain metallic chlorides and other
salts.[45]

  [43] When an unsaturated hydrocarbon, or, in general, an unsaturated
       compound, assimilates to itself the molecules Cl_{2}, HCl, SO_{3},
       H_{2}SO_{4}, &c., the cause of the reaction is most simple. As
       nitrogen, besides the type NX_{3} to which NH_{3}, belongs, gives
       compounds of the type NX_{5}--for example, NO_{2}(OH)--the
       formation of the salts of ammonium should be understood in this
       way. NH_{3} gives NH_{4}Cl because NX_{3} is capable of giving
       NX_{5}. But as saturated compounds--for instance, SO_{3},H_{2}O,
       NaCl, &c.--are also capable of combination even between
       themselves, it is impossible to deny the capacity of HCl also for
       combination. SO_{3} combines with H_{2}O, and also with HCl and
       the unsaturated hydrocarbons. It is impossible to recognise the
       distinction formerly sought to be established between atomic and
       molecular compounds, and regarding, for instance, PCl_{3} as an
       atomic compound and PCl_{5} as a molecular one, only because it
       easily splits up into molecules PCl_{3} and Cl_{2}.

  [44] Sal-ammoniac is prepared from ammonium carbonate, obtained in the
       dry distillation of nitrogenous substances (Chapter VI.), by
       saturating the resultant solution with hydrochloric acid. A
       solution of sal-ammoniac is thus produced, which is evaporated,
       and in the residue a mass is obtained containing a mixture of
       various other, especially tarry, products of dry distillation. The
       sal-ammoniac is generally purified by sublimation. For this
       purpose iron vessels covered with hemispherical metallic covers
       are employed, or else simply clay crucibles covered by other
       crucibles. The upper portion, or head, of the apparatus of this
       kind will have a lower temperature than the lower portion, which
       is under the direct action of the flame. The sal-ammoniac
       volatilises when heated, and settles on the cooler portion of the
       apparatus. It is thus freed from many impurities, and is obtained
       as a crystalline crust, generally several centimetres thick, in
       which form it is commonly sold. The solubility of sal-ammoniac
       rises rapidly with the temperature: at 0°, 100 parts of water
       dissolve about 28 parts of NH_{4}Cl, at 50° about 50 parts, and at
       the ordinary temperature about 35 parts. This is sometimes taken
       advantage of for separating NH_{4}Cl from solutions of other
       salts.

  [45] The solubility of sal-ammoniac in 100 parts of water (according to
       Alluard) is--

           0°    10°    20°    30°    40°  60°  80°  100°  110°
          28·40  32·48  37·28  41·72  46   55   64    73    77

       A saturated solution boils at 115°·8. The specific gravity at
       15°/4° of solutions of sal-ammoniac (water 4° = 10,000) =
       9,991·6-31·26_p_-0·085_p_^2, where _p_ is the amount by weight of
       ammonium chloride in 100 parts of solution. With the majority of
       salts the differential _ds_/_dp_ increases, but here it decreases
       with the increase of _p_. For (unlike the sodium and potassium
       salts) a solution of the alkali _plus_ a solution of acid occupy a
       greater volume than that of the resultant ammonium salt. In the
       solution of _solid_ ammonium chloride a contraction, and not
       expansion, generally takes place. It may further be remarked that
       solutions of sal-ammoniac have an acid reaction even when prepared
       from the salt remaining after prolonged washing of the sublimed
       salt with water (A. Stcherbakoff).




                               CHAPTER XI

          THE HALOGENS: CHLORINE, BROMINE, IODINE, AND FLUORINE


Although hydrochloric acid, like water, is one of the most stable
substances, it is nevertheless decomposed not only by the action of
a galvanic current,[1] but also by a high temperature. Sainte-Claire
Deville showed that decomposition already occurs at 1,300°, because a
cold tube (as with CO, Chapter IX.) covered with an amalgam of silver
absorbs chlorine from hydrochloric acid in a red-hot tube, and the
escaping gas contains hydrogen. V. Meyer and Langer (1885) observed
the decomposition of hydrochloric acid at 1,690° in a platinum vessel;
the decomposition in this instance was proved not only from the fact
that hydrogen diffused through the platinum (p. 142), owing to which
the volume was diminished, but also from chlorine being obtained in the
residue (the hydrogen chloride was mixed with nitrogen), which liberated
iodine from potassium iodide.[2] The usual method for the preparation
of chlorine consists in the abstraction of the hydrogen by oxidising
agents.[2 bis]

  [1] The decomposition of fused sodium chloride by an electric current
      has been proposed in America and Russia (N. N. Beketoff) as a means
      for the preparation of chlorine and sodium. A strong solution of
      hydrochloric acid is decomposed into equal volumes of chlorine and
      hydrogen by the action of an electric current. If sodium chloride
      and lead be melted in a crucible, the former being connected with
      the cathode and a carbon anode immersed in the lead, then the lead
      dissolves sodium and chlorine is disengaged as gas. This
      electrolytic method has not yet been practised on a large scale,
      probably because gaseous chlorine has not many applications, and
      because of the difficulty there is in dealing with it.

  [2] To obtain so high a temperature (at which the best kinds of
      porcelain soften) Langer and Meyer employed the dense graphitoidal
      carbon from gas retorts, and a powerful blast. They determined the
      temperature by the alteration of the volume of nitrogen in the
      platinum vessel, for this gas does not permeate through platinum,
      and is unaltered by heat.

  [2 bis] The acid properties of hydrochloric acid were known when
      Lavoisier pointed out the formation of acids by the combination of
      water with the oxides of the non-metals, and therefore there was
      reason for thinking that hydrochloric acid was formed by the
      combination of water with the oxide of some element. Hence when
      Scheele obtained chlorine by the action of hydrochloric acid on
      manganese peroxide he considered it as the acid contained in common
      salt. When it became known that chlorine gives hydrochloric acid
      with hydrogen, Lavoisier and Berthollet supposed it to be a
      compound with oxygen of an anhydride contained in hydrochloric
      acid. They supposed that hydrochloric acid contained water and the
      oxide of a particular radicle, and that chlorine was a higher
      degree of oxidation of this radicle _muvias_ (from the Latin neme
      of hydrochloric acid, _acidum muriaticum_). It was only in 1811
      that Gay-Lussac and Thénard in France and Davy in England arrived
      at the conclusion that the substance obtained by Scheele does not
      contain oxygen, nor under any conditions give water with hydrogen,
      and that there is no water in hydrochloric acid gas, and therefore
      concluded that chlorine is an elementary substance. They named it
      'chlorine' from the Greek word [Greek: chlôros], signifying a green
      colour, because of the peculiar colour by which this gas is
      characterised.

An aqueous solution of hydrochloric acid is generally employed for t he
evolution of chlorine. The hydrogen has to be abstracted from the
hydrochloric acid. This is accomplished by nearly all oxidising
substances, and especially by those which are able to evolve oxygen at a
red heat (besides bases, such as mercury and silver oxides, which are
able to give salts with hydrogen chloride); for example, manganese
peroxide, potassium chlorate, chromic acid, &c. The decomposition
essentially consists in the oxygen of the oxidising substance displacing
the chlorine from 2HCl, forming water, H_{2}O, and setting the chlorine
free, 2HCl + O (disengaged by the oxidising substances) = H_{2}O +
Cl_{2}. Even nitric acid partially produces a like reaction; but as we
shall afterwards see its action is more complicated, and it is therefore
not suitable for the preparation of pure chlorine.[3] But other oxidising
substances which do not give any other volatile products with
hydrochloric acid may be employed for the preparation of chlorine. Among
these may be mentioned: potassium chlorate, acid potassium chromate,
sodium manganate, manganese peroxide, &c. Manganese peroxide is commonly
employed in the laboratory, and on a large scale, for the preparation of
chlorine. The chemical process in this case may be represented as
follows: an exchange takes place between 4HCl and MnO_{2}, in which the
manganese takes the place of the four atoms of hydrogen, or the chlorine
and oxygen exchange places--that is, MnCl_{4} and 2H_{2}O are produced.
The chlorine compound, MnCl_{4}, obtained is very unstable; it splits up
into chlorine, which as a gas passes from the sphere of action, and a
lower compound containing less chlorine than the substance first formed,
which remains in the apparatus in which the mixture is heated, MnCl_{4} =
MnCl_{2} + Cl_{2}.[3 bis] The action of hydrochloric acid requires a
temperature of about 100°. In the laboratory the _preparation of
chlorine_ is carried on in flasks, heated over a water-bath, by acting on
manganese peroxide with hydrochloric acid or a mixture of common salt
and sulphuric acid[4] and washing the gas with water to remove
hydrochloric acid.[5] Chlorine cannot be collected over mercury, because
it combines with it as with many other metals, and it is soluble in
water; however, it is but slightly soluble in hot water or brine. Owing
to its great weight, chlorine may be directly collected in a dry vessel
by carrying the gas-conducting tube down to the bottom of the vessel. The
chlorine will lie in a heavy layer at the bottom of the vessel, displace
the air, and the extent to which it fills the vessel may be followed by
its colour.[6]

  [3] However, nitric acid has been proposed as a means for obtaining
      chlorine, but by methods which have the drawback of being very
      complicated

  [3 bis] This representation of the process of the reaction is most
      natural. However, this decomposition is generally represented as if
      chlorine gave only one degree of combination with manganese,
      MnCl_{2}, and therefore directly reacts in the following
      manner--MnO_{2} + 4HCl = MnCl_{2} + 2H_{2}O + Cl_{2}, in which case
      it is supposed that manganese peroxide, MnO_{2}, breaks up, as it
      were, into manganous oxide, MnO and oxygen, both of which react
      with hydrochloric acid, the manganous oxide acting upon HCl as a
      base, giving MnCl_{2} and at the same time 2HCl + O = H_{2}O +
      Cl_{2}. In reality, a mixture of oxygen and hydrochloric acid does
      give chlorine at a red heat, and this reaction may also take place
      at the moment of its evolution in this case.

      All the oxides of manganese (Mn_{2}O_{3}, MnO_{2}, MnO_{3},
      Mn_{2}O_{7}), with the exception of manganous oxide, MnO, disengage
      chlorine from hydrochloric acid, because manganous chloride,
      MnCl_{2}, is the only compound of chlorine and manganese which
      exists as a stable compound, all the higher chlorides of manganese
      being unstable and evolving chlorine. Hence we here take note of
      two separate changes: (1) an exchange between oxygen and chlorine,
      and (2) the instability of the higher chlorine compounds. As
      (according to the law of substitution) in the substitution of
      oxygen by chlorine, Cl_{2} takes the place of O, the chlorine
      compounds will contain more atoms than the corresponding oxygen
      compounds. It is not surprising, therefore, that certain of the
      chlorine compounds corresponding with oxygen compounds do not
      exist, or if they are formed are very unstable. And furthermore, an
      atom of chlorine is heavier than an atom of oxygen, and therefore a
      given element would have to retain a large mass of chlorine if in
      the higher oxides the oxygen were replaced by chlorine. For this
      reason equivalent compounds of chlorine do not exist for all oxygen
      compounds. Many of the former are immediately decomposed, when
      formed, with the evolution of chlorine. From this it is evident
      that there should exist such chlorine compounds as would evolve
      chlorine as peroxides evolve oxygen, and indeed a large number of
      such compounds are known. Amongst them may be mentioned antimony
      pentachloride, SbCl_{5}, which splits up into chlorine and antimony
      trichloride when heated. Cupric chloride, corresponding with copper
      oxide, and having a composition CuCl_{2}, similar to CuO, when
      heated parts with half its chlorine, just as barium peroxide
      evolves half its oxygen. This method may even be taken advantage of
      for the preparation of chlorine and cuprous chloride, CuCl. The
      latter attracts oxygen from the atmosphere, and in so doing is
      converted from a colourless substance into a green compound whose
      composition is Cu_{2}Cl_{2}O. With hydrochloric acid this substance
      gives cupric chloride (Cu_{2}Cl_{2}O + 2HCl = H_{2}O + 2CuCl_{2}),
      which has only to be dried and heated in order again to obtain
      chlorine. Thus, in solution, and at the ordinary temperature, the
      compound CuCl_{2} is stable, but when heated it splits up. On this
      property is founded Deacon's process for the preparation of
      chlorine from hydrochloric acid with the aid of air and copper
      salts, by passing a mixture of air and hydrochloric acid at about
      440° over bricks saturated with a solution of a copper salt (a
      mixture of solutions of CuSO_{4} and Na_{2}SO_{4}). CuCl_{2} is
      then formed by the double decomposition of the salt of copper and
      the hydrochloric acid; the CuCl_{2} liberates chlorine, and the
      CuCl forms Cu_{2}Cl_{2}O with the oxygen of the air, which again
      gives CuCl_{2} with 2HCl, and so on.

      Magnesium chloride, which is obtained from sea-water, carnallite,
      &c., may serve not only as a means for the preparation of
      hydrochloric acid, but also of chlorine, because its basic salt
      (magnesium oxychloride) when heated in the air gives magnesium
      oxide and chlorine (Weldon-Pechiney's process, 1888). Chlorine is
      now prepared on a large scale by this method. Several new methods
      based upon this reaction have been proposed for procuring chlorine
      from the bye-products of other chemical processes. Thus, Lyte and
      Tattars (1891) obtained up to 67 p.c. of chlorine from CaCl_{2} in
      this manner. A solution of CaCl_{2}, containing a certain amount of
      common salt, is evaporated and oxide of magnesium added to it. When
      the solution attains a density of 1·2445 (at 15°), it is treated
      with carbonic acid, which precipitates carbonate of calcium, while
      chloride of magnesium remains in solution. After adding ammonium
      chloride, the solution is evaporated to dryness and the double
      chloride of magnesium and ammonium formed is ignited, which drives
      off the chloride of ammonium. The chloride of magnesium which
      remains behind is used in the Weldon-Pechiney process. The De
      Wilde-Reychler (1892) process for the manufacture of chlorine
      consists in passing alternate currents of hot air and hydrochloric
      acid gas through a cylinder containing a mixture of the chlorides
      of magnesium and manganese. A certain amount of sulphate of
      magnesium which does not participate in any way in the reaction, is
      added to the mixture to prevent its fusing. The reactions may be
      expressed by the following equations: (1) 3MgCl_{2} + 3MnCl_{2} +
      8O = Mg_{3}Mn_{3}O_{8} + 12Cl; (2) Mg_{3}Mn_{3}O_{8} + 16HCl =
      3MgCl_{2} + 3MnCl_{2} + 8H_{2}O + 4Cl. As nitric acid is able to
      take up the hydrogen from hydrochloric acid, a heated mixture of
      these acids is also employed for the preparation of chlorine. The
      resultant mixture of chlorine and lower oxides of nitrogen is mixed
      with air and steam which regenerates the HNO_{3}, while the
      chlorine remains as a gas together with nitrogen, in which form it
      is quite capable of bleaching, forming chloride of lime, &c.
      Besides these, Solvay and Mond's methods of preparing chlorine must
      be mentioned. The first is based upon the reaction CaCl_{2} +
      SiO_{2} + O(air) = CaOSiO_{2} + Cl_{2}, the second on the action of
      the oxygen of the air (heated) upon MgCl_{2} (and certain similar
      chlorides) MgCl_{2} + O = MgO + Cl_{2} The remaining MgO is treated
      with sal-ammoniac to re-form MgCl_{2} (MgO + 2NH_{4}Cl = MgCl_{2} +
      H_{2}O + 2NH_{3}) and the resultant NH_{3} again converted into
      sal-ammoniac, so that hydrochloric acid is the only substance
      consumed. The latter processes have not yet found much application.

  [4] The following proportions are accordingly taken by weight: 5 parts
      of powdered manganese peroxide, 11 parts of salt (best fused, to
      prevent its frothing), and 14 parts of sulphuric acid previously
      mixed with an equal volume of water. The mixture is heated in a
      salt bath, so as to obtain a temperature above 100°. The corks in
      the apparatus must be soaked in paraffin (otherwise they are
      corroded by the chlorine), and black india-rubber tubing smeared
      with vaseline must be used, and not vulcanised rubber (which
      contains sulphur, and becomes brittle under the action of the
      chlorine).

      The reaction which proceeds may be expressed thus: MnO_{2} + 2NaCl
      + 2H_{2}SO_{4} = MnSO_{4} + Na_{2}SO_{4} + 2H_{2}O + Cl_{2}. The
      method of preparation of Cl_{2} from manganese peroxide and
      hydrochloric acid was discovered by Scheele, and from sodium
      chloride by Berthollet.

  [5] The reaction of hydrochloric acid upon bleaching powder gives
      chlorine without the aid of heat, CaCl_{2}O_{2} + 4HCl = CaCl_{2} +
      2H_{2}O + 2Cl_{2} and is therefore also used for the preparation of
      chlorine. This reaction is very violent if all the acid be added at
      once; it should be poured in drop by drop (Mermé, Kämmerer). C.
      Winkler proposed to mix bleaching powder with one quarter of burnt
      and powdered gypsum, and having damped the mixture with water, to
      press and cut it up into cubes and dry at the ordinary temperature.
      These cubes can be used for the preparation of chlorine in the same
      apparatus as that used for the evolution of hydrogen and carbonic
      anhydride--the disengagement of the chlorine proceeds uniformly.

      A mixture of potassium dichromate and hydrochloric acid evolves
      chlorine perfectly free from oxygen (V. Meyer and Langer).

  [6] [Illustration: FIG. 66.--Clay retort for the preparation of chlorine
      on a large scale.]

      Chlorine is manufactured on a _large scale_ from manganese peroxide
      and hydrochloric acid. It is most conveniently prepared in the
      apparatus shown in fig. 66, which consists of a three-necked
      earthenware vessel whose central orifice is the largest. A clay or
      lead funnel, furnished with a number of orifices, is placed in the
      central wide neck of the vessel. Roughly-ground lumps of natural
      manganese peroxide are placed in the funnel, which is then closed
      by the cover N, and luted with clay. One orifice is closed by a
      clay stopper, and is used for the introduction of the hydrochloric
      acid and withdrawal of the residues. The chlorine disengaged passes
      along a leaden gas-conducting tube placed in the other orifice. A
      row of these vessels is surrounded by a water-bath to ensure their
      being uniformly heated. Manganese chloride is found in the residue.
      In Weldon's process lime is added to the acid solution of manganese
      chloride. A double decomposition takes place, resulting in the
      formation of manganous hydroxide and calcium chloride. When the
      insoluble manganous hydroxide has settled, a further excess of milk
      of lime is added (to make a mixture 2Mn(OH)_{2} + CaO +
      _x_CaCl_{2}, which is found to be the best proportion, judging from
      experiment), and then air is forced through the mixture. The
      hydroxide is thus converted from a colourless to a brown substance,
      containing peroxide, MnO_{2}, and oxide of manganese, Mn_{2}O_{3}.
      This is due to the manganous oxide absorbing oxygen from the air.
      Under the action of hydrochloric acid this mixture evolves
      chlorine, because of all the compounds of chlorine and manganese
      the chloride MnCl_{2} is the only one which is stable (_see_ Note
      3). Thus one and the same mass of manganese may be repeatedly used
      for the preparation of chlorine. The same result is attained in
      other ways. If manganous oxide be subjected to the action of oxides
      of nitrogen and air (Coleman's process), then manganese nitrate is
      formed, which at a red heat gives oxides of nitrogen (which are
      again used in the process) and manganese peroxide, which is thus
      renewed for the fresh evolution of chlorine.

Chlorine is a _gas_ of a yellowish green colour, and has a very
suffocating and characteristic odour. On lowering the temperature to
-50° or increasing the pressure to six atmospheres (at 0°) chlorine
condenses[7] into a liquid which has a yellowish-green colour, a density
of 1·3, and boils at -34°. The density and atomic weight of chlorine is
35·5 times greater than that of hydrogen, hence the molecule contains
Cl_{2}[8]. At 0° one volume of water dissolves about 1-1/2 volume of
chlorine, at 10° about 3 volumes, at 50° again 1-1/2 volume.[9] Such
a solution of chlorine is termed 'chlorine water;' and is employed in
a diluted form in medicine and as a laboratory reagent. It is prepared
by passing chlorine through a series of Woulfe's bottles or into an
inverted retort filled with water. Under the action of light, chlorine
water gives oxygen and hydrochloric acid. At 0° a saturated solution of
chlorine yields a crystallo-hydrate, Cl_{2},8H_{2}O, which easily splits
up into chlorine and water when heated, so that if it be sealed up in a
tube and heated to 35°, two layers of liquid are formed--a lower stratum
of chlorine containing a small quantity of water, and an upper stratum
of water containing a small quantity of chlorine.[10]

  [7] Davy and Faraday liquefied chlorine in 1823 by heating the
      crystallo-hydrate Cl_{2}8H_{2}O in a bent tube (as with NH_{3}),
      surrounded by warm water, while the other end of the tube was
      immersed in a freezing mixture. Meselan condensed chlorine in
      freshly-burnt charcoal (placed in a glass tube), which when cold
      absorbs an equal weight of chlorine. The tube was then fused up,
      the bent end cooled, and the charcoal heated, by which means the
      chlorine was expelled from the charcoal, and the pressure
      increased.

  [8] Judging from Ludwig's observations (1868), and from the fact that
      the coefficient of expansion of gases increases with their
      molecular weight (Chapter II., Note 26, for hydrogen = 0·367,
      carbonic anhydride = 0·373, hydrogen bromide = 0·386), it might be
      expected that the expansion of chlorine would be greater than that
      of air or of the gases composing it. V. Meyer and Langer (1885)
      having remarked that at 1,400° the density of chlorine (taking its
      expansion as equal to that of nitrogen) = 29, consider that the
      molecules of chlorine split up and partially give molecules Cl, but
      it might be maintained that the decrease in density observed only
      depends on the increase of the coefficient of expansion.

  [9] Investigations on the solubility of chlorine in water (the
      solutions evolve all their chlorine on boiling and passing air
      through them) show many different peculiarities. First Gay-Lussac,
      and subsequently Pelouze, determined that the solubility increases
      between 0° and 8°-10° (from 1-1/2 to 2 vols. of chlorine per 100
      vols. of water at 0° up to 3 to 2-3/4 at 10°). In the following
      note we shall see that this is not due to the breaking-up of the
      hydrate at about 8° to 10°, but to its formation below 9°. Roscoe
      observed an increase in the solubility of chlorine in the presence
      of hydrogen--even in the dark. Berthelot determined an increase of
      solubility with the progress of time. Schönbein and others suppose
      that chlorine acts on water, forming hypochlorous and hypochloric
      acids, (HClO + HCl).

      The equilibrium between chlorine and steam as gases and between
      water, liquid chlorine, ice, and the solid crystallo-hydrate of
      chlorine is evidently very complex. Gibbs, Guldberg (1870) and
      others gave a theory for similar states of equilibrium, which was
      afterwards developed by Roozeboom (1887), but it would be
      inopportune here to enter into its details. It will be sufficient
      in the first place to mention that there is now no doubt (according
      to the theory of heat, and the direct observations of Ramsay and
      Young) that the vapour tensions at one and the same temperature are
      different for the liquid and solid states of substances; secondly,
      to call attention to the following note; and, thirdly, to state
      that, in the presence of the crystallo-hydrate, water between O°·24
      and +28°·7 (when the hydrate and a solution may occur
      simultaneously) dissolves a different amount of chlorine than it
      does in the absence of the crystallo-hydrate.

  [10] According to Faraday's data the hydrate of chlorine contains
       Cl_{2},10H_{2}O, but Roozeboom (1885) showed that it is poorer in
       water and = Cl_{2},8H_{2}O. At first small, almost colourless,
       crystals are obtained, but they gradually form (if the temperature
       be below their critical point 28°·7, above which they do not
       exist) large yellow crystals, like those of potassium chromate.
       The specific gravity is 1·23. The hydrate is formed if there be
       more chlorine in a solution than it is able to dissolve under the
       dissociation pressure corresponding with a given temperature. _In
       the presence of the hydrate_ the percentage amount of chlorine at
       0° = 0·5, at 9° = 0·9, and at 20° = 1·82. At temperatures below 9°
       the solubility (determined by Gay-Lussac and Pelouze, _see_ Note
       9) is dependent on the formation of the hydrate; whilst at higher
       temperatures under the ordinary pressure the hydrate cannot be
       formed, and the solubility of chlorine falls, as it does for all
       gases (Chapter I.). If the crystallo-hydrate is not formed, then
       below 9° the solubility follows the same rule (6° 1·07 p.c. Cl, 9°
       0·95 p.c.). According to Roozeboom, the chlorine evolved by the
       hydrate presents the following tensions of dissociation: at 0° =
       249 mm., at 4° = 398, at 8° = 620, at 10° = 797, at 14° = 1,400
       mm. In this case a portion of the crystallo-hydrate remains solid.
       At 9°·6 the tension of dissociation is equal to the atmospheric
       pressure. At a higher pressure the crystallo-hydrate may form at
       temperatures above 9° up to 28°·7, when the vapour tension of the
       hydrate equals the tension of the chlorine. It is evident that the
       equilibrium which is established is on the one hand a case of a
       complex heterogeneous system, and on the other hand a case of the
       solution of solid and gaseous substances in water.

       The crystallo-hydrate or chlorine water must be kept in the dark,
       or the access of light be prevented by  glass, otherwise
       oxygen is evolved and hydrochloric acid formed.

Chlorine explodes _with hydrogen_, if a mixture of equal volumes be
exposed to the direct action of the sun's rays[11] or brought into
contact with spongy platinum, or a strongly heated substance, or when
subjected to the action of an electric spark. The explosion in this case
takes place for exactly the same reasons--_i.e._ the evolution of heat
and expansion of the resultant product--as in the case of detonating
gas (Chapter III.) Diffused light acts in the same way, but slowly,
whilst direct sunlight causes an explosion.[12] The hydrochloric acid
gas produced by the reaction of chlorine on hydrogen occupies (at the
original temperature and pressure) a volume equal to the sum of the
original volumes; that is, a reaction of substitution here takes place:
H_{2} + Cl_{2} = HCl + HCl. In this reaction twenty-two thousand heat
units are evolved for one part by weight [1 gram] of hydrogen.[13]

  [11] The chemical action of light on a mixture of chlorine and
       hydrogen was discovered by Gay-Lussac and Thénard (1809). It has
       been investigated by many savants, and especially by Draper,
       Bunsen, and Roscoe. Electric or magnesium light, or the light
       emitted by the combustion of carbon bisulphide in nitric oxide,
       and actinic light in general, acts in the same manner as sunlight,
       in proportion to its intensity. At temperatures below -12° light
       no longer brings about reaction, or at all events does not give an
       explosion. It was long supposed that chlorine that had been
       subjected to the action of light was afterwards able to act on
       hydrogen in the dark, but it was shown that this only takes place
       with moist chlorine, and depends on the formation of oxides of
       chlorine. The presence of foreign gases, and even of excess of
       chlorine or of hydrogen, very much enfeebles the explosion, and
       therefore the experiment is conducted with a detonating mixture
       prepared by the action of an electric current on a strong solution
       (sp. gr. 1·15) of hydrochloric acid, in which case the water is
       not decomposed--that is, no oxygen becomes mixed with the
       chlorine.

  [12] The quantity of chlorine and hydrogen which combine is
       proportional to the intensity of the light--not of all the rays,
       but only those so-termed chemical (actinic) rays which produce
       chemical action. Hence a mixture of chlorine and hydrogen, when
       exposed to the action of light in vessels of known capacity and
       surface, may be employed as an actinometer--that is, as a means
       for estimating the intensity of the chemical rays, the influence
       of the heat rays being previously destroyed, which may be done by
       passing the rays through water. Investigations of this kind
       (photo-chemical) showed that chemical action is chiefly limited to
       the violet end of the spectrum, and that even the invisible
       ultra-violet rays produce this action. A colourless gas flame
       contains no chemically active rays; the flame  green by a
       salt of copper evinces more chemical action than the colourless
       flame, but the flame brightly  yellow by salts of sodium
       has no more chemical action than that of the colourless flame.

       As the chemical action of light becomes evident in plants,
       photography, the bleaching of tissues, and the fading of colours
       in the sunlight, and as a means for studying the phenomenon is
       given in the reaction of chlorine on hydrogen, this subject has
       been the most fully investigated in _photo-chemistry_. The
       researches of Bunsen and Roscoe in the fifties and sixties are the
       most complete in this respect. Their actinometer contains hydrogen
       and chlorine, and is surrounded by a solution of chlorine in
       water. The hydrochloric acid is absorbed as it forms, and
       therefore the variation in volume indicates the progress of the
       combination. As was to be expected, the action of light proved to
       be proportional to the time of exposure and intensity of the
       light, so that it was possible to conduct detailed photometrical
       investigations respecting the time of day and season of the year,
       various sources of light, its absorption, &c. This subject is
       considered in detail in special works, and we only stop to mention
       one circumstance, that a small quantity of a foreign gas decreases
       the action of light; for example, 1/330 of hydrogen by 38 p.c.,
       1/200 of oxygen by 10 p.c., 1/100 of chlorine by 60 p.c., &c.
       According to the researches of Klimenko and Pekatoros (1889), the
       photo-chemical alteration of chlorine water is retarded by the
       presence of traces of metallic chlorides, and this influence
       varies with different metals.

       As much heat is evolved in the reaction of chlorine on hydrogen,
       and as this reaction, being exothermal, may proceed by itself, the
       action of light is essentially the same as that of heat--that is,
       it brings the chlorine and hydrogen into the condition necessary
       for the reaction--it, as we may say, disturbs the original
       equilibrium; this is the work done by the luminous energy. It
       seems to me that the action of light on the mixed gases should be
       understood in this sense, as Pringsheim (1877) pointed out.

  [13] In the formation of steam (from one part by weight [1 gram]
       of hydrogen) 29,000 heat units are evolved. The following are the
       quantities of heat (thousands of units) evolved in the formation
       of various other _corresponding_ compounds of oxygen and of
       chlorine (from Thomsen's, and, for Na_{2}O, Beketoff's results):

       {2NaCl,        195; CaCl_{2},    170; HgCl_{2}, 63; 2AgCl,    59.
       { Na_{2}O,     100;  CaO,        131;  HgO,     42;  Ag_{2}O,  6.
       {2AsCl_{3},    143; 2PCl_{5},    210;  CCl_{4}, 21; 2HCl,     44
                                                                   (gas).
       { As_{2}O_{3}, 155;  P_{2}O_{5}, 370;  CO_{2},  97;  H_{2}O,  58
                                                                   (gas).

       With the first four elements the formation of the chlorine
       compound gives the most heat, and with the four following the
       formation of the oxygen compound evolves the greater amount of
       heat. The first four chlorides are true salts formed from HCl and
       the oxide, whilst the remainder have other properties, as is seen
       from the fact that they are not formed from hydrochloric acid and
       the oxide, but give hydrochloric acid with water.

These relations show that the affinity of chlorine for hydrogen is
very great and analogous to the affinity between hydrogen and oxygen.
Thus[14] on the one hand by passing a mixture of steam and chlorine
through a red-hot tube, or by exposing water and chlorine to the
sunlight, oxygen is disengaged, whilst on the other hand, as we saw
above, oxygen in many cases displaces chlorine from its compound with
hydrogen, and therefore the reaction H_{2}O + Cl_{2} = 2HCl + O belongs
to the number of reversible reactions, and hydrogen will distribute
itself between oxygen and chlorine. This determines the relation of Cl
to substances containing hydrogen and its reactions in the presence of
water, to which we shall turn our attention after having pointed out the
relation of chlorine to other elements.

  [14] This has been already pointed out in Chapter III., Note 5.

Many _metals_ when brought into contact with chlorine immediately
combine with it, and form those metallic chlorides which correspond
with hydrogen chloride and with the oxide of the metal taken. This
combination may proceed rapidly with the evolution of heat and light;
that is, metals are able to burn in chlorine. Thus, for example,
sodium[15] burns in chlorine, synthesising common salt. Metals in
the form of powders burn without the aid of heat, and become highly
incandescent in the process; for instance, antimony, which is a metal
easily converted into a powder.[16] Even such metals as gold and
platinum,[17] which do not combine directly with oxygen and give very
unstable compounds with it, unite directly with chlorine to form
metallic chlorides. Either chlorine water or aqua regia may be employed
for this purpose instead of gaseous chlorine. These dissolve gold and
platinum, converting them into metallic chlorides. _Aqua regia_ is a
mixture of 1 part of nitric acid with 2 to 3 parts of hydrochloric
acid. This mixture converts into soluble chlorides not only those
metals which are acted on by hydrochloric and nitric acids, but also
gold and platinum, which are insoluble in either acid separately.
This action of aqua regia depends on the fact that nitric acid in
acting on hydrochloric acid evolves chlorine. If the chlorine evolved
be transferred to a metal, then a fresh quantity is formed from the
remaining acids and also combines with the metal.[18] Thus the aqua
regia acts by virtue of the chlorine which it contains and disengages.

  [15] Sodium remains unaltered in perfectly dry chlorine at the ordinary
       temperature, and even when slightly warmed; but the combination is
       exceedingly violent at a red heat.

  [16] An instructive experiment on combustion in chlorine may be
       conducted as follows: leaves of Dutch metal (used instead of gold
       for gilding) are placed in a glass globe, and a gas-conducting
       tube furnished with a glass cock is placed in the cork closing it,
       and the air is pumped out of the globe. The gas-conducting tube is
       then connected with a vessel containing chlorine, and the cock
       opened; the chlorine rushes in, and the metallic leaves are
       consumed.

  [17] The behaviour of platinum to chlorine at a high temperature
       (1,400°) is very remarkable, because platinous chloride, PtCl_{2},
       is then formed, whilst this substance decomposes at a much lower
       temperature into chlorine and platinum. Hence, when chlorine comes
       into contact with platinum at such high temperatures, it forms
       fumes of platinous chloride, and they on cooling decompose, with
       the liberation of platinum, so that the phenomenon appears to be
       dependent on the volatility of platinum. Deville proved the
       formation of platinous chloride by inserting a cold tube inside a
       red-hot one (as in the experiment on carbonic oxide). However, V.
       Meyer was able to observe the density of chlorine in a platinum
       vessel at 1,690°, at which temperature chlorine does not exert
       this action on platinum, or at least only to an insignificant
       degree.

  [18] When left exposed to the air aqua regia disengages chlorine, and
       afterwards it no longer acts on gold. Gay-Lussac, in explaining
       the action of aqua regia, showed that when heated it evolves,
       besides chlorine, the vapours of two chloranhydrides--that of
       nitric acid, NO_{2}Cl (nitric acid, NO_{2}OH, in which HO is
       replaced by chlorine; _see_ Chapter on Phosphorus), and that of
       nitrous acid, NOCl--but these do not act on gold. The formation of
       aqua regia may therefore be expressed by 4NHO_{3} + 8HCl =
       2NO_{2}Cl + 2NOCl + 6H_{2}O + 2Cl_{2}. The formation of the
       chlorides NO_{2}Cl and NOCl is explained by the fact that the
       nitric acid is deoxidised, gives the oxides NO and NO_{2}, and
       they directly combine with chlorine to form the above anhydrides.

The majority of _non-metals_ also react directly on chlorine; hot
sulphur and phosphorus burn in it and combine with it at the ordinary
temperature. Only nitrogen, carbon, and oxygen do not combine directly
with it. The chlorine compounds formed by the non-metals--for instance,
phosphorus trichloride, PCl_{3}, and sulphurous chloride, &c., do not
have the properties of salts, and, as we shall afterwards see more
fully, correspond to acid anhydrides and acids; for example, PCl_{3}--to
phosphorous acid, P(OH)_{3}:

               NaCl    FeCl_{2}    SnCl_{4}    PCl_{3}      HCl
               Na(HO)  Fe(HO)_{2}  Sn(HO)_{4}  P(HO)_{3}    H(HO)

As the above-mentioned relation in composition--_i.e._ substitution of
Cl by the aqueous residue--exists between many chlorine compounds and
their corresponding hydrates, and as furthermore some (acid) hydrates
are obtained from chlorine compounds by the action of water, for
instance,

          PCl_{3}    +  3H_{2}0   =   P(HO)_{3}   +    3HCl
         Phosphorus      Water       Phosphorous    Hydrochloric
         trichloride                    acid           acid

whilst other chlorine compounds are formed from hydroxides and
hydrochloric acid, with the liberation of water, for example,

                       NaHO + HCl = NaCl + H_{2}O

we endeavour to express this intimate connection between the hydrates
and chlorine compounds by calling the latter _chloranhydrides_. In
general terms, if the hydrate be basic, then,

        M(HO)   +      HCl           =      MCl       + H_{2}O
       hydrate  + hydrochloric acid  = chloranhydride + water

and if the hydrate ROH be acid, then,

             RCl       + H_{2}O  =  R(HO)   +      HCl
        Chloranhydride + water   =  hydrate + hydrochloric acid

The chloranhydrides MCl corresponding to the bases are evidently
metallic chlorides or salts corresponding to HCl. In this manner a
distinct equivalency is marked between the compounds of chlorine and the
so-called hydroxyl radicle (HO), which is also expressed in the analogy
existing between chlorine, Cl_{2}, and hydrogen peroxide, (HO)_{2}.

As regards the chloranhydrides corresponding to acids and non-metals,
they bear but little resemblance to metallic salts. They are nearly all
volatile, and have a powerful suffocating smell which irritates the eyes
and respiratory organs. They react on water like many anhydrides of the
acids, with the evolution of heat and liberation of hydrochloric acid,
forming acid hydrates. For this reason they cannot usually be obtained
from hydrates--that is, acids--by the action of hydrochloric acid, as in
that case water would be formed together with them, and water decomposes
them, converting them into hydrates. There are many intermediate
chlorine compounds between true saline metallic chlorides like sodium
chloride and true acid chloranhydrides, just as there are all kinds of
transitions between bases and acids. Acid chloranhydrides are not only
obtained from chlorine and non-metals, but also from many lower oxides,
by the aid of chlorine. Thus, for example, CO, NO, NO_{2}, SO_{2}, and
other lower oxides which are capable of combining with oxygen may also
combine with a corresponding quantity of chlorine. Thus COCl_{2},
NOCl, NO_{2}Cl, SO_{2}Cl_{2}, &c., are obtained. They correspond with
the hydrates CO(OH)_{2}, NO(OH), NO_{2}(OH), SO_{2}(OH)_{2}, &c., and
to the anhydrides CO_{2}, N_{2}O_{3}, N_{2}O_{5}, SO_{3}, &c. Here
we should notice two aspects of the matter: (1) chlorine combines
with that with which oxygen is able to combine, because it is in many
respects equally if not more energetic than oxygen and replaces it in
the proportion Cl_{2} : O; (2) that highest limit of possible combination
which is proper to a given element or grouping of elements is very
easily and often attained by combination with chlorine. If phosphorus
gives PCl_{3} and PCl_{5}, it is evident that PCl_{5} is the higher
form of combination compared with PCl_{3}. To the form PCl_{5}, or
in general PX_{5}, correspond PH_{4}I, PO(OH)_{3}, POCl_{3}, &c. If
chlorine does not always directly give compounds of the highest possible
forms for a given element, then generally the lower forms combine
with it in order to reach or approach the limit. This is particularly
clear in hydrocarbons, where we see the limit C_{_n_}H_{2_n_+2} very
distinctly. The unsaturated hydrocarbons are sometimes able to combine
with chlorine with the greatest ease and thus reach the limit. Thus
ethylene, C_{2}H_{4}, combines with Cl_{2}, forming the so-called Dutch
liquid or ethylene chloride, C_{2}H_{4}Cl_{2}, because it then reaches
the limit C_{_n_}X_{2_n_+2}. In this and all similar cases the combined
chlorine is able by reactions of substitution to give a hydroxide and
a whole series of other derivatives. Thus a hydroxide called glycol,
C_{2}H_{4}(OH)_{2}, is obtained from C_{2}H_{4}Cl_{2}.

Chlorine _in the presence of water_ very often acts directly _as an
oxidising agent_. A substance A combines with chlorine and gives, for
example, ACl_{2}, and this in turn a hydroxide, A(OH)_{2}, which on
losing water forms AO. Here the chlorine has oxidised the substance
A. This frequently happens in the simultaneous action of water and
chlorine: A + H_{2}O + Cl_{2} = 2HCl + AO. Examples of this oxidising
action of chlorine may frequently be observed both in practical
chemistry and technical processes. Thus, for instance, chlorine in the
presence of water oxidises sulphur and metallic sulphides. In this
case the sulphur is converted into sulphuric acid, and the chlorine
into hydrochloric acid, or a metallic chloride if a metallic sulphide
be taken. A mixture of carbonic oxide and chlorine passed into water
gives carbonic anhydride and hydrochloric acid. Sulphurous anhydride
is oxidised by chlorine in the presence of water into sulphuric acid,
just as it is by the action of nitric acid: SO_{2} + 2H_{2}O + Cl_{2} =
H_{2}SO_{4} + 2HCl.

The oxidising action of chlorine in the presence of water is taken
advantage of in practice for the rapid bleaching of tissues and
fibres. The colouring matter of the fibres is altered by oxidation and
converted into a colourless substance, but the chlorine afterwards acts
on the tissue itself. Bleaching by means of chlorine therefore requires
a certain amount of technical skill in order that the chlorine should
not act on the fibres themselves, but that its action should be limited
to the colouring matter only. The fibre for making writing paper, for
instance, is bleached in this manner. The bleaching property of chlorine
was discovered by Berthollet, and forms an important acquisition to the
arts, because it has in the majority of cases replaced that which before
was the universal method of bleaching--namely, exposure to the sun of
the fabrics damped with water, which is still employed for linens, &c.
Time and great trouble, and therefore money also, have been considerably
saved by this change.[19]

  [19] Ozone and peroxide of hydrogen also bleach tissues. As the action
       of peroxide of hydrogen is easily controlled by taking a weak
       solution, and as it has hardly any action upon the tissues
       themselves, it is replacing chlorine more and more as a bleaching
       agent. The oxidising property of chlorine is apparent in
       destroying the majority of organic tissues, and proves fatal to
       organisms. This action of chlorine is taken advantage of in
       quarantine stations. But the simple fumigation by chlorine must be
       carried on with great care in dwelling places, because chlorine
       disengaged into the atmosphere renders it harmful to the health.

The power of chlorine for combination is intimately connected with
its capacity for substitution, because, according to the law of
substitution, if chlorine combines with hydrogen, then it also replaces
hydrogen, and furthermore the combination and substitution are
accomplished in the same quantities. Therefore _the atom of chlorine_
which combines with the atom of hydrogen is also able _to replace the
atom of hydrogen_. We mention this property of chlorine not only because
it illustrates the application of the law of substitution in clear and
historically important examples, but more especially because reactions
of this kind explain those _indirect methods_ of the formation of many
substances which we have often mentioned and to which recourse is had
in many cases in chemistry. Thus chlorine does not act on carbon,[20]
oxygen, or nitrogen, but nevertheless its compounds with these elements
may be obtained by the indirect method of the substitution of hydrogen
by chlorine.

  [20] A certain propensity of carbon to attract chlorine is evidenced in
       the immense absorption of chlorine by charcoal (Note 7), but, so
       far as is at present known (if I am not mistaken, no one has tried
       the aid of light), no combination takes place between the chlorine
       and carbon.

As chlorine easily combines with hydrogen, and does not act on carbon,
it decomposes hydrocarbons (and many of their derivatives) at a high
temperature, depriving them of their hydrogen and liberating the carbon,
as, for example, is clearly seen when a lighted candle is placed in a
vessel containing chlorine. The flame becomes smaller, but continues
to burn for a certain time, a large amount of soot is obtained, and
hydrochloric acid is formed. In this case the gaseous and incandescent
substances of the flame are decomposed by the chlorine, the hydrogen
combines with it, and the carbon is disengaged as soot.[21] This
action of chlorine on hydrocarbons, &c., proceeds otherwise at lower
temperatures, as we will now consider.

  [21] The same reaction takes place under the action of oxygen, with
       the difference that it burns the carbon, which chlorine is not
       able to do. If chlorine and oxygen compete together at a high
       temperature, the oxygen will unite with the carbon, and the
       chlorine with the hydrogen.

A very important epoch in the history of chemistry was inaugurated by
the discovery of Dumas and Laurent that chlorine is able to displace
and _replace hydrogen_. This discovery is important from the fact
that chlorine proved to be an element which combines with great ease
simultaneously with both the hydrogen and the element with which the
hydrogen was combined. This clearly proved that there is no opposite
polarity between elements forming stable compounds. Chlorine does not
combine with hydrogen because it has opposite properties, as Dumas and
Laurent stated previously, accounting hydrogen to be electro-positive
and chlorine electro-negative; this is not the reason of their combining
together, for the same chlorine which combines with hydrogen is also
able to replace it without altering many of the properties of the
resultant substance. This substitution of hydrogen by chlorine is termed
_metalepsis_. The mechanism of this substitution is very constant. If we
take a hydrogen compound, preferably a hydrocarbon, and if chlorine acts
directly on it, then there is produced on the one hand hydrochloric acid
and on the other hand a compound containing chlorine in the place of the
hydrogen--so that the chlorine divides itself into two equal portions,
one portion is evolved as hydrochloric acid, and the other portion takes
the place of the hydrogen thus liberated. _Hence this metalepsis is
always accompanied by the formation of hydrochloric acid._[22] The scheme
of the process is as follows:

      C_{n}H_{m}X  +    Cl_{2}  =   C_{n}H_{m-1}ClX  +       HCl
      Hydrocarbon        Free          Product of        Hydrochloric
                       chlorine        metalepsis           acid

Or, in general terms--

                        RH + Cl_{2} = RCl + HCl.

  [22] This division of chlorine into two portions may at the same time
       be taken as a clear confirmation of the conception of molecules.
       According to Avogadro-Gerhardt's law, the molecule of chlorine (p.
       310) contains two atoms of this substance; one atom replaces
       hydrogen, and the other combines with it.

The conditions under which metalepsis takes place are also very
constant. In the dark chlorine does not usually act on hydrogen
compounds, but the action commences under the influence of light.
The direct action of the sun's rays is particularly propitious to
metalepsis. It is also remarkable that the presence of traces of certain
substances,[23] especially of iodine, aluminium chloride, antimony
chloride, &c., promotes the action. A trace of iodine added to the
substance subjected to metalepsis often produces the same effect as
sunlight.[24]

  [23] Such carriers or media for the transference of chlorine and the
       halogens in general were long known to exist in iodine and
       antimonious chloride, and have been most fully studied by
       Gustavson and Friedel, of the Petroffsky Academy--the former with
       respect to aluminium bromide, and the latter with respect to
       aluminium chloride. Gustavson showed that if a trace of metallic
       aluminium be dissolved in bromine (it floats on bromine, and when
       combination takes place much heat and light are evolved), the
       latter becomes endowed with the property of entering into
       metalepsis, which it is not able to do of its own accord. When
       pure, for instance, it acts very slowly on benzene, C_{6}H_{6},
       but in the presence of a trace of aluminium bromide the reaction
       proceeds violently and easily, so that each drop of the
       hydrocarbon gives a mass of hydrobromic acid, and of the product
       of metalepsis. Gustavson showed that the _modus operandi_ of this
       instructive reaction is based on the property of aluminium bromide
       to enter into combination with hydrocarbons and their derivatives.
       The details of this and all researches concerning the metalepsis
       of the hydrocarbons must be looked for in works on organic
       chemistry.

  [24] As small admixtures of iodine, aluminium bromide, &c., aid the
       metalepsis of large quantities of a substance, just as nitric
       oxide aids the reaction of sulphurous anhydride on oxygen and
       water, so the principle is essentially the same in both cases.
       Effects of this kind (which should also be explained by a chemical
       reaction proceeding at the surfaces) only differ from true contact
       phenomena in that the latter are produced by solid bodies and are
       accomplished at their surfaces, whilst in the former all is in
       solution. Probably the action of iodine is founded on the
       formation of iodine chloride, which reacts more easily than
       chlorine.

If marsh gas be mixed with chlorine and the mixture ignited, then the
hydrogen is entirely taken up from the marsh gas and hydrochloric
acid and carbon formed, but there is no metalepsis.[25] But if a
mixture of equal volumes of chlorine and marsh gas be exposed to the
action of diffused light, then the greenish yellow mixture gradually
becomes colourless, and hydrochloric acid and the first product of
metalepsis--namely, methyl chloride--are formed:

         CH_{4}    +    Cl_{2}  =      CH_{3}Cl      +       HCl
        Marsh gas      Chlorine     Methyl chloride     Hydrochloric acid

  [25] Metalepsis belongs to the number of delicate reactions--if it may
       be so expressed--as compared with the energetic reaction of
       combustion. Many cases of substitution are of this kind. Reactions
       of metalepsis are accompanied by an evolution of heat, but in a
       less quantity than that evolved in the formation of the resulting
       quantity of the halogen acids. Thus the reaction C_{2}H_{6} +
       Cl_{2} = C_{2}H_{5}Cl + HCl, according to the data given by
       Thomsen, evolves about 20,000 heat units, whilst the formation of
       hydrochloric acid evolves 22,000 units.

The volume of the mixture remains unaltered. The methyl chloride which
is formed is a gas. If it be separated from the hydrochloric acid (it
is soluble in acetic acid, in which hydrochloric acid is but sparingly
soluble) and be again mixed with chlorine, then it may be subjected
to a further metalepsical substitution--the second atom of hydrogen
may be substituted by chlorine, and a liquid substance, CH_{2}Cl_{2},
called methylene chloride, will be obtained. In the same manner
the substitution may be carried on still further, and CHCl_{3}, or
chloroform, and lastly carbon tetrachloride, CCl_{4}, will be produced.
Of these substances the best known is chloroform, owing to its being
formed from many organic substances (by the action of bleaching powder)
and to its being used in medicine as an anæsthetic; chloroform boils
at 62° and carbon tetrachloride at 78°. They are both colourless
odoriferous liquids, heavier than water. The progressive substitution of
hydrogen by chlorine is thus evident, and it can be clearly seen that
the double decompositions are accomplished between molecular quantities
of the substance--that is, between equal volumes in a gaseous state.

_Carbon tetrachloride_, which is obtained by the metalepsis of marsh
gas, cannot be obtained directly from chlorine and carbon, but it
may be obtained from certain compounds of carbon--for instance, from
carbon bisulphide--if its vapour mixed with chlorine be passed through
a red-hot tube. Both the sulphur and carbon then combine with the
chlorine. It is evident that by ultimate metalepsis a corresponding
carbon chloride may be obtained from any hydrocarbon--indeed, the number
of chlorides of carbon C_{_n_}Cl_{2_m_} already known is very large.

As a rule, the fundamental chemical characters of hydrocarbons are not
changed by metalepsis; that is, if a neutral substance be taken, then
the product of metalepsis is also a neutral substance, or if an acid
be taken the product of metalepsis also has acid properties. Even the
crystalline form not unfrequently remains unaltered after metalepsis.
The metalepsis of acetic acid, CH_{3}·COOH, is historically the most
important. It contains three of the atoms of the hydrogen of marsh gas,
the fourth being replaced by carboxyl, and therefore by the action of
chlorine it gives three products of metalepsis (according to the amount
of the chlorine and conditions under which the reaction takes place),
mono-, di-, and tri-chloracetic acids--CH_{2}Cl·COOH, CHCl_{2}·COOH, and
CCl_{3}·COOH; they are all, like acetic acid, monobasic. The resulting
products of metalepsis, in containing an element which so easily acts
on metals as chlorine, possess the possibility of attaining a further
complexity of molecules of which the original hydrocarbon is often in
no way capable. Thus on treating with an alkali (or first with a salt
and then with an alkali, or with a basic oxide and water, &c.) the
chlorine forms a salt with its metal, and the hydroxyl radicle takes
the place of the chlorine--for example, CH_{3}·OH is obtained from
CH_{3}Cl. By the action of metallic derivatives of hydrocarbons--for
example, CH_{3}Na--the chlorine also gives a salt, and the hydrocarbon
radicle--for instance, CH_{3}--takes the place of the chlorine. In
this, or in a similar manner, CH_{3}·CH_{3}, or C_{2}H_{6} is obtained
from CH_{3}Cl and C_{6}H_{5}·CH_{3} from C_{6}H_{6}. The products of
metalepsis also often react on ammonia, forming hydrochloric acid (and
thence NH_{4}Cl) and an amide; that is, the product of metalepsis,
with the ammonia radicle NH_{2}, &c. in the place of chlorine. Thus
by means of metalepsical substitution methods were found in chemistry
for an artificial and general means of the formation of complex carbon
compounds from more simple compounds which are often totally incapable
of direct reaction. Besides which, this key opened the doors of that
secret edifice of complex organic compounds into which man had up to
then feared to enter, supposing the hydrocarbon elements to be united
only under the influence of those mystic forces acting in organisms.[26]

  [26] With the predominance of the representation of compound radicles
       (this doctrine dates from Lavoisier and Gay-Lussac) in organic
       chemistry, it was a very important moment in its history when it
       became possible to gain an insight into the structure of the
       radicles themselves. It was clear, for instance, that ethyl,
       C_{2}H_{5}, or the radicle of common alcohol, C_{2}H_{5}·OH,
       passes, without changing, into a number of ethyl derivatives, but
       its relation to the still simpler hydrocarbons was not clear, and
       occupied the attention of science in the 'forties' and 'fifties.'
       Having obtained ethyl hydride, C_{2}H_{5}H = C_{2}H_{6}, it was
       looked on as containing the same ethyl, just as methyl hydride,
       CH_{4} = CH_{3}H, was considered as existing in methane. Having
       obtained free methyl, CH_{3}CH_{3} = C_{2}H_{6}, from it, it was
       considered as a derivative of methyl alcohol, CH_{3}OH, and as
       only isomeric with ethyl hydride. By means of the products of
       metalepsis it was proved that this is not a case of isomerism but
       of strict identity, and it therefore became clear that ethyl is
       methylated methyl, C_{2}H_{5} = CH_{2}CH_{3}. In its time a still
       greater impetus was given by the study of the reactions of
       monochloracetic acid, CH_{2}Cl·COOH, or CO(CH_{2}Cl)(OH). It
       appeared that metalepsical chlorine, like the chlorine of
       chloranhydrides--for instance, of methyl chloride, CH_{3}Cl, or
       ethyl chloride, C_{2}H_{5}Cl--is capable of substitution; for
       example, glycollic acid, CH_{2}(OH)(CO_{2}H), or
       CO(CH_{2}·OH)(OH), was obtained from it, and it appeared that the
       OH in the group CH_{2}(OH) reacted like that in alcohols, and it
       became clear, therefore, that it was necessary to examine the
       radicles themselves by analysing them from the point of view of
       the bonds connecting the constituent atoms. Whence arose the
       present doctrine of the structure of the carbon compounds. (_See_
       Chapter VIII., Note 42.)

It is not only hydrocarbons which are subject to metalepsis. Certain
other hydrogen compounds, under the action of chlorine, also give
corresponding chlorine derivatives in exactly the same manner; for
instance, ammonia, caustic potash, caustic lime, and a whole series of
_alkaline_ substances.[27] In fact, just as the hydrogen in marsh gas
can be replaced by chlorine and form methyl chloride, so the hydrogen
in caustic potash, KHO, ammonia, NH_{3}, and calcium hydroxide,
CaH_{2}O_{2} or Ca(OH)_{2}, may be replaced by chlorine and give
potassium hypochlorite, KClO, calcium hypochlorite, CaCl_{2}O_{2},
and the so-called chloride of nitrogen, NCl_{3}. For not only is the
correlation in composition the same as in the substitution in marsh gas,
but the whole mechanism of the reaction is the same. Here also two atoms
of chlorine act: one takes the place of the hydrogen whilst the other is
evolved as hydrochloric acid, only in the former case the hydrochloric
acid evolved remained free, and in the latter, in presence of alkaline
substances, it reacts on them. Thus, in the action of chlorine on
caustic potash, the hydrochloric acid formed acts on another quantity of
caustic potash and gives potassium chloride and water, and therefore not
only KHO + Cl_{2} = HCl + KClO, but also KHO + HCl = H_{2}O + KCl, and
the result of both simultaneous phases will be 2KHO + Cl_{2} = H_{2}O +
KCl + KClO. We will here discuss certain special cases.

  [27] By including many instances of the action of chlorine under
       metalepsis we not only explain the indirect formation of CCl_{4},
       NCl_{3}, and Cl_{2}O by one method, but we also arrive at the fact
       that the reactions of the metalepsis of the hydrocarbons lose that
       exclusiveness which was often ascribed to them. Also by subjecting
       the chemical representations to the law of substitution we may
       foretell metalepsis as a particular case of a general law.

The action of chlorine on ammonia may either result in the entire
breaking up of the ammonia, with the evolution of gaseous nitrogen, or
in a product of metalepsis (as with CH_{4}). With an excess of chlorine
and the aid of heat the ammonia is decomposed, with the disengagement
of free nitrogen.[28] This reaction evidently results in the formation
of sal-ammoniac, 8NH_{3} + 3Cl_{2} = 6NH_{4}Cl + N_{2}. But if the
ammonium salt be in excess, then the reaction takes the direction of the
replacement of the hydrogen in the ammonia by chlorine. The principal
result is that NH_{3} + 3Cl_{2} forms NCl_{3} + 3HCl.[29] The resulting
product of metalepsis, or _chloride of nitrogen_, NCl_{3}, discovered by
Dulong, is a liquid having the property of decomposing with excessive
ease not only when heated, but even under the action of mechanical
influences, as by a blow or by contact with certain solid substances.
The explosion which accompanies the decomposition is due to the fact
that the liquid chloride of nitrogen gives gaseous products, nitrogen
and chlorine.[29 bis]

  [28] This may be taken advantage of in the preparation of nitrogen.
       If a large excess of chlorine water be poured into a beaker, and a
       small quantity of a solution of ammonia be added, then, after
       shaking, nitrogen is evolved. If chlorine act on a dilute solution
       of ammonia, the volume of nitrogen does not correspond with the
       volume of the chlorine taken, because ammonium hypochlorite is
       formed. If ammonia gas be passed through a fine orifice into a
       vessel containing chlorine, the reaction of the formation of
       nitrogen is accompanied by the emission of light and the
       appearance of a cloud of sal-ammoniac. In all these instances an
       excess of chlorine must be present.

  [29] The hydrochloric acid formed combines with ammonia, and therefore
       the final result is 4NH_{3} + 3Cl_{2} = NCl_{3} + 3NH_{4}Cl. For
       this reason, more ammonia must enter into the reaction, but the
       metalepsical reaction in reality only takes place with an excess
       of ammonia or its salt. If bubbles of chlorine be passed through a
       fine tube into a vessel containing ammonia gas, each bubble gives
       rise to an explosion. If, however, chlorine be passed into a
       solution of ammonia, the reaction at first brings about the
       formation of nitrogen, because chloride of nitrogen acts on
       ammonia like chlorine. But when sal-ammoniac has begun to form,
       then the reaction directs itself towards the formation of chloride
       of nitrogen. The first action of chlorine on a solution of
       sal-ammoniac always causes the formation of chloride of nitrogen,
       which then reacts on ammonia thus: NCl_{3} + 4NH_{3} = N_{2} +
       3NH_{4}Cl. Therefore, so long as the liquid is alkaline from the
       presence of ammonia the chief product will be nitrogen. The
       reaction NH_{4}Cl + 3Cl_{2} = NCl_{3} + 4HCl is reversible; with a
       dilute solution it proceeds in the above-described direction
       (perhaps owing to the affinity of the hydrochloric acid for the
       excess of water), but with a strong solution of hydrochloric acid
       it takes the opposite direction (probably by virtue of the
       affinity of hydrochloric acid for ammonia). Therefore there must
       exist a very interesting case of equilibrium between ammonia,
       hydrochloric acid, chlorine, water, and chloride of nitrogen which
       has not yet been investigated. The reaction NCl_{3} + 4HCl =
       NH_{4}Cl + 3Cl_{2} enabled Deville and Hautefeuille to determine
       the composition of chloride of nitrogen. When slowly decomposed by
       water, chloride of nitrogen gives, like a chloranhydride, nitrous
       acid or its anhydride, 2NCl_{3} + 3H_{2}O = N_{2}O_{3} + 6HCl.
       From these observations it is evident that chloride of nitrogen
       presents great chemical interest, which is strengthened by its
       analogy with trichloride of phosphorus. The researches of F. F.
       Selivanoff (1891-94) prove that NCl_{3} may be regarded as an
       ammonium derivative of hypochlorous acid. Chloride of nitrogen is
       decomposed by dilute sulphuric acid in the following manner:
       NCl_{3} + 3H_{2}O + H_{2}SO_{4} = NH_{4}HSO_{4} + 3HClO. This
       reaction is reversible and is only complete when some substance,
       combining with HClO (for instance, succinimide) or decomposing it,
       is added to the liquid. This is easily understood from the fact
       that hypochlorous acid itself, HClO, may, according to the view
       held in this book, be regarded as the product of the metalepsis of
       water, and consequently bears the same relation to NCl_{3} as
       H_{2}O does to NH_{3}, or as RHO to RNH_{2}, R_{2}NH, and
       R_{3}N--that is to say, NCl_{3} corresponds as an ammonium
       derivative to ClOH and Cl_{2} in exactly the same manner as NR_{3}
       corresponds to ROH and R_{2}. The connection of NCl_{3} and other
       similar explosive chloro-nitrogen compounds (called chloryl
       compounds by Selivanoff; for example, the C_{2}H_{5}NCl_{2} of
       Wurtz is chloryl ethylamine), such as NRCl_{2} (as
       NC_{2}H_{5}Cl_{2}), and NR_{2}Cl (for instance, N(CH_{3}CO)HCl,
       chlorylacetamide, and N(C_{2}H_{5})_{2}Cl, chloryl diethylamine)
       with HClO is evident from the fact that under certain
       circumstances these compounds give hypochlorous acid, with water,
       for instance, NR_{2}Cl + H_{2}O = NR_{2}H + HClO, and frequently
       act (like NCl_{3} and HClO, or Cl_{2}) in an oxidising and
       chloridising manner. We may take chloryl succinimide,
       C_{2}H_{4}(CO)_{2}NCl for example. It was obtained by <DW12> by
       the action of HClO upon succinimide, C_{2}H_{4}(CO)_{2}NH, and is
       decomposed by water with the re-formation of amide and HClO (the
       reaction is reversible). Selivanoff obtained, investigated, and
       classified many of the compounds NR_{2}Cl and NRCl_{2}, where R is
       a residue of organic acids or alcohols, and showed their
       distinction from the chloranhydrides, and thus supplemented the
       history of chloride of nitrogen, which is the simplest of the
       amides containing chlorine, NR_{3}, where R is fully substituted
       by chlorine.

  [29 bis] In preparing NCl_{3} every precaution must he used to guard
       against an explosion, and care should he taken that the NCl_{3}
       remains under a layer of water. Whenever an ammoniacal substance
       comes into contact with chlorine great care must be taken, because
       it may be a case of the formation of such products and a very
       dangerous explosion may ensue. The liquid product of the
       metalepsis of ammonia may be most safely prepared in the form of
       small drops by the action of a galvanic current on a slightly warm
       solution of sal-ammoniac; chlorine is then evolved at the positive
       pole, and this chlorine acting on the ammonia gradually forms the
       product of metalepsis which floats on the surface of the liquid
       (being carried up by the gas), and if a layer of turpentine be
       poured on to it these small drops, on coming into contact with the
       turpentine, give feeble explosions, which are in no way dangerous
       owing to the small mass of the substance formed. Drops of chloride
       of nitrogen may with great caution be collected for investigation
       in the following manner. The neck of a funnel is immersed in a
       basin containing mercury, and first a saturated solution of common
       salt is poured into the funnel, and above it a solution of
       sal-ammoniac in 9 parts of water. Chlorine is then slowly passed
       through the solutions, when drops of chloride of nitrogen fall
       into the salt water.

Chloride of nitrogen is a yellow oily liquid of sp. gr. 1·65, which
boils at 71°, and breaks up into N + Cl_{3} at 97°. The contact of
phosphorus, turpentine, india-rubber, &c. causes an explosion, which
is sometimes so violent that a small drop will pierce through a thick
board. The great ease with which chloride of nitrogen decomposes is
dependent upon the fact that it is formed with an absorption of heat,
which it evolves when decomposed, to the amount of about 38,000 heat
units for NCl_{3}, as Deville and Hautefeuille determined.

Chlorine, when absorbed by a solution of caustic soda (and also of other
alkalis) at the ordinary temperature, causes the replacement of the
hydrogen in the caustic soda by the chlorine, with the formation of
sodium chloride by the hydrochloric acid, so that the reaction may be
represented in two phases, as described above. In this manner, sodium
hypochlorite, NaClO, and sodium chloride are simultaneously formed: 2NaHO
+ Cl_{2} = NaCl + NaClO + H_{2}O. The resultant solution contains NaClO
and is termed 'eau de Javelle.' An exactly similar reaction takes place
when chlorine is passed over dry hydrate of lime at the ordinary
temperature: 2Ca(HO)_{2} + 2Cl_{2} = CaCl_{2}O_{2} + CaCl_{2} + 2H_{2}O.
A mixture of the product of metalepsis with calcium chloride is obtained.
This mixture is employed in practice on a large scale, and is termed
'bleaching powder,' owing to its acting, especially when mixed with
acids, as a bleaching agent on tissues, so that it resembles chlorine in
this respect. It is however preferable to chlorine, because the
destructive action of the chlorine can be moderated in this case, and
because it is much more convenient to deal with a solid substance than
with gaseous chlorine. Bleaching powder is also called _chloride of
lime_, because it is obtained from chlorine and hydrate of lime, and
contains[30] both these substances. It may be prepared in the laboratory
by passing a current of chlorine through a cold mixture of water and lime
(milk of lime). The mixture must be kept cold, as otherwise 3Ca(ClO)_{2}
passes into 2CaCl_{2} + Ca(ClO_{3})_{2}. In the manufacture of bleaching
powder in large quantities at chemical works, the purest possible slaked
lime is taken and laid in a thin layer in large flat chambers, M (whose
walls are made of Yorkshire flags or tarred wood, on which chlorine has
no action), and into which chlorine gas is introduced by lead tubes. The
distribution of the plant is shown in the annexed drawing (fig. 67).

[Illustration: FIG. 67.--Apparatus for the manufacture of bleaching
powder (on a small scale) by the action of chlorine, which is generated
in the vessels C, on lime, which is charged into M.]

  [30] Quicklime, CaO (or calcium carbonate, CaCO_{3}), does not absorb
       chlorine when cold, but at a red heat, in a current of chlorine,
       it forms calcium chloride, with the evolution of oxygen. (This was
       confirmed in 1893 by Wells, at Oxford.) This reaction corresponds
       with the decomposing action of chlorine on methane, ammonia, and
       water. Slaked lime (calcium hydroxide, CaH_{2}O_{2}) also, when
       dry, does not absorb chlorine at 100°. The absorption proceeds at
       the ordinary temperature (below 40°). The dry mass thus obtained
       contains not less than three equivalents of calcium hydroxide to
       four equivalents of chlorine, so that its composition is
       [Ca(HO)_{2}]_{5}Cl_{4}. In all probability a simple absorption of
       chlorine by the lime at first takes place in this case, as may be
       seen from the fact that even carbonic anhydride, when acting on
       the dry mass obtained as above, disengages all the chlorine from
       it, leaving only calcium carbonate. But if the bleaching powder be
       obtained by a wet method, or if it be dissolved in water (in which
       it is very soluble), and carbonic anhydride be passed into it,
       then chlorine is no longer disengaged, but chlorine oxide,
       Cl_{2}O, and only half of the chlorine is converted into this
       oxide, while the other half remains in the liquid as calcium
       chloride. From this it may be inferred that calcium chloride is
       formed by the action of water on bleaching powder, and this is
       proved to be the case by the fact that small quantities of water
       extract a considerable amount of calcium chloride from bleaching
       powder. If a large quantity of water act on bleaching powder an
       excess of calcium hydroxide remains, a portion of which is not
       subjected to change. The action of the water may be expressed by
       the following formulæ: From the dry mass Ca_{3}(HO)_{6}Cl_{4}
       there is formed lime, Ca(HO)_{2}, calcium chloride, CaCl_{2}, and
       a saline substance, Ca(ClO)_{2}. Ca_{3}H_{6}O_{6}Cl_{4} =
       CaH_{2}O_{2} + CaCl_{2}O_{2} + CaCl_{2} + 2H_{2}O. The resulting
       substances are not equally soluble; water first extracts the
       calcium chloride, which is the most soluble, then the compound
       Ca(ClO)_{2} and ultimately calcium hydroxide is left. A mixture of
       calcium chloride and hypochlorite passes into solution. On
       evaporation there remains Ca_{2}O_{2}Cl_{4}3H_{2}O. The dry
       bleaching powder does not absorb more chlorine, but the solution
       is able to absorb it in considerable quantity. If the liquid be
       boiled, a considerable amount of chlorine monoxide is evolved.
       After this calcium chloride alone remains in solution, and the
       decomposition may be expressed as follows: CaCl_{2} +
       CaCl_{2}O_{2} + 2Cl_{2} = 2CaCl_{2} + 2Cl_{2}O. Chlorine monoxide
       may be prepared in this manner.

       It is sometimes said that bleaching powder contains a substance,
       Ca(OH)_{2}Cl_{2}, that is calcium peroxide, CaO_{2}, in which one
       atom of oxygen is replaced by (OH)_{2}, and the other by Cl_{2};
       but, judging from what has been said above, this can only be the
       case in the dry state, and not in solutions.

       On being kept for some time, bleaching powder sometimes
       decomposes, with the evolution of oxygen (because CaCl_{2}O_{2} =
       CaCl_{2} + O_{2}, _see_ p. 163); the same takes place when it is
       heated.

The products of the metalepsis of alkaline hydrates, NaClO and
Ca(ClO)_{2}, which are present in solutions of 'Javelle salt' and
bleaching powder (they are not obtained free from metallic chlorides),
must be counted as salts, because their metals are capable of
substitution. But the hydrate HClO corresponding with these salts,
or _hypochlorous acid_, is not obtained in a free or pure state, for
two reasons: in the first place, because this hydrate, as a very
feeble acid, splits up (like H_{2}CO_{3} or HNO_{3}) into water and the
anhydride, or _chlorine monoxide_, Cl_{2}O = 2HClO-H_{2}O; and, in the
second place, because, in a number of instances, it evolves oxygen
with great facility, forming hydrochloric acid: HClO = HCl + O. Both
hypochlorous acid and chlorine monoxide may be regarded as products of
the metalepsis of water, because HOH corresponds with ClOH and ClOCl.
Hence in many instances bleaching salts (a mixture of hypochlorites and
chlorides) break up, with the evolution of (1) _chlorine_, under the
action of an excess of a powerful acid capable of evolving hydrochloric
acid from sodium or calcium chlorides, and this takes place most simply
under the action of hydrochloric acid itself, because (p. 462) NaCl +
NaClO + 3HCl = 2NaCl + HCl + Cl_{2} + H_{2}O; (2) _oxygen_, as we saw
in Chapter III.--The bleaching properties and, in general, _oxidising
action_ of bleaching salts is based on this evolution of oxygen (or
chlorine); oxygen is also disengaged on heating the dry salts--for
instance, NaCl + NaClO = 2NaCl + O; (3) and, lastly, _chlorine
monoxide_, which contains both chlorine and oxygen. Thus, if a little
sulphuric, nitric, or similar acid (not enough to liberate hydrochloric
acid from the CaCl_{2}) be added to a solution of a bleaching salt
(which has an alkaline reaction, owing either to an excess of alkali
or to the feeble acid properties of HClO), then the hypochlorous acid
set free gives water and chlorine monoxide. If carbonic anhydride
(or boracic or a similar very feeble acid) act on the solution of a
bleaching salt, then hydrochloric acid is not evolved from the sodium
or calcium chlorides, but the hypochlorous acid is displaced and gives
chlorine monoxide,[31] because hypochlorous acid is one of the most
feeble acids. Another method for the preparation of chlorine monoxide
is based on these feeble acid properties of hypochlorous acid. Zinc
oxide and mercury oxide, under the action of chlorine in the presence
of water, do not give a salt of hypochlorous acid, but form a chloride
and hypochlorous acid, which fact shows the incapacity of this acid to
combine with the bases mentioned. Therefore, if such oxides as those of
zinc or mercury be shaken up in water, and chlorine be passed through
the turbid liquid,[32] a reaction occurs which may be expressed in the
following manner: 2HgO + 2Cl_{2} = Hg_{2}OCl_{2} + Cl_{2}O. In this case,
a compound of mercury oxide with mercury chloride, or the so-called
mercury oxychloride, is obtained: Hg_{2}OCl_{2} = HgO + HgCl_{2}. This is
insoluble in water, and is not affected by hypochlorous anhydride, so
that the solution will contain hypochlorous acid only, but the greater
part of it splits up into the anhydride and water.[32 bis]

  [31] For this reason it is necessary that in the preparation of
       bleaching powder the chlorine should be free from hydrochloric
       acid, and even the lime from calcium chloride. An excess of
       chlorine, in acting on a solution of bleaching powder, may also
       give chlorine monoxide, because calcium carbonate also gives
       chlorine monoxide under the action of chlorine. This reaction may
       be brought about by treating freshly precipitated calcium
       carbonate with a stream of chlorine in water: 2Cl_{2} + CaCO_{3} =
       CO_{2} + CaCl_{2} + Cl_{2}O. From this we may conclude that,
       although carbonic anhydride displaces hypochlorous anhydride, it
       may be itself displaced by an excess of the latter.

  [32] Dry red mercury oxide acts on chlorine, forming dry hypochlorous
       anhydride (chlorine monoxide) (Balard); when mixed with water, red
       mercury oxide acts feebly on chlorine, and when freshly
       precipitated it evolves oxygen and chlorine. An oxide of mercury
       which easily and abundantly evolves chlorine monoxide under the
       action of chlorine in the presence of water may be prepared as
       follows: the oxide of mercury, precipitated from a mercuric salt
       by an alkali, is heated to 300° and cooled (Pelouze). If a salt,
       MClO, be added to a solution of mercuric salt, HgX_{2}, mercuric
       oxide is liberated, because the hypochlorite is decomposed.

  [32 bis] A solution of hypochlorous anhydride is also obtained by the
       action of chlorine on many salts; for example, in the action of
       chlorine on a solution of sodium sulphate the following reaction
       takes place: Na_{2}SO_{4} + H_{2}O + Cl_{2} = NaCl + HClO +
       NaHSO_{4}. Here the hypochlorous acid is formed, together with
       HCl, at the expense of chlorine and water, for Cl_{2} + H_{2}O =
       HCl + HClO. If the crystallo-hydrate of chlorine be mixed with
       mercury oxide, the hydrochloric acid formed in the reaction gives
       mercury chloride, and hypochlorous acid remains in solution. A
       dilute solution of hypochlorous acid or chlorine monoxide may be
       concentrated by distillation, and if a substance which takes up
       water (without destroying the acid)--for instance, calcium
       nitrate--be added to the stronger solution, then the anhydride of
       hypochlorous acid--_i.e._ chlorine monoxide--is disengaged.

Chlorine monoxide, which corresponds to bleaching and hypochlorous
salts, containing as it does the two elements oxygen and chlorine,
forms a characteristic example of a compound of elements which, in the
majority of cases, act chemically in an analogous manner. Chlorine
monoxide, as prepared from an aqueous solution by the abstraction of
water or by the action of dry chlorine on cold mercury oxide, is, at
the ordinary temperature, a gas or vapour which condenses into a red
liquid boiling at +20° and giving a vapour whose density (43 referred
to hydrogen) shows that 2 vols. of chlorine and 1 vol. of oxygen
give 2 vols. of chlorine monoxide. In an anhydrous form the gas or
liquid easily explodes, splitting up into chlorine and oxygen. This
explosiveness is determined by the fact that heat is _evolved_ in the
decomposition to the amount of about 15,000 heat units for Cl_{2}O.[33]
The explosion may even take place spontaneously, and also in the
presence of many oxidisable substances (for instance, sulphur, organic
compounds, &c.), but the solution, although unstable and showing a
strong oxidising tendency, does not explode.[34] It is evident that
the presence of hypochlorous acid, HClO, may be assumed in an aqueous
solution of Cl_{2}O, since Cl_{2}O + H_{2}O = 2HClO.

  [33] All explosive substances are of this kind--ozone, hydrogen
       peroxide, chloride of nitrogen, nitro-compounds, &c. Hence they
       cannot be formed directly from the elements or their simplest
       compounds, but, on the contrary, decompose into them. In a liquid
       state chlorine monoxide explodes even on contact with powdery
       substances, or when rapidly agitated--for instance, if a file be
       rasped over the vessel in which it is contained.

  [34] A solution of chlorine monoxide, or hypochlorous acid, does not
       explode, owing to the presence of the mass of water. In
       dissolving, chlorine monoxide evolves about 9,000 heat units, so
       that its store of heat becomes less.

       The capacity of hypochlorous acid (studied by Carius and others)
       for entering into combination with the unsaturated hydrocarbons is
       very often taken advantage of in organic chemistry. Thus its
       solution absorbs ethylene, forming the chlorhydrin C_{2}H_{4}ClOH.

       The oxidising action of hypochlorous acid and its salts is not
       only applied to bleaching but also to many reactions of oxidation.
       Thus it converts the lower oxides of manganese into the peroxide.

Hypochlorous acid, its salts, and chlorine monoxide serve as a transition
between hydrochloric acid, chlorides, and chlorine, and a whole series of
compounds containing the same elements combined with a still greater
quantity of oxygen. The higher oxides of chlorine, as their origin
indicates, are closely connected with hypochlorous acid and its salts:

         Cl_{2},       NaCl,       HCl,       hydrochloric acid.
         Cl_{2}O,      NaClO,      HClO,      hypochlorous acid.
         Cl_{2}O_{3},  NaClO_{2},  HClO_{2},  chlorous acid.[35]
         Cl_{2}O_{5},  NaClO_{3},  HClO_{3},  chloric acid.
         Cl_{2}O_{7},  NaClO_{4},  HClO_{4},  perchloric acid.

When heated, solutions of hypochlorites undergo a remarkable change.
Themselves so unstable, they, without any further addition, yield two
fresh salts which are both much more stable; one contains more oxygen
than MClO, the other contains none at all.

                     3MClO       =   MClO_{3} +    2MCl
                  hypochlorite       chlorate    chloride

  [35] _Chlorous acid_, HClO_{2} (according to the data given by Millon,
       Brandau, and others) in many respects resembles hypochlorous acid,
       HClO, whilst they both differ from chloric and perchloric acids in
       their degree of stability, which is expressed, for instance, in
       their bleaching properties; the two higher acids do not bleach,
       but both the lower ones do so (oxidise at the ordinary
       temperature). On the other hand, chlorous acid is analogous to
       nitrous acid, HNO_{2}. The anhydride of chlorous acid,
       Cl_{2}O_{3}, is not known in a pure state, but it probably occurs
       in admixture with chlorine dioxide, ClO_{2}, which is obtained by
       the action of nitric and sulphuric acids on a mixture of potassium
       chlorate with such reducing substances as nitric oxide, arsenious
       oxide, sugar, &c. All that is at present known is that pure
       chlorine dioxide ClO_{2} (_see_ Notes 39-43) is gradually
       converted into a mixture of hypochlorous and chlorous acids under
       the action of water (and alkalis); that is, it acts like nitric
       peroxide, NO_{2} (giving HNO_{3} and HNO_{2}), or as a mixed
       anhydride, 2ClO_{2} + H_{2}O = HClO_{3} + HClO_{2}. The silver
       salt, AgClO_{2}, is sparingly soluble in water. The investigations
       of Garzarolli-Thurnlackh and others seem to show that the
       anhydride Cl_{2}O_{3} does not exist in a free state.

Part of the salt--namely, two-thirds of it--parts with its oxygen in
order to oxidise the remaining third.[36] From an intermediate substance,
RX, two extremes, R and RX_{3} are formed, just as nitrous anhydride
splits up into nitric oxide and nitric anhydride (or nitric acid). The
resulting salt, MClO_{3}, corresponds with _chloric acid_ and potassium
chlorate, KClO_{3}. It is evident that a similar salt may be obtained
directly by the action of chlorine on an alkali if its solution be
heated, because RClO will be first formed, and then RClO_{3}; for
example, 6KHO + 3Cl_{2} = KClO_{3} + 5KCl + 3H_{2}O. Chlorates are so
prepared; for instance, _potassium chlorate_, which is easily separated
from potassium chloride, being sparingly soluble in cold water.[37]

  [36] Hydrochloric acid, which is an example of compounds of this kind,
       is a saturated substance which does not combine directly with
       oxygen, but in which, nevertheless, a considerable quantity of
       oxygen may be inserted between the elements forming it. The same
       may be observed in a number of other cases. Thus oxygen may be
       added or inserted between the elements, sometimes in considerable
       quantities, in the saturated hydrocarbons; for instance, in
       C_{3}H_{8}, three atoms of oxygen produce an alcohol, glycerin or
       glycerol, C_{3}H_{5}(OH)_{3}. We shall meet with similar examples
       hereafter. This is generally explained by regarding oxygen as a
       bivalent element--that is, as capable of combining with two
       different elements, such as chlorine, hydrogen, &c. On the basis
       of this view, it may be inserted between each pair of combined
       elements; the oxygen will then be combined with one of the
       elements by one of its affinities and with the other element by
       its other affinity. This view does not, however, express the
       entire truth of the matter, even when applied to the compounds of
       chlorine. Hypochlorous acid, HOCl--that is, hydrochloric acid in
       which one atom of oxygen is inserted--is, as we have already seen,
       a substance of small stability; it might therefore be expected
       that on the addition of a fresh quantity of oxygen, a still less
       stable substance would be obtained, because, according to the
       above view, the chlorine and hydrogen, which form such a stable
       compound together, are then still further removed from each other.
       But it appears that chloric and perchloric acid, HClO_{3} and
       HClO_{4}, are much more stable substances. Furthermore, the
       addition of oxygen has also its limit, it can only be added to a
       certain extent. If the above representation were true and not
       merely hypothetical, there would be no limit to the combination of
       oxygen, and the more it entered into one continuous chain the more
       unstable would be the resultant compound. But not more than four
       atoms of oxygen can be added to hydrogen sulphide, nor to
       hydrochloric acid, nor to hydrogen phosphide. This peculiarity
       must lie in the properties of oxygen itself; four atoms of oxygen
       seem to have the power of forming a kind of radicle which retains
       two or several atoms of various other substances--for example,
       chlorine and hydrogen, hydrogen and sulphur, sodium and manganese,
       phosphorus and metals, &c., forming comparatively stable
       compounds, NaClO_{4}, Na_{2}SO_{4}, NaMnO_{4}, Na_{3}PO_{4}, &c.
       _See_ Chapter X. Note 1 and Chapter XV.

  [37] If chlorine be passed through a _cold_ solution of potash, a
       bleaching compound, potassium chloride and hypochlorite, KCl +
       KClO, is formed, but if it be passed through a _hot_ solution
       potassium chlorate is formed. As this is sparingly soluble in
       water, it chokes the gas-conducting tube, which should therefore
       be widened out at the end.

       Potassium chlorate is usually obtained on a large scale from
       calcium chlorate, which is prepared by passing chlorine (as long
       as it is absorbed) into water containing lime, the mixture being
       kept warm. A mixture of calcium chlorate and chloride is thus
       formed in the solution. Potassium chloride is then added to the
       warm solution, and on cooling a precipitate of potassium chlorate
       is formed as a substance which is sparingly soluble in cold water,
       especially in the presence of other salts. The double
       decomposition taking place is Ca(ClO_{3})_{2} + 2KCl = CaCl_{2} +
       2KClO_{3}. On a small scale in the laboratory potassium chlorate
       is best prepared from a strong solution of bleaching powder by
       passing chlorine through it and then adding potassium chloride.
       KClO_{3} is always formed by the action of an electric current on
       a solution of KCl, especially at 80° (Häussermann and Naschold,
       1894), so that this method is now used on a large scale.

       Potassium chlorate crystallises easily in large colourless tabular
       crystals. Its solubility in 100 parts of water at 0° = 3 parts,
       20° = 8 parts, 40° = 14 parts, 60° = 25 parts, 80° = 40 parts. For
       comparison we will cite the following figures showing the
       solubility of potassium chloride and perchlorate in 100 parts of
       water: potassium chloride at O° = 28 parts, 20° = 35 parts, 40° =
       40 parts, 100° = 57 parts; potassium perchlorate at 0° about 1
       part, 20° about 1-3/4 part, 100° about 18 parts. When heated,
       potassium chlorate melts (the melting point has been given as from
       335°-376°; according to the latest determination by Carnelley,
       359°) and decomposes with the evolution of oxygen, potassium
       perchlorate being at first formed, as will afterwards be described
       (_see_ Note 47). A mixture of potassium chlorate and nitric and
       hydrochloric acids effects oxidation and chlorination in
       solutions. It deflagrates when thrown upon incandescent carbon,
       and when mixed with sulphur (1/3 by weight) it ignites it on being
       struck, in which case an explosion takes place. The same occurs
       with many metallic sulphides and organic substances. Such mixtures
       are also ignited by a drop of sulphuric acid. All these effects
       are due to the large amount of oxygen contained in potassium
       chlorate, and to the ease with which it is evolved. A mixture of
       two parts of potassium chlorate, one part of sugar, and one part
       of yellow prussiate of potash acts like gunpowder, but burns too
       rapidly, and therefore bursts the guns, and it also has a very
       strong oxidising action on their metal. The sodium salt,
       NaClO_{3}, is much more soluble than the potassium salt, and it is
       therefore more difficult to free it from sodium chloride, &c. The
       barium salt is also more soluble than the potassium salt; O° = 24
       parts, 20° = 37 parts, 80° = 98 parts of salt per 100 of water.

If dilute sulphuric acid be added to a solution of potassium chlorate,
_chloric acid_ is liberated, but it cannot be separated by distillation,
as it is decomposed in the process. To obtain the free acid, sulphuric
acid must be added to a solution of barium chlorate.[38] The sulphuric
acid gives a precipitate of barium sulphate, and free chloric acid
remains in solution. The solution may be evaporated under the receiver of
an air-pump. This solution is colourless, has no smell, and acts as a
powerful acid (it neutralises sodium hydroxide, decomposes sodium
carbonate, gives hydrogen with zinc, &c.); when heated above 40°,
however, it decomposes, forming chlorine, oxygen, and perchloric acid:
4HClO_{3} = 2HClO_{4} + H_{2}O + Cl_{2} + O_{3}. In a concentrated
condition the acid acts as an exceedingly energetic oxidiser, so that
organic substances brought into contact with it burst into flame. Iodine,
sulphurous acid, and similar oxidisable substances form higher oxidation
products and reduce the chloric acid to hydrochloric acid. Hydrochloric
acid gas gives chlorine with chloric acid (and consequently with
KClO_{3} also) acting in the same manner as it acts on the lower acids:
HClO_{3} + 5HCl = 3H_{2}O + 3Cl_{2}.

  [38] Barium chlorate, Ba(ClO_{3})_{2},H_{2}O, is prepared in the
       following way: impure chloric acid is first prepared and saturated
       with baryta, and the barium salt purified by crystallisation. The
       impure free chloric acid is obtained by converting the potassium
       in potassium chlorate into an insoluble salt. This is done by
       adding tartaric or hydrofluosilicic acid to a solution of
       potassium chlorate, because potassium tartrate and potassium
       silicofluoride are very sparingly soluble in water. Chloric acid
       is easily soluble in water.

By cautiously acting on potassium chlorate with sulphuric acid, the
_dioxide_ (_chloric peroxide_), ClO_{2},[39] is obtained (Davy, Millon).
This gas is easily liquefied in a freezing mixture, and boils at +10°.
The vapour density (about 35 if H = 1) shows that the molecule of this
substance is ClO_{2}.[40] In a gaseous or liquid state it very easily
explodes (for instance, at 60°, or by contact with organic compounds or
finely divided substances, &c.), forming Cl and O_{2}, and in many
instances[41] therefore it acts as an oxidising agent, although (like
nitric peroxide) it may itself be further oxidised.[42] In dissolving in
water or alkalis chloric peroxide gives chlorous and hypochlorous
acids--2ClO_{2} + 2KHO = KClO_{3} + KClO_{2} + H_{2}O--and therefore,
like nitric peroxide, the dioxide may be regarded as an intermediate
oxide between the (unknown) anhydrides of chlorous and chloric acids:
4ClO_{2} = Cl_{2}O_{3} + Cl_{2}O_{3}.[43]

  [39] To prepare ClO_{2} 100 grams of sulphuric acid are cooled in a
       mixture of ice and salt, and 15 grams of powdered potassium
       chlorate are gradually added to the acid, which is then carefully
       distilled at 20° to 40°, the vapour given off being condensed in a
       freezing mixture. Potassium perchlorate is then formed: 3KClO_{3}
       + 2H_{2}SO_{4} = 2KHSO_{4} + KClO_{4} + 2ClO_{2} + H_{2}O. The
       reaction may result in an explosion. Calvert and Davies obtained
       chloric peroxide without the least danger by heating a mixture of
       oxalic acid and potassium chlorate in a test tube in a water-bath.
       In this case 2KClO_{3} + 3C_{2}H_{2}O_{4},2H_{2}O = 2C_{2}HKO_{4}
       + 2CO_{2} + 2ClO_{2} + 8H_{2}O. The reaction is still further
       facilitated by the addition of a small quantity of sulphuric acid.
       If a solution of HCl acts upon KClO_{3} at the ordinary
       temperature, a mixture of Cl_{2} and ClO_{2} is formed, but if the
       temperature be raised to 80° the greater part of the ClO_{2}
       decomposes, and when passed through a hot solution of MnCl_{2} it
       oxidises it. Gooch and Kreider proposed (1894) to employ this
       method for preparing small quantities of chlorine in the
       laboratory.

  [40] By analogy with nitric peroxide it might be expected that at low
       temperatures a doubling of the molecule into Cl_{2}O_{4} would
       take place, as the reactions of ClO_{2} point to its being a mixed
       anhydride of HClO_{2} and HClO_{3}.

  [41] Owing to the formation of this chlorine dioxide, a mixture of
       potassium chlorate and sugar is ignited by a drop of sulphuric
       acid. This property was formerly made use of for making matches,
       and is now sometimes employed for setting fire to explosive
       charges by means of an arrangement in which the acid is caused to
       fall on the mixture at the moment required. An interesting
       experiment on the combustion of phosphorus under water may be
       conducted with chlorine dioxide. Pieces of phosphorus and of
       potassium chlorate are placed under water, and sulphuric acid is
       poured on to them (through a long funnel); the phosphorus then
       burns at the expense of the chlorine dioxide.

  [42] Potassium permanganate oxidises chlorine dioxide into chloric acid
       (Fürst).

  [43] The euchlorine obtained by Davy by gently heating potassium
       chlorate with hydrochloric acid is (Pebal) a mixture of chlorine
       dioxide and free chlorine. The liquid and gaseous chlorine oxide
       (Note 35), which Millon considered to be Cl_{2}O_{3}, probably
       contains a mixture of ClO_{2} (vapour density 35), Cl_{2}O_{3}
       (whose vapour density should be 59), and chlorine (vapour density
       35·5), since its vapour density was determined to be about 40.

As the salts of chloric acid, HClO_{3}, are produced by the splitting
up of the salts of hypochlorous acid, so in the same way the salts
of perchloric acid, HClO_{4}, are produced from the salts of chloric
acid, HClO_{3}. But this is the highest form of the oxidation of HCl.
_Perchloric acid_, HClO_{4}, is the most stable of all the acids of
chlorine. When fused potassium chlorate begins to swell up and solidify,
after having parted with one-third of its oxygen, potassium chloride and
potassium perchlorate have been formed according to the equation 2KClO_{3}
= KClO_{4} + KCl + O_{2}.

The formation of this salt is easily observed in the preparation of
oxygen from potassium chlorate, owing to the fact that the potassium
perchlorate fuses with greater difficulty than the chlorate, and
therefore appears in the molten salt as solid grains (_see_ Chapter III.
Note 12). Under the action of certain acids--for instance, sulphuric and
nitric--potassium chlorate also gives potassium perchlorate. This latter
may be easily purified, because it is but sparingly soluble in water,
although all the other salts of perchloric acid are very soluble and even
deliquesce in the air. The perchlorates, although they contain more
oxygen than the chlorates, are decomposed with greater difficulty, and
even when thrown on ignited charcoal give a much feebler deflagration
than the chlorates. Sulphuric acid (at a temperature not below 100°)
evolves volatile and to a certain extent stable perchloric acid from
potassium perchlorate. Neither sulphuric nor any other acid will further
decompose perchloric acid as it decomposes chloric acid. Of all the acids
of chlorine, perchloric acid alone can be distilled.[44] The pure hydrate
HClO_{4}[45] is a colourless and exceedingly caustic substance which
fumes in the air and has a specific gravity 1·78 at 15° (sometimes, after
being kept for some time, it decomposes with a violent explosion). It
explodes violently when brought into contact with charcoal, paper, wood,
and other organic substances. If a small quantity of water be added to
this hydrate, and it be cooled, a crystallo-hydrate, ClHO_{4},H_{2}O,
separates out. This is much more stable, but the liquid hydrate
HClO_{4},2H_{2}O is still more so. The acid dissolves in water in all
proportions, and its solutions are distinguished for their stability.[46]
When ignited both the acid and its salts are decomposed, with the
evolution of oxygen.[47]

  [44] If a solution of chloric acid, HClO_{3}, be first concentrated
       over sulphuric acid under the receiver of an air-pump and
       afterwards distilled, chlorine and oxygen are evolved and
       perchloric acid is formed: 4HClO_{3} = 2HClO_{4} + Cl_{2} + 3O +
       H_{2}O. Roscoe accordingly decomposed directly a solution of
       potassium chlorate by hydrofluosilicic acid, decanted it from the
       precipitate of potassium silicofluoride, K_{2}SiF_{6},
       concentrated the solution of chloric acid, and then distilled it,
       perchloric acid being then obtained (_see_ following footnote).
       That chloric acid is capable of passing into perchloric acid is
       also seen from the fact that potassium permanganate is
       decolorised, although slowly, by the action of a solution of
       chloric acid. On decomposing a solution of potassium chlorate by
       the action of an electric current, potassium perchlorate is
       obtained at the positive electrode (where the oxygen is evolved).
       Perchloric acid is also formed by the action of an electric
       current on solutions of chlorine and chlorine monoxide. Perchloric
       acid was obtained by Count Stadion and afterwards by Serullas, and
       was studied by Roscoe and others.

  [45] Perchloric acid, which is obtained in a free state by the action
       of sulphuric acid on its salts, may be separated from a solution
       very easily by distillation, being volatile, although it is
       partially decomposed by distillation. The solution obtained after
       distillation may be concentrated by evaporation in open vessels.
       In the distillation the solution reaches a temperature of 200°,
       and then a very constant liquid hydrate of the composition
       HClO_{4},2H_{2}O is obtained in the distillate. If this hydrate be
       mixed with sulphuric acid, it begins to decompose at 100°, but
       nevertheless a portion of the acid passes over into the receiver
       without decomposing, forming a crystalline hydrate HClO_{4},H_{2}O
       which melts at 50°. On carefully heating this hydrate it breaks up
       into perchloric acid, which distills over below 100°, and into the
       liquid hydrate HClO_{4},2H_{2}O. The acid HClO_{4} may also be
       obtained by adding one-fourth part of strong sulphuric acid to
       potassium chlorate, carefully distilling and subjecting the
       crystals of the hydrate HClO_{4},H_{2}O obtained in the distillate
       to a fresh distillation. Perchloric acid, HClO_{4}, itself does
       not distil, and is decomposed on distillation until the more
       stable hydrate HClO_{4}, H_{2}O is formed; this decomposes into
       HClO_{4} and HClO_{4},2H_{2}O, which latter hydrate distils without
       decomposition. This forms an excellent example of the influence of
       water on stability, and of the property of chlorine of giving
       compounds of the type ClX_{7}, of which all the above hydrates,
       ClO_{3}(OH), ClO_{2}(OH)_{3}, and ClO(OH)_{5}, are members.
       Probably further research will lead to the discovery of a hydrate
       Cl(OH)_{7}.

  [46] According to Roscoe the specific gravity of perchloric acid = 1·782
       and of the hydrate HClO_{4},H_{2}O in a liquid state (50°) 1·811;
       hence a considerable contraction takes place in the combination of
       HClO_{4} with H_{2}O.

  [47] The decomposition of salts analogous to potassium chlorate has been
       more fully studied in recent years by Potilitzin and P. Frankland.
       Professor Potilitzin, by decomposing, for example, lithium
       chlorate LiClO_{3}, found (from the quantity of lithium chloride
       and oxygen) that at first the decomposition of the fused salt
       (368°) takes place according to the equation, 3LiClO_{3} = 2LiCl +
       LiClO_{4} + 5O, and that towards the end the remaining salt is
       decomposed thus: 5LiClO_{3} = 4LiCl + LiClO_{4} + 10O. The
       phenomena observed by Potilitzin obliged him to admit that lithium
       perchlorate is capable of decomposing simultaneously with lithium
       chlorate, with the formation of the latter salt and oxygen; and
       this was confirmed by direct experiment, which showed that lithium
       chlorate is always formed in the decomposition of the perchlorate.
       Potilitzin drew particular attention to the fact that the
       decomposition of potassium chlorate and of salts analogous to it,
       although exothermal (Chapter III., Note 12), not only does not
       proceed spontaneously, but requires time and a rise of temperature
       in order to attain completion, which again shows that chemical
       equilibria are not determined by the heat effects of reactions
       only.

       P. Frankland and J. Dingwall (1887) showed that at 448° (in the
       vapour of sulphur) a mixture of potassium chlorate and powdered
       glass is decomposed almost in accordance with the equation
       2KClO_{3} = KClO_{4} + KCl + O_{2}, whilst the salt by itself
       evolves about half as much oxygen, in accordance with the
       equation, 8KClO_{3} = 5KClO_{4} + 3KCl + 2O_{2}. The decomposition
       of potassium perchlorate in admixture with manganese peroxide
       proceeds to completion, KClO_{4} = KCl + 2O_{2}. But in
       decomposing by itself the salt at first gives potassium chlorate,
       approximately according to the equation 7KClO_{4} = 2KClO_{3} +
       5KCl + 11O_{2}. Thus there is now no doubt that when potassium
       chlorate is heated, the perchlorate is formed, and that this salt,
       in decomposing with evolution of oxygen, again gives the former
       salt.

       In the decomposition of barium hypochlorite, 50 per cent. of the
       whole amount passes into chlorate, in the decomposition of
       strontium hypochlorite (Potilitzin, 1890) 12·5 per cent., and of
       calcium hypochlorite about 2·5 per cent. Besides which Potilitzin
       showed that the decomposition of the hypochlorites and also of the
       chlorates is always accompanied by the formation of a certain
       quantity of the oxides and by the evolution of chlorine, the
       chlorine being displaced by the oxygen disengaged. Spring and
       Prost (1889) represent the evolution of oxygen from KClO_{3} as
       due to the salt first splitting up into base and anhydride, thus
       (1) 2MClO_{3} = M_{2}O + Cl_{2}O_{5}; (2) Cl_{2}O_{5} = Cl_{2} +
       O_{3}; and (3) M_{2}O + Cl = 2MCl + O.

       I may further remark that the decomposition of potassium chlorate
       as a reaction evolving heat easily lends itself for this very
       reason to the contact action of manganese peroxide and other
       similar admixtures; for such very feeble influences as those of
       contact may become evident either in those cases (for instance,
       detonating gas, hydrogen peroxide, &c.), when the reaction is
       accompanied by the evolution of heat, or when (for instance, H_{2}
       + I_{2}, &c.) little heat is absorbed or evolved. In these cases
       it is evident that the existing equilibrium is not very stable,
       and that a small alteration in the conditions at the surfaces of
       contact may suffice to upset it. In order to conceive the _modus
       operandi_ of contact phenomena, it is enough to imagine, for
       instance, that at the surface of contact the movement of the atoms
       in the molecules changes from a circular to an elliptical path.
       Momentary and transitory compounds may he formed, but their
       formation cannot affect the explanation of the phenomena.

On comparing chlorine as an element not only with nitrogen and carbon but
with all the other non-metallic elements (chlorine has so little analogy
with the metals that a comparison with them would be superfluous), we
find in it the following fundamental properties of _the halogens_ or
salt-producers. With metals chlorine gives salts (such as sodium
chloride, &c.); with hydrogen a very energetic and monobasic acid HCl,
and the same quantity of chlorine is able by metalepsis to replace the
hydrogen; with oxygen it forms unstable oxides of an acid character.
These properties of chlorine are possessed by three other elements,
bromine, iodine, and fluorine. They are members of one natural family.
Each representative has its peculiarities, its individual properties and
points of distinction, in combination and in the free state--otherwise
they would not be independent elements; but the repetition in all of them
of the same chief characteristics of the family enables one more quickly
to grasp all their various properties and to classify the elements
themselves.

In order to have a guiding thread in forming comparisons between the
elements, attention must however be turned not only to their points of
resemblance but also to those of their properties and characters in
which they differ most from each other. And the atomic weights of the
elements must be considered as their most elementary property, since
this is a quantity which is most firmly established, and must be taken
account of in all the reactions of the element. The halogens have the
following atomic weights--

                 F = 19,  Cl = 35·5,  Br = 80,  I = 127.

All the properties, physical and chemical, of the elements and their
corresponding compounds must evidently be in a certain dependence on
this fundamental point, if the grouping in one family be natural.[47 bis]
And we find in reality that, for instance, the properties of bromine,
whose atomic weight is almost the mean between those of iodine and
chlorine, occupy a mean position between those of these two elements. The
second measurable property of the elements is their equivalence or their
capacity for forming _compounds of definite forms_. Thus carbon or
nitrogen in this respect differs widely from the halogens. Although the
form ClO_{2} corresponds with NO_{2} and CO_{2}, yet the last is the
highest oxide of carbon, whilst that of nitrogen is N_{2}O_{5}, and for
chlorine, if there were an anhydride of perchloric acid, its composition
would be Cl_{2}O_{7}, which is quite different from that of carbon. In
respect to the forms of their compounds the halogens, like all elements
of one family or group, are perfectly analogous to each other, as is seen
from their hydrogen compounds:

                            HF, HCl, HBr, HI.

  [47 bis] See, for example the melting point of NaCl, NaBr, NaI in
       Chapter II. Note 27. According to F. Freyer and V. Meyer (1892),
       the following are the boiling points of some of the corresponding
       compounds of chlorine and bromine:

                      BCl_{5}    17°  BBr_{3}   90°
                      SiCl_{3}   59°  SiBr_{4} 153°
                      PCl_{3}    76°  PBr_{3}  175°
                      SbCl_{3}  223°  SbBr_{3} 275°
                      BiCl_{3}  447°  BiBr_{3} 453°
                      SnCl_{4}  606°  SnBr_{4} 619°
                      ZnCl_{2}  730°  ZnBr_{2} 650°

       Thus for all the more volatile compounds the replacement of
       chlorine by bromine raises the boiling point, but in the ease of
       ZnX_{2} it lowers it (Chapter XV. Note 19).

Their oxygen compounds exhibit a similar analogy. Only fluorine does not
give any oxygen compounds. The iodine and bromine compounds corresponding
with HClO_{3} and HClO_{4} are HBrO_{3} and HBrO_{4}, HIO_{3} and
HIO_{4}. On comparing the properties of these acids we can even predict
that fluorine will not form any oxygen compound. For iodine is easily
oxidised--for instance, by nitric acid--whilst chlorine is not directly
oxidised. The oxygen acids of iodine are comparatively more stable than
those of chlorine; and, generally speaking, the affinity of iodine for
oxygen is much greater than that of chlorine. Here also bromine occupies
an intermediate position. In fluorine we may therefore expect a still
smaller affinity for oxygen than in chlorine--and up to now it has not
been combined with oxygen. If any oxygen compounds of fluorine should be
obtained, they will naturally be exceedingly unstable. The relation of
these elements to hydrogen is the reverse of the above. Fluorine has so
great an affinity for hydrogen that it decomposes water at the ordinary
temperature; whilst iodine has so little affinity for hydrogen that
hydriodic acid, HI, is formed with difficulty, is easily decomposed, and
acts as a reducing agent in a number of cases.

From the form of their compounds the halogens are _univalent elements_
with respect to hydrogen and septivalent with respect to oxygen, N being
trivalent to hydrogen (it gives NH_{3}) and quinqui-valent to oxygen (it
gives N_{2}O_{5}), and C being quadrivalent to both H and O as it forms
CH_{4} and CO_{2}. And as not only their oxygen compounds, but also their
hydrogen compounds, have acid properties, the halogens are _elements_
of an exclusively _acid character_. Such metals as sodium, potassium,
barium only give basic oxides. In the case of nitrogen, although it
forms acid oxides, still in ammonia we find that capacity to give an
alkali with hydrogen which indicates a less distinctly acid character
than in the halogens. In no other elements is the acid-giving property
so strongly developed as in the halogens.

In describing certain peculiarities characterising the halogens, we
shall at every step encounter a confirmation of the above-mentioned
general relations.

As _fluorine_ decomposes water with the evolution of oxygen, F_{2} +
H_{2}O = 2HF + O, for a long time all efforts to obtain it in free state
by means of methods similar to those for the preparation of chlorine
proved fruitless.[48] Thus by the action of hydrofluoric acid on
manganese peroxide, or by decomposing a solution of hydrofluoric acid by
an electric current, either oxygen or a mixture of oxygen and fluorine
were obtained instead of fluorine. Probably a certain quantity of
fluorine[48 bis] was set free by the action of oxygen or an electric
current on incandescent and fused calcium fluoride, but at a high
temperature fluorine acts even on platinum, and therefore it was not
obtained. When chlorine acted on silver fluoride, AgF, in a vessel of
natural fluor spar, CaF_{2}, fluorine was also liberated; but it was
mixed with chlorine, and it was impossible to study the properties of
the resultant gas. Brauner (1881) also obtained fluorine by igniting
cerium fluoride, 2CeF_{4} = 2CeF_{3} + F_{2}; but this, like all
preceding efforts, only showed fluorine to be a gas which decomposes
water, and is capable of acting in a number of instances like chlorine,
but gave no possibility of testing its properties. It was evident that it
was necessary to avoid as far as possible the presence of water and a
rise of temperature; this Moissan succeeded in doing in 1886. He
decomposed anhydrous hydrofluoric acid, liquefied at a temperature of
-23° and contained in a U-shaped tube (to which a small quantity of
potassium fluoride had been added to make it a better conductor), by the
action of a powerful electric current (twenty Bunsen's elements in
series). Hydrogen was then evolved at the negative pole, and fluorine
appeared at the positive pole (of iridium platinum) as a pale green gas
which decomposed water with the formation of ozone and hydrofluoric acid,
and combined directly with silicon (forming silicon fluoride, SiF_{4}),
boron (forming BF_{3}), sulphur, &c. Its density (H = 1) is 18, so that
its molecule is F_{2}. But the action of fluorine on metals at the
ordinary temperature is comparatively feeble, because the metallic
fluoride formed coats the remaining mass of the metals; it is, however,
completely absorbed by iron. Hydrocarbons (such as naphtha), alcohol,
&c., immediately absorb fluorine, with the formation of hydrofluoric
acid. Fluorine when mixed with hydrogen can easily be made to explode
violently, forming hydrofluoric acid.[49]

  [48] Even before free fluorine was obtained (1886) it was evident from
       experience gained in the efforts made to obtain it, and from
       analogy, that it would decompose water (_see_ first Russian
       edition of the _Principles of Chemistry_).

  [48 bis] It is most likely that in this experiment of Fremy's, which
       corresponds with the action of oxygen on calcium chloride,
       fluorine was set free, but that a converse reaction also
       proceeded, CaO + F_{2} = CaF_{2} + O--that is, the calcium
       distributed itself between the oxygen and fluorine. MnF_{4}, which
       is capable of splitting up into MnF_{2} and F_{2}, is without
       doubt formed by the action of a strong solution of hydrofluoric
       acid on manganese peroxide, but under the action of water the
       fluorine gives hydrofluoric acid, and probably this is aided by
       the affinity of the manganese fluoride and hydrofluoric acid. In
       all the attempts made (by Davy, Knox, Louget, Fremy, Gore, and
       others) to decompose fluorides (those of lead, silver, calcium,
       and others) by chlorine, there were doubtless also cases of
       distribution, a portion of the metal combined with chlorine and a
       portion of the fluorine was evolved; but it is improbable that any
       decisive results were obtained. Fremy probably obtained fluorine,
       but not in a pure state.

  [49] According to Moissan, fluorine is disengaged by the action of an
       electric current on fused hydrogen potassium fluoride, KHF_{2}.
       The present state of chemical knowledge is such that the knowledge
       of the properties of an element is much more general than the
       knowledge of the free element itself. It is useful and
       satisfactory to learn that even fluorine in the free state has not
       succeeded in eluding experiment and research, that the efforts to
       isolate it have been crowned with success, but the sum total of
       chemical data concerning fluorine as an element gains but little
       by this achievement. The gain will, however, be augmented if it be
       now possible to subject fluorine to a comparative study in
       relation to oxygen and chlorine. There is particular interest in
       the phenomena of the distribution of fluorine and oxygen, or
       fluorine and chlorine, competing under different conditions and
       relations. We may add that Moissan (1892) found that free fluorine
       decomposes H_{2}S, HCl, HBr, CS_{2}, and CNH with a flash; it does
       not act upon O_{2}, N_{2}, CO, and CO_{2}; Mg, Al, Ag, and Ni,
       when heated, burn in it, as also do S, Se, P (forms PF_{5}); it
       reacts upon H_{2} even in the dark, with the evolution of 366·00
       units of heat. At a temperature of -95°, F_{2} still retains its
       gaseous state. Soot and carbon in general (but not the diamond)
       when heated in gaseous fluorine form _fluoride of carbon_, CF_{4}
       (Moissan, 1890); this compound is also formed at 300° by the
       double decomposition of CCl_{4} and AgF; it is a gas which
       liquefies at 10° under a pressure of 5 atmospheres. With an
       alcoholic solution of KHO, CF_{4} gives K_{2}CO_{3}, according to
       the equation CF_{4} + 6KHO = K_{2}CO_{3} + 4KF + 3H_{2}O. CF_{4}
       is not soluble in water, but it is easily soluble in CCl_{4} and
       alcohol.

In 1894 Brauner obtained fluorine directly by igniting the easily
formed[49 bis] double lead salt HF,3KF,PbF_{4}, which first, at 230°,
decomposes with the evolution of HF, and then splits up forming
3KF,PbF_{2} and fluorine F_{2}, which is recognised by the fact that it
liberates iodine from KI and easily combines with silicon, forming
SiF_{4}. This method gives chemically pure fluorine, and is based upon
the breaking up of the higher compound--tetrafluoride of lead, PbF_{4},
corresponding to PbO_{2}, into free fluorine, F_{2}, and the lower more
stable form--bifluoride of lead, PbF_{2}, which corresponds to PbO; that
is, this method resembles the ordinary method of obtaining chlorine by
means of MnO_{2}, as MnCl_{4} here breaks up into MnCl_{2} and chlorine,
just as PbF_{4} splits up into PbF_{2} and fluorine.

  [49 bis] T. Nikolukin (1885) and subsequently Friedrich and Classen
       obtained PbCl_{4} and a double ammonium salt of tetrachloride of
       lead (starting from the binoxide), PbCl_{4}2NH_{4}Cl; Hutchinson
       and Pallard obtained a similar salt of acetic acid (1893)
       corresponding to PbX_{4} by treating red lead with strong acetic
       acid; the composition of this salt is Pb(C_{2}H_{3}O_{2})_{4}; it
       melts (and decomposes) at about 175°. Brauner (1894) obtained a
       salt corresponding to tetrafluoride of lead, PbF_{4}, and the acid
       corresponding to it, H_{4}PbF_{8}. For example, by treating
       potassium plumbate (Chapter XVIII. Note 55) with strong HF, and
       also the above-mentioned tetra-acetate with a solution of KHF_{2},
       Brauner obtained crystalline HK_{3}PbF_{8}--i.e. the salt from
       which he obtained fluorine.

Among the compounds of fluorine, calcium fluoride, CaF_{2}, is somewhat
widely distributed in nature as fluor spar,[50] whilst _cryolite_, or
aluminium sodium fluoride, Na_{3}AlF_{6}, is found more rarely (in large
masses in Greenland). Cryolite, like fluor spar, is also insoluble in
water, and gives hydrofluoric acid with sulphuric acid. Small quantities
of fluorine have also in a number of cases been found in the bodies of
animals, in the blood, urine, and bones. If fluorides occur in the bodies
of animals, they must have been introduced in food, and must occur in
plants and in water. And as a matter of fact river, and especially sea,
water always contains a certain, although small, quantity of fluorine
compounds.

  [50] It is called spar because it very frequently occurs as crystals of
       a clearly laminar structure, and is therefore easily split up into
       pieces bounded by planes. It is called fluor spar because when
       used as a flux it renders ores fusible, owing to its reacting with
       silica, SiO_{2} + 2CaF_{2} = 2CaO + SiF_{4}; the silicon fluoride
       escapes as a gas and the lime combines with a further quantity of
       silica, and gives a vitreous slag. Fluor spar occurs in mineral
       veins and rocks, sometimes in considerable quantities. It always
       crystallises in the cubic system, sometimes in very large
       semi-transparent cubic crystals, which are colourless or of
       different colours. It is insoluble in water. It melts under the
       action of heat, and crystallises on cooling. The specific gravity
       is 3·1. When steam is passed over incandescent fluor spar, lime
       and hydrofluoric acid are formed: CaF_{2} + H_{2}O = CaO + 2HF. A
       double decomposition is also easily produced by fusing fluor spar
       with sodium or potassium hydroxides, or potash, or even with their
       carbonates; the fluorine then passes over to the potassium or
       sodium, and the oxygen to the calcium. In solutions--for example,
       Ca(NO_{3})_{2} + 2KF = CaF_{2} (precipitate) + 2KNO_{3} (in
       solution)--the formation of calcium fluoride takes place, owing to
       its very sparing solubility. 26,000 parts of water dissolve one
       part of fluor spar.

Hydrifluoric acid, HF, cannot be obtained from fluor spar in glass
retorts, because glass is acted on by and destroys the acid. It is
prepared in lead vessels, and when it is required pure, in platinum
vessels, because lead also acts on hydrofluoric acid, although only very
feebly on the surface, and when once a coating of fluoride and sulphate
of lead is formed no further action takes place. Powdered fluor spar and
sulphuric acid evolve hydrofluoric acid (which fumes in the air) even at
the ordinary temperature, CaF_{2} + H_{2}SO_{4} = CaSO_{4} + 2HF. At 130°
fluor spar is completely decomposed by sulphuric acid. The acid is then
evolved as vapour, which may be condensed by a freezing mixture into an
anhydrous acid. The condensation is aided by pouring water into the
receiver of the condenser, as the acid is easily soluble in cold water.

In the liquid anhydrous form hydrofluoric acid boils at +19°, and its
specific gravity at 12·8° = 0·9849.[51] It dissolves in water with the
evolution of a considerable amount of heat, and gives a solution of
constant boiling point which distils over at 120°; showing that the acid
is able to combine with water. The specific gravity of the compound is
1·15, and its composition HF,2H_{2}O.[52] With an excess of water a
dilute solution distils over first. The aqueous solution and the acid
itself must be kept in platinum vessels, but the dilute acid may be
conveniently preserved in vessels made of various organic materials,
such as gutta-percha, or even in glass vessels having an interior
coating of paraffin. Hydrofluoric acid does not act on hydrocarbons
and many other substances, but it acts in a highly corrosive manner
on metals, glass, porcelain, and the majority of rock substances.[53]
It also attacks the skin, and is distinguished by its poisonous
properties, so that in working with the acid a strong draught must be
kept up, to prevent the possibility of the fumes being inhaled. The
non-metals do not act on hydrofluoric acid, but all metals--with the
exception of mercury, silver, gold, and platinum, and, to a certain
degree, lead--decompose it with the evolution of hydrogen. With bases
it gives directly metallic fluorides, and behaves in many respects like
hydrochloric acid. There are, however, several distinct individual
differences, which are furthermore much greater than those between
hydrochloric, hydrobromic, and hydriodic acids. Thus the silver
compounds of the latter are insoluble in water, whilst silver fluoride
is soluble. Calcium fluoride, on the contrary, is insoluble in water,
whilst calcium chloride, bromide, and iodide are not only soluble, but
attract water with great energy. Neither hydrochloric, hydrobromic, nor
hydriodic acid acts on sand and glass, whilst hydrofluoric acid corrodes
them, forming gaseous silicon fluoride. The other halogen acids only form
normal salts, KCl, NaCl, with Na or K, whilst hydrofluoric acid gives
acid salts, for instance HKF_{2} (and by dissolving KF in liquid HF,
KHF_{2}2HF is obtained). This latter property is in close connection with
the fact that at the ordinary temperature the vapour density of
hydrofluoric acid is nearly 20, which corresponds with a formula
H_{2}F_{2}, as Mallet (1881) showed; but a depolymerisation occurs with a
rise of temperature, and the density approaches 10, which answers to the
formula HF.[54]

  [51] According to Gore. Fremy obtained anhydrous hydrofluoric acid by
       decomposing lead fluoride at a red heat, by hydrogen, or by
       beating the double salt HKF_{2}, which easily crystallises (in
       cubes) from a solution of hydrofluoric acid, half of which has
       been saturated with potassium hydroxide. Its vapour density
       corresponds to the formula HF.

  [52] This composition corresponds to the crystallo-hydrate HCl,2H_{2}O.
       All the properties of hydrofluoric acid recall those of
       hydrochloric acid, and therefore the comparative ease with which
       hydrofluoric acid is liquefied (it boils at +19°, hydrochloric
       acid at -35°) must be explained by a polymerisation taking place at
       low temperatures, as will be afterwards explained, H_{2}F_{2} being
       formed, and therefore in a liquid state it differs from
       hydrochloric acid, in which a phenomenon of a similar kind has not
       yet been observed.

  [53] The corrosive action of hydrofluoric acid on glass and similar
       siliceous compounds is based upon the fact that it acts on silica,
       SiO_{2}, as we shall consider more fully in describing that
       compound, forming gaseous silicon fluoride, SiO_{2} + 4HF =
       SiF_{4} + 2H_{2}O. Silica, on the other hand, forms the binding
       (acid) element of glass and of the mass of mineral substances
       forming the salts of silica. When it is removed the cohesion is
       destroyed. This is made use of in the arts, and in the laboratory,
       for etching designs and scales, &c., on glass. In _engraving on
       glass_ the surface is covered with a varnish composed of four
       parts of wax and one part of turpentine. This varnish is not acted
       on by hydrofluoric acid, and it is soft enough to allow of designs
       being drawn upon it whose lines lay bare the glass. The drawing is
       made with a steel point, and the glass is afterwards laid in a
       lead trough in which a mixture of fluor spar and sulphuric acid is
       placed. The sulphuric acid must be used in considerable excess, as
       otherwise transparent lines are obtained (owing to the formation
       of hydrofluosilicic acid). After being exposed for some time, the
       varnish is removed (melted) and the design drawn by the steel
       point is found reproduced in dull lines. The drawing may be also
       made by the direct application of a mixture of a silicofluoride
       and sulphuric acid, which forms hydrofluoric acid.

  [54] Mallet (1881) determined the density at 30° and 100°, previous to
       which Gore (1869) had determined the vapour density at 100°,
       whilst Thorpe and Hambly (1888) made fourteen determinations
       between 26° and 88°, and showed that within this limit of
       temperature the density gradually diminishes, just like the vapour
       of acetic acid, nitrogen dioxide, and others. The tendency of HF
       to polymerise into H_{2}F_{2} is probably connected with the
       property of many fluorides of forming acid salts--for example,
       KHF_{2} and H_{2}SiF_{6}. We saw above that HCl has the same
       property (forming, for instance, H_{2}PtCl_{6}, &c., p. 457), and
       hence this property of hydrofluoric acid does not stand isolated
       from the properties of the other halogens.

The analogy between chlorine and the other two halogens, bromine and
iodine, is much more perfect. Not only have their hydrates or halogen
acids much in common, but they themselves resemble chlorine in many
respects,[55] and even the properties of the corresponding metallic
compounds of bromine and iodine are very much alike. Thus, the chlorides,
bromides, and iodides of sodium and potassium crystallise in the cubic
system, and are soluble in water; the chlorides of calcium, aluminium,
magnesium, and barium are just as soluble in water as the bromides and
iodides of these metals. The iodides and bromides of silver and lead are
sparingly soluble in water, like the chlorides of these metals. The
oxygen compounds of bromine and iodine also present a very strong analogy
to the corresponding compounds of chlorine. A hypobromous acid is known
corresponding with hypochlorous acid. The salts of this acid have the
same bleaching property as the salts of hypochlorous acid. Iodine was
discovered in 1811 by Courtois in kelp, and was shortly afterwards
investigated by Clement, Gay-Lussac, and Davy. Bromine was discovered in
1826 by Balard in the mother liquor of sea water.

  [55] For instance, the experiment with Dutch metal foil (Note 16) may
       be made with bromine just as well as with chlorine. A very
       instructive experiment on the direct combination of the halogens
       with metals maybe made by throwing a small piece (a shaving) of
       aluminium into a vessel containing liquid bromine; the aluminium,
       being lighter, floats on the bromine, and after a certain time
       reaction sets in accompanied by the evolution of heat, light, and
       fumes of bromine. The incandescent piece of metal moves rapidly
       over the surface of the bromine in which the resultant aluminium
       bromide dissolves. For the sake of comparison we will proceed to
       cite several thermochemical data (Thomsen) for analogous actions
       of (1) chlorine, (2) bromine, and (3) iodine, with respect to
       metals; the halogen being expressed by the symbol X, and the plus
       sign connecting the reacting substances. All the figures are given
       in thousands of calories, and refer to molecular quantities in
       grams and to the ordinary temperature:--

                                          1     2     3
                        K_{2}  + X_{2}   211   191   160
                        Na_{2} + X_{2}   195   172   138
                        Ag_{2} + X_{2}    59    45    28
                        Hg_{2} + X_{2}    83    68    48
                        Hg     + X_{2}    63    51    34
                        Ca     + X_{2}   170   141    --
                        Ba     + X_{2}   195   170    --
                        Zn     + X_{2}    97    76    49
                        Pb     + X_{2}    83    64    40
                        Al     + X_{2}   161   120    70

       We may remark that the latent heat of vaporisation of the
       molecular weight Br_{2} is about 7·2, and of iodine 6·0 thousand
       heat units, whilst the latent heat of fusion of Br_{2} is about
       0·3, and of I_{2} about 3·0 thousand heat units. From this it is
       evident that the difference between the amounts of heat evolved
       does not depend on the difference in physical state. For instance,
       the vapour of iodine in combining with Zn to form ZnI_{2} would
       give 48 + 8 + 3, or about sixty thousand heat units, or 1-1/2
       times less than Zn + Cl_{2}.

_Bromine_ and iodine, like chlorine, occur in sea water in combination
with metals. However, the amount of bromides, and especially of iodides,
in sea water is so small that their presence can only be discovered by
means of sensitive reactions.[56] In the extraction of salt from sea
water the bromides remain in the mother liquor. Iodine and bromine also
occur combined with silver, in admixture with silver chloride, as a rare
ore which is mainly found in America. Certain mineral waters (those of
Kreuznach and Staro-rossüsk) contain metallic bromides and iodides,
always in admixture with an excess of sodium chloride. Those upper strata
of the Stassfurt rock salt (Chapter X.) which are a source of potassium
salts also contain metallic bromides,[57] which collect in the mother
liquors left after the crystallisation of the potassium salts; and this
now forms the chief source (together with certain American springs) of
the bromine in common use. Bromine may be easily liberated from a mixture
of bromides and chlorides, owing to the fact that chlorine displaces
bromine from its compounds with sodium, magnesium, calcium, &c. A
colourless solution of bromides and chlorides turns an orange colour
after the passage of chlorine, owing to the disengagement of bromine.[58]
Bromine may be extracted on a large scale by a similar method, but it is
simpler to add a small quantity of manganese peroxide and sulphuric acid
to the mother liquid direct. This sets free a portion of the chlorine,
and this chlorine liberates the bromine.

  [56] One litre of sea-water contains about 20 grams of chlorine, and
       about 0·07 gram of bromine. The Dead Sea contains about ten times
       as much bromine.

  [57] But there is no iodine in Stassfurt carnallite.

  [58] The chlorine must not, however, be in large excess, as otherwise
       the bromine would contain chlorine. Commercial bromine not
       unfrequently contains chlorine, as bromine chloride; this is more
       soluble in water than bromine, from which it may thus be freed. To
       obtain pure bromine the commercial bromine is washed with water,
       dried by sulphuric acid, and distilled, the portion coming over at
       58° being collected; the greater part is then converted into
       potassium bromide and dissolved, and the remainder is added to the
       solution in order to separate iodine, which is removed by shaking
       with carbon bisulphide. By heating the potassium bromide thus
       obtained with manganese peroxide and sulphuric acid, bromine is
       obtained quite free from iodine, which, however, is not present in
       certain kinds of commercial bromine (the Stassfurt, for instance).
       By treatment with potash, the bromine is then converted into a
       mixture of potassium bromide and bromate, and the mixture (which
       is in the proportion given in the equation) is distilled with
       sulphuric acid, bromine being then evolved: 5KBr + KBrO_{3} +
       6H_{2}SO_{4} = 6KHSO_{4} + 3H_{2}O + 3Br_{2}. After dissolving the
       bromine in a strong solution of calcium bromide and precipitating
       with an excess of water, it loses all the chlorine it contained,
       because chlorine forms calcium chloride with CaBr_{2}.

Bromine is a _dark brown liquid_, giving brown fumes, and having a
poisonous suffocating smell, whence its name (from the Greek [Greek:
brômos], signifying evil smelling). The vapour density of bromine shows
that its molecule is Br_{2}. In the cold bromine freezes into brown-grey
scales like iodine. The melting point of pure bromine is -7°·05.[59] The
density of liquid bromide at 0° is 3·187, and at 15° about 3·0. The
boiling point of bromine is about 58°·7. Bromine, like chlorine, is
soluble in water; 1 part of bromine at 5° requires 27 parts of water, and
at 15° 29 parts of water. The aqueous solution of bromine is of an
orange colour, and when cooled to -2° yields crystals containing 10
molecules of water to 1 molecule of bromine.[60] Alcohol dissolves a
greater quantity of bromine, and ether a still greater amount. But after
a certain time products of the action of the bromine on these organic
substances are formed in the solutions. Aqueous solutions of the bromides
also absorb a large amount of bromine.

  [59] There has long existed a difference of opinion as to the melting
       point of pure bromine. By some investigators (Regnault, Pierre) it
       was given as between -7° and -8°, and by others (Balard, Liebig,
       Quincke, Baumhauer) as between -20° and -25°. There is now no doubt,
       thanks more especially to the researches of Ramsay and Young
       (1885), that pure bromine melts at about -7°. This figure is not
       only established by direct experiment (Van der Plaats confirmed
       it), but also by means of the determination of the vapour
       tensions. For solid bromine the vapour tension _p_ in mm. at _t_
       was found to be--

                  _p_ =  20     25   30   35  40    45 mm.
                  _t_ = -16°·6 -14° -12° -10° -8·5° -7°

       For liquid bromine--

             _p_ =   50    100    200    400    600   760 mm.
             _t_ = -5°·0   8°·2  23°·4  40°·4  51°·9  58°·7

       These curves intersect at -7°·05. Besides which, in comparing the
       vapour tension of many liquids (for example, those given in
       Chapter II., Note 27), Ramsay and Young observed that the ratio of
       the absolute temperatures (_t_ + 273) corresponding with equal
       tension _varies_ for every pair of substances in rectilinear
       proportion in dependence upon _t_, and, therefore, for the above
       pressure _p_, Ramsay and Young determined the ratio of _t_ + 273
       for water and bromine, and found that the straight lines
       expressing these ratios for liquid and solid bromine intersect
       also at 7°·05; thus, for example, for solid bromine--

             _p_  =   20       25       30       35       40      45
       273 + _t_  =  256·4    259      261      263      264·6    266
       273 + _t_´ =  295·3    299      302·1    304·8    307·2    309·3
             _c_  =    1·152    1·154    1·157    1·159    1·161    1·163

       where _t_´ indicates the temperature of water corresponding with a
       vapour tension _p_, and where _c_ is the ratio of 273 + _t_´ to
       273 + _t_. The magnitude of _c_ is evidently expressed with great
       accuracy by the straight line _c_ = 1·1703 + 0·0011_t_. In exactly
       the same way we find the ratio for liquid bromine and water to be
       _c__{1} = 1·1585 + 0·00057t. The intersection of these straight
       lines in fact corresponds with -7°·06, which again confirms the
       melting point given above for bromine. In this manner it is
       possible with the existing store of data to accurately establish
       and _verify_ the melting point of substances. Ramsay and Young
       established the thermal constants of iodine by exactly the same
       method.

  [60] The observations made by Paterno and Nasini (by Raoult's method,
       Chapter I. Note 49) on the temperature of the formation of ice
       (-1°·115, with 1·391 gram of bromine in 100 grams of water) in an
       aqueous solution of bromine, showed that bromine is contained in
       solutions as the molecule Br_{2}. Similar experiments conducted on
       iodine (Kloboukoff 1889 and Beckmann 1890) show that in solution
       the molecule is I_{2}.

       B. Roozeboom investigated the hydrate of bromine as completely as
       the hydrate of chlorine (Notes 9, 10). The temperature of the
       complete decomposition of the hydrate is +6°·2; the density of
       Br_{2},10H_{2}O = 1·49.

With respect to _iodine_, it is almost exclusively extracted from the
mother liquors after the crystallisation of natural sodium nitrate (Chili
saltpetre) and from the ashes of the sea-weed cast upon the shores of
France, Great Britain, and Spain, sometimes in considerable quantities,
by the high tides. The majority of these sea-weeds are of the genera
_Fucus_, _Laminaria_, &c. The fused ashes of these sea-weeds are called
'kelp' in Scotland and 'varech' in Normandy. A somewhat considerable
quantity of iodine is contained in these sea-weeds. After being burnt
(or subjected to dry distillation) an ash is left which chiefly contains
salts of potassium, sodium, and calcium. The metals occur in the sea-weed
as salts of organic acids. On being burnt these organic salts are
decomposed, forming carbonates of potassium and sodium. Hence, sodium
carbonate is found in the ash of sea plants. The ash is dissolved in hot
water, and on evaporation sodium carbonate and other salts separate, but
a portion of the substances remains in solution. These mother liquors
left after the separation of the sodium carbonate contain chlorine,
bromine, and iodine in combination with metals, the chlorine and iodine
being in excess of the bromine. 13,000 kilos of kelp give about 1,000
kilos of sodium carbonate and 15 kilos of iodine.

The liberation of the iodine from the mother liquor is effected with
comparative ease, because chlorine disengages iodine from potassium
iodide and its other combinations with the metals. Not only chlorine,
but also sulphuric acid, liberates iodine from sodium iodide. Sulphuric
acid, in acting on an iodide, sets hydriodic acid free, but the latter
easily decomposes, especially in the presence of substances capable
of evolving oxygen, such as chromic acid, nitrous acid, and even
ferric salts.[61] Owing to its sparing solubility in water, the iodine
liberated separates as a precipitate. To obtain pure iodine it is
sufficient to distil it, and neglect the first and last portions of
the distillate, the middle portion only being collected. Iodine passes
directly from a state of vapour into a crystalline form, and settles on
the cool portions of the apparatus in tabular crystals, having a black
grey colour and metallic lustre.[62]

  [61] In general, 2HI + O = I_{2} + H_{2}O, if the oxygen proceed from a
       substance from which it is easily evolved. For this reason
       compounds corresponding with the higher stages of oxidation or
       chlorination frequently give a lower stage when treated with
       hydriodic acid. Ferric oxide, Fe_{2}O_{3}, is a higher oxide, and
       ferrous oxide, FeO, a lower oxide; the former corresponds with
       FeX_{3}, and the latter with FeX_{2}, and this passage from the
       higher to the lower takes place under the action of hydriodic
       acid. Thus hydrogen peroxide and ozone (Chapter IV.) are able to
       liberate iodine from hydriodic acid. Compounds of copper oxide,
       CuO or CuX_{2}, give compounds of the suboxide Cu_{2}O, or CuX.
       Even sulphuric acid, which corresponds to the higher stage SO_{3},
       is able to act thus, forming the lower oxide SO_{2}. The
       liberation of iodine from hydriodic acid proceeds with still
       greater ease under the action of substances capable of disengaging
       oxygen. In practice, many methods are employed for liberating
       iodine from acid liquids containing, for example, sulphuric acid
       and hydriodic acid. The higher oxides of nitrogen are most
       commonly used; they then pass into nitric oxide. Iodine may even
       be disengaged from hydriodic acid by the action of iodic acid, &c.
       But there is a limit in these reactions of the oxidation of
       hydriodic acid because, under certain conditions, especially in
       dilute solutions, the iodine set free is itself able to act as an
       oxidising agent--that is, it exhibits the character of chlorine,
       and of the halogens in general, to which we shall again have
       occasion to refer. In Chili, where a large quantity of iodine is
       extracted in the manufacture of Chili nitre, which contains
       NaIO_{3}, it is mixed with the acid and normal sulphites of sodium
       in solution; the iodine is then precipitated according to the
       equation 2NaIO_{3} + 3Na_{2}SO_{3} + 2NaHSO_{3} = 5Na_{2}SO_{4}
       +I_{2} + H_{2}O. The iodine thus obtained is purified by
       sublimation.

  [62] For the final purification of iodine, Stas dissolved it in a
       strong solution of potassium iodide, and precipitated it by the
       addition of water (_see_ Note 58).

The specific gravity of the crystals of iodine is 4·95. It melts at 114°
and boils at 184°. Its vapour is formed at a much lower temperature, and
is of a violet colour, whence iodine receives its name ([Greek: ioeidês],
violet). The smell of iodine recalls the characteristic smell of
hypochlorous acid; it has a sharp sour taste. It destroys the skin and
organs of the body, and is therefore frequently employed for cauterising
and as an irritant for the skin. In small quantities it turns the skin
brown, but the coloration disappears after a certain time, partly owing
to the volatility of the iodine. Water dissolves only 1/5000 part of
iodine. A brown solution is thus obtained, which bleaches, but much more
feebly than bromine and chlorine. Water which contains salts, and
especially iodides, in solution dissolves iodine in considerable
quantities, and the resultant solution is of a dark brown colour. Pure
alcohol dissolves a small amount of iodine, and in so doing acquires a
brown colour, but the solubility of iodine is considerably increased by
the presence of a small quantity of an iodine compound--for instance,
ethyl iodide--in the alcohol.[63] Ether dissolves a larger amount of
iodine than alcohol; but iodine is particularly soluble in liquid
hydrocarbons, in carbon bisulphide, and in chloroform. A small quantity
of iodine dissolved in carbon bisulphide tints it rose-colour, but in a
somewhat larger amount it gives a violet colour. Chloroform (quite free
from alcohol) is also tinted rose colour by a small amount of iodine.
This gives an easy means for detecting the presence of free iodine in
small quantities. The blue coloration which free iodine gives with
_starch_ may also, as has already been frequently mentioned (_see_
Chapter IV.), serve for the detection of iodine.

  [63] The solubility of iodine in solutions containing iodides, and
       compounds of iodine in general, may serve, on the one hand, as an
       indication that solution is due to a similarity between the
       solvent and dissolved substance, and, on the other hand, as an
       indirect proof of that view as to solutions which was cited in
       Chapter I., because in many instances unstable highly iodised
       compounds, resembling crystallo-hydrates, have been obtained from
       such solutions. Thus iodide of tetramethylammonium,
       N(CH_{3})_{4}I, combines with I_{2}, and I_{4}. Even a solution of
       iodine in a saturated solution of potassium iodide presents
       indications of the formation of a definite compound KI_{3}. Thus,
       an alcoholic solution of KI_{3} does not give up iodine to carbon
       bisulphide, although this solvent takes up iodine from an
       alcoholic solution of iodine itself (Girault, Jörgensen, and
       others). The instability of these compounds resembles the
       instability of many crystallo-hydrates, for instance of
       HCl,2H_{2}O.

If we compare the four elements, fluorine, chlorine, bromine, and iodine,
we see in them an example of analogous substances which arrange
themselves by their physical properties in the same order as they stand
in respect to their atomic and molecular weights. If the weight of the
molecule be large, the substance has a higher specific gravity, a higher
melting and boiling point, and a whole series of properties depending on
this difference in its fundamental properties. Chlorine in a free state
boils at about -35°, bromine boils at 60°, and iodine only above 180°.
According to Avogadro-Gerhardt's law, the vapour densities of these
elements in a gaseous state are proportional to their atomic weights, and
here, at all events approximately, the densities in a liquid (or solid)
state are also almost in the ratio of their atomic weights. Dividing the
atomic weight of chlorine (35·5) by its specific gravity in a liquid
state (1·3), we obtain a volume = 27, for bromine (80/3·1) 26, and for
iodine also (127/4·9) 26.[64]

  [64] The equality of the atomic volumes of the halogens themselves is
       all the more remarkable because in all the halogen compounds the
       volume augments with the substitution of fluorine by chlorine,
       bromine, and iodine. Thus, for example, the volume of sodium
       fluoride (obtained by dividing the weight expressed by its formula
       by its specific gravity) is about 15, of sodium chloride 27, of
       sodium bromide 32, and of sodium iodide 41. The volume of silicon
       chloroform, SiHCl_{3}, is 82, and those of the corresponding
       bromine and iodine compounds are 108 and 122 respectively. The
       same difference also exists in solutions; for example, NaCl +
       200H_{2}O has a sp. gr. (at 15°/4°) of 1·0106, consequently the
       volume of the solution 3,658·5/1·0106 = 3,620, hence the volume of
       sodium chloride in solution = 3,620-3,603 (this is the volume of
       200 H_{2}O) = 17, and in similar solutions, NaBr = 26 and NaI =
       35.

The metallic bromides and iodides are in the majority of cases, in most
respects analogous to the corresponding chlorides,[65] but chlorine
displaces the bromine and iodine from them, and bromine liberates iodine
from iodides, which is taken advantage of in the preparation of these
halogens. However, the researches of Potilitzin showed that a _reverse_
displacement of chlorine by bromine may occur both in solutions and in
ignited metallic chlorides in an atmosphere of bromine vapour--that is, a
distribution of the metal (according to Berthollet's doctrine) takes
place between the halogens, although however the larger portion, still
unites with the chlorine, which shows its greater affinity for metals as
compared with that of bromine and iodine.[66] The latter, however,
sometimes behave with respect to metallic oxides in exactly the same
manner as chlorine. Gay-Lussac, by igniting potassium carbonate in iodine
vapour, obtained (as with chlorine) an evolution of oxygen and carbonic
anhydride, K_{2}CO_{3} + I_{2} = 2KI + CO_{2} + O, only the reactions
between the halogens and oxygen are more easily reversible with bromine
and iodine than with chlorine. Thus, at a red heat oxygen displaces
iodine from barium iodide. Aluminium iodide burns in a current of oxygen
(Deville and Troost), and a similar, although not so clearly marked,
relation exists for aluminium chloride, and shows that the halogens have
a distinctly smaller affinity for those metals which only form feeble
bases. This is still more the case with the non-metals, which form acids
and evolve much more heat with oxygen than with the halogens (Note 13).
But in all these instances the affinity (and amount of heat evolved) of
iodine and bromine is less than that of chlorine, probably because the
atomic weights are greater. The smaller store of energy in iodine and
bromine is seen still more clearly in the relation of the halogens to
hydrogen. In a gaseous state they all enter, with more or less ease, into
direct combination with gaseous hydrogen--for example, in the presence of
spongy platinum, forming halogen acids, HX--but the latter are far from
being equally stable; hydrogen chloride is the most stable, hydrogen
iodide the least so, and hydrogen bromide occupies an intermediate
position. A very high temperature is required to decompose hydrogen
chloride even partially, whilst hydrogen iodide is decomposed by light
even at the ordinary temperature and very easily by a red heat. Hence the
reaction I_{2} + H_{2} = HI + HI is very easily reversible, and
consequently has a limit, and hydrogen iodide easily dissociates.[67]
Judging by the direct measurement of the heat evolved (22,000 heat units)
in the formation of HCl, the conversion of 2HCl into H_{2} + Cl_{2}
requires the expenditure of 44,000 heat units. The decomposition of 2HBr
into H_{2} + Br_{2} only requires, if the bromine be obtained in a
gaseous state, a consumption of about 24,000 units, whilst in the
decomposition of 2HI into H_{2} + I_{2} as vapour about 3,000 heat units
are _evolved_;[68] these facts, without doubt, stand in causal connection
with the great stability of hydrogen chloride, the easy decomposability
of hydrogen iodide, and the intermediate properties of hydrogen bromide.
From this it would be expected that chlorine is capable of decomposing
water with the evolution of oxygen, whilst iodine has not the energy to
produce this disengagement,[69] although it is able to liberate the
oxygen from the oxides of potassium and sodium, the affinity of these
metals for the halogens being very considerable. For this reason oxygen,
especially in compounds from which it can be evolved readily (for
instance, ClHO, CrO_{3}, &c.), easily decomposes hydrogen iodide. A
mixture of hydrogen iodide and oxygen burns in the presence of an ignited
substance, forming water and iodine. Drops of nitric acid in an
atmosphere of hydrogen iodide cause the disengagement of violet fumes of
iodine and brown fumes of nitric peroxide. In the presence of alkalis and
an excess of water, however, iodine is able to effect oxidation like
chlorine--that is, it decomposes water; the action is here aided by the
affinity of hydrogen iodide for the alkali and water, just as sulphuric
acid helps zinc to decompose water. But the relative instability of
hydriodic acid is best seen in comparing the acids in a gaseous state. If
the halogen acids be dissolved in water, they evolve so much heat that
they approach much nearer to each other in properties. This is seen from
thermochemical data, for in the formation of HX in solution (in a large
excess of water) from the _gaseous_ elements there is _evolved_ for HCl
39,000, for HBr 32,000, and for HI 18,000 heat units.[70] But it is
especially evident from the fact that solutions of hydrogen bromide and
iodide in water have many points in common with solutions of hydrogen
chloride, both in their capacity to form hydrates and fuming solutions of
constant boiling point, and in their capacity to form haloid salts, &c.
by reacting on bases.

  [65] But the density (and also molecular volume, Note 64) of a bromine
       compound is always greater than that of a chlorine compound,
       whilst that of an iodine compound is still greater. The order is
       the same in many other respects. For example, an iodine compound
       has a higher boiling point than a bromine compound, &c.

  [66] A. L. Potilitzin showed that in heating various metallic chlorides
       in a closed tube, with an equivalent quantity of bromine, a
       distribution of the metal between the halogens always occurs, and
       that the amounts of chlorine replaced by the bromine in the
       ultimate product are proportional to the atomic weights of the
       metals taken and inversely proportional to their equivalence.
       Thus, if NaCl + Br be taken, then out of 100 parts of chlorine,
       5·54 are replaced by the bromine, whilst with AgCl + Br 27·28
       parts are replaced. These figures are in the ratio 1 : 4·9, and
       the atomic weights Na : Ag = 1 : 4·7. In general terms, if a
       chloride MCl_{_n_} be taken, it gives with _n_Br a percentage
       substitution = 4M/_n_^2 where M is the atomic weight of the metal.
       This law was deduced from observations on the chlorides of Li, K,
       Na, Ag (_n_ = 1), Ca, Sr, Ba, Co, Ni, Hg, Pb (_n_ = 2), Bi (_n_ =
       3), Sn (_n_ = 4), and Fe_{2} (_n_ = 6).

       In these determinations of Potilitzin we see not only a brilliant
       confirmation of Berthollet's doctrine, but also the first effort
       to directly determine the affinities of elements by means of
       displacement. The chief object of these researches consisted in
       proving whether a displacement occurs in those cases where heat is
       absorbed, and in this instance it should be absorbed, because the
       formation of all metallic bromides is attended with the evolution
       of less heat than that of the chlorides, as is seen by the figures
       given in Note 55.

       If the mass of the bromine be increased, then the amount of
       chlorine displaced also increases. For example, if masses of
       bromine of 1 and 4 equivalents act on a molecule of sodium
       chloride, then the percentages of the chlorine displaced will be
       6·08 p.c. and 12·46 p.c.; in the action of 1, 4, 25, and 100
       molecules of bromine on a molecule of barium chloride, there will
       be displaced 7·8, 17·6, 35·0, and 45·0 p.c. of chlorine. If an
       equivalent quantity of hydrochloric acid act on metallic bromides
       in closed tubes, and in the absence of water at a temperature of
       300°, then the percentages of the substitution of the bromine by
       the chlorine in the double decomposition taking place between
       univalent metals are inversely proportional to their atomic
       weights. For example, NaBr + HCl gives at the limit 21 p.c. of
       displacement, KCl 12 p.c. and AgCl 4-1/4 p.c. Essentially the same
       action takes place in an aqueous solution, although the phenomenon
       is complicated by the participation of the water. The reactions
       proceed spontaneously in one or the other direction at the
       ordinary temperature but at different _rates_. In the action of a
       dilute solution (1 equivalent per 5 litres) of sodium chloride on
       silver bromide at the ordinary temperature the amount of bromine
       replaced in six and a half days is 2·07 p.c., and with potassium
       chloride 1·5 p.c. With an excess of the chloride the magnitude of
       the substitution increases. These conversions also proceed with
       the absorption of heat. The reverse reactions evolving heat
       proceed incomparably more rapidly, but also to a certain limit;
       for example, in the reaction AgCl + RBr the following percentages
       of silver bromide are formed in different times:

               hours    2       3      22      96     120
                K     79·82   87·4    88·22    --     94·21
                Na    83·63   90·74   91·70   95·49    --

       That is, the conversions which are accompanied by an evolution of
       heat proceed with very much greater rapidity than the reverse
       conversions.

  [67] _The dissociation of hydriodic acid_ has been studied in detail by
       Hautefeuille and Lemoine, from whose researches we extract the
       following information. The decomposition of hydriodic acid is
       decided, but proceeds slowly at 180°; the rate and limit of
       decomposition increase with a rise of temperature. The reverse
       action--that is, I_{2} + H_{2} = 2HI--proceeds not only under the
       influence of spongy platinum (Corenwinder), which also accelerates
       the decomposition of hydriodic acid, but also by itself, although
       slowly. The limit of the reverse reaction remains the same with or
       without spongy platinum. An increase of pressure has a very
       powerful accelerative effect on the rate of formation of hydriodic
       acid, and therefore spongy platinum by condensing gases has the
       same effect as increase of pressure. At the atmospheric pressure
       the decomposition of hydriodic acid reaches the limit at 250° in
       several months, and at 440° in several hours. The limit at 250° is
       about 18 p.c. of decomposition--that is, out of 100 parts of
       hydrogen previously combined in hydriodic acid, about 18 p.c. may
       be disengaged at this temperature (this hydrogen may be easily
       measured, and the measure of dissociation determined), but not
       more; the limit at 440° is about 26 p.c. If the pressure under
       which 2HI passes into H_{2} + I_{2} be 4-1/2 atmospheres, then the
       limit is 24 p.c.; under a pressure of 1/5 atmosphere the limit is
       29 p.c. The small influence of pressure on the dissociation of
       hydriodic acid (compared with N_{2}O_{4}, Chapter VI. Note 46) is
       due to the fact that the reaction 2HI = I_{2} + H_{2} is not
       accompanied by a change of volume. In order to show the influence
       of time, we will cite the following figures referring to 350°: (1)
       Reaction H_{2} + I_{2}; after 3 hours, 88 p.c. of hydrogen remained
       free; 8 hours, 69 p.c.; 34 hours, 48 p.c.; 76 hours, 29 p.c.; and
       327 hours, 18·5 p.c. (2) The reverse decomposition of 2HI; after 9
       hours, 3 p.c. of hydrogen was set free, and after 250 hours 18·6
       p.c.--that is, the limit was reached. The addition of extraneous
       hydrogen diminishes the limit of the reaction of decomposition, or
       increases the formation of hydriodic acid from iodine and
       hydrogen, as would be expected from Berthollet's doctrine (Chapter
       X.). Thus at 440° 26 p.c. of hydriodic acid is decomposed if there
       be no admixture of hydrogen, while if H_{2} be added, then at the
       limit only half as large a mass of HI is decomposed. Therefore, if
       an infinite mass of hydrogen be added there will be no
       decomposition of the hydriodic acid. Light aids the decomposition
       of hydriodic acid very powerfully. At the ordinary temperature 80
       p.c. is decomposed under the influence of light, whilst under the
       influence of heat alone this limit corresponds with a very high
       temperature. The distinct action of light, spongy platinum, and of
       impurities in glass (especially of sodium sulphate, which
       decomposes hydriodic acid), not only render the investigations
       difficult, but also show that in reactions like 2HI = I_{2} + H_{2},
       which are accompanied by slight heat effects, all foreign and
       feeble influences may strongly affect the progress of the action
       (Note 47).

  [68] The thermal determinations of Thomsen (at 18°) gave in thousands
       of calories, Cl + H = +22, HCl + Aq (that is, on dissolving HCl in
       a large amount of water) = +17·3, and therefore H + Cl + Aq =
       +39·3. In taking molecules, all these figures must be doubled. Br
       + H = +8·4; HBr + Aq = 19·9; H + Br + Aq = +28·3. According to
       Berthelot 7·2 are required for the vaporisation of Br_{2}, hence
       Br_{2} + H_{2} = 16·8 + 7·2 = +24, if Br_{2} be taken as vapour
       for comparison with Cl_{2}. H + I =-6·0, HI + Aq = 19·2; H + I +
       Aq= +13·2, and, according to Berthelot, the heat of fusion of
       I_{2} = 3·0, and of vaporisation 6·0 thousand heat units, and
       therefore I_{2} + H_{2} =-2(6·0) + 3 + 6 =-3·0, if the iodine be
       taken as vapour. Berthelot, on the basis of his determinations,
       gives, however, +0·8 thousand heat units. Similar contradictory
       results are often met with in thermochemistry owing to the
       imperfection of the existing methods, and particularly the
       necessity of depending on indirect methods for obtaining the
       fundamental figures. Thus Thomsen decomposed a dilute solution of
       potassium iodide by gaseous chlorine; the reaction gave +26·2,
       whence, having first determined the heat effects of the reactions
       KHO + HCl, KHO + HI and Cl + H in aqueous solutions, it was
       possible to find H + I + Aq; then, knowing HI + Aq, to find I + H.
       It is evident that unavoidable errors may accumulate.

  [69] One can believe, however, on the basis of Berthollet's doctrine,
       and the observations of Potilitzin (Note 66), that a certain slow
       decomposition of water by iodine takes place. On this view the
       observations of Dossios and Weith on the fact that the solubility
       of iodine in water increases after the lapse of several months
       will be comprehensible. Hydriodic acid is then formed, and it
       increases the solubility. If the iodine be extracted from such a
       solution by carbon bisulphide, then, as the authors showed, after
       the action of nitrous anhydride iodine may be again detected in
       the solution by means of starch. It can easily be understood that
       a number of similar reactions, requiring much time and taking
       place in small quantities, have up to now eluded the attention of
       investigators, who even still doubt the universal application of
       Berthollet's doctrine, or only see the thermochemical side of
       reactions, or else neglect to pay attention to the element of time
       and the influence of mass.

  [70] On the basis of the data in Note 68.

In consequence of what has been said above, it follows that _hydrobromic
and hydriodic acids_, being substances which are but slightly stable,
cannot be evolved in a gaseous state under many of those conditions under
which hydrochloric acid is formed. Thus if sulphuric acid in solution
acts on sodium iodide, all the same phenomena take place as with sodium
chloride (a portion of the sodium iodide gives hydriodic acid, and all
remains in solution), but if sodium iodide be mixed with strong sulphuric
acid, then the oxygen of the latter decomposes the hydriodic acid set
free, with liberation of iodine, H_{2}SO_{4} + 2HI = 2H_{2}O + SO_{2} +
I_{2}. This reaction takes place in the reverse direction in the presence
of a _large quantity_ of water (2,000 parts of water per 1 part of
SO_{2}), in which case not only the affinity of hydriodic acid for water
is brought to light but also the action of water in directing chemical
reactions in which it participates.[71] Therefore, with a halogen salt,
it is easy to obtain gaseous hydrochloric acid by the action of sulphuric
acid, but neither hydrobromic nor hydriodic acid can be so obtained in
the free state (as gases).[72] Other methods have to be resorted to for
their preparation, and recourse must not be had to compounds of oxygen,
which are so easily able to destroy these acids. Therefore hydrogen
sulphide, phosphorus, &c., which themselves easily take up oxygen, are
introduced as means for the conversion of bromine and iodine into
hydrobromic and hydriodic acids in the presence of water. For example, in
the action of phosphorus the essence of the matter is that the oxygen of
the water goes to the phosphorus, and the union of the remaining
elements leads to the formation of hydrobromic or hydriodic acid; but the
matter is complicated by the reversibility of the reaction, the affinity
for water, and other circumstances which are understood by following
Berthollet's doctrine. Chlorine (and bromine also) directly decomposes
hydrogen sulphide, forming hydrochloric acid and liberating sulphur, both
in a gaseous form and in solutions, whilst iodine only decomposes
hydrogen sulphide in weak solutions, when its affinity for hydrogen is
aided by the affinity of hydrogen iodide for water. In a gaseous state
iodine does not act on hydrogen sulphide,[73] whilst sulphur is able to
decompose gaseous hydriodic acid, forming hydrogen sulphide and a
compound of sulphur and iodine which with water forms hydriodic acid.[74]

  [71] A number of similar cases confirm what has been said in Chapter X.

  [72] This is prevented by the reducibility of sulphuric acid. If
       volatile acids be taken they pass over, together with the
       hydrobromic and hydriodic acids, when distilled; whilst many
       non-volatile acids which are not reduced by hydrobromic and
       hydriodic acids only act feebly (like phosphoric acid), or do not
       act at all (like boric acid).

  [73] This is in agreement with the thermochemical data, because if all
       the substances be taken in the gaseous state (for sulphur the heat
       of fusion is 0·3, and the heat of vaporisation 2·3) we have H_{2}
       + S = 4·7; H_{2} + Cl_{2} = 44; H_{2} + Br_{2} = 24, and H_{2} +
       I_{2} = -3 thousand heat units; hence the formation of H_{2}S
       gives less heat than that of HCl and HBr, but more than that of
       HI. In dilute solutions H_{2} + S + Aq = 9·3, and consequently
       less than the formation of all the halogen acids, as H_{2}S
       evolves but little heat with water, and therefore in dilute
       solutions chlorine, bromine, and iodine decompose hydrogen
       sulphide.

  [74] Here there are three elements, hydrogen, sulphur, and iodine, each
       pair of which is able to form a compound, HI, H_{2}S, and SI,
       besides which the latter may unite in various proportions. The
       complexity of chemical mechanics is seen in such examples as
       these. It is evident that only the study of the simplest cases can
       give the key to the more complex problems, and on the other hand
       it is evident from the examples cited in the last pages that,
       without penetrating into the conditions of chemical equilibria, it
       would be impossible to explain chemical phenomena. By following
       the footsteps of Berthollet the possibility of unravelling the
       problems will be reached; but work in this direction has only been
       begun during the last ten years, and much remains to be done in
       collecting experimental material, for which occasions present
       themselves at every step. In speaking of the halogens I wished to
       turn the reader's attention to problems of this kind.

If hydrogen sulphide be passed through water containing iodine, the
reaction H_{2}S + I_{2} = 2HI + S proceeds so long as the solution is
dilute, but when the mass of free HI increases the reaction stops,
because the iodine then passes into solution. A solution having a
composition approximating to 2HI + 4I_{2} + 9H_{2}O (according to Bineau)
does not react with H_{2}S, notwithstanding the quantity of free iodine.
Therefore only weak solutions of hydriodic acid can be obtained by
passing hydrogen sulphide into water with iodine.[74 bis]

  [74 bis] The same essentially takes place when sulphurous anhydride,
       in a dilute solution, gives hydriodic acid and sulphuric acid with
       iodine. On concentration a reverse reaction takes place. The
       equilibrated systems and the part played by water are everywhere
       distinctly seen.

To obtain[75] gaseous hydrobromic and hydriodic acids it is most
convenient to take advantage of the reactions between phosphorus, the
halogens, and water, the latter being present in small quantity
(otherwise the halogen acids formed are dissolved by it); the halogen is
gradually added to the phosphorus moistened with water. Thus if red
phosphorus be placed in a flask and moistened with water, and bromine be
added drop by drop (from a tap funnel), hydrobromic acid is abundantly
and uniformly disengaged.[76] Hydrogen iodide is prepared by adding 1
part of common (yellow) dry phosphorus to 10 parts of dry iodine in a
glass flask. On shaking the flask, union proceeds quietly between them
(light and heat being evolved), and when the mass of iodide of phosphorus
which is formed has cooled, water is added drop by drop (from a tap
funnel) and hydrogen iodide is evolved directly without the aid of heat.
These methods of preparation will be at once understood when it is
remembered (p. 468) that phosphorus chloride gives hydrogen chloride with
water. It is exactly the same here--the oxygen of the water passes over
to the phosphorus, and the hydrogen to the iodine, thus, PI_{3} +
3H_{2}O = PH_{3}O_{3} + 3HI.[77]

  [75] Methods of formation and preparation are nothing more than
       particular cases of chemical reaction. If the knowledge of
       chemical mechanics were more exact and complete than it now is it
       would be possible to foretell all cases of preparation _with every
       detail_ (of the quantity of water, temperature, pressure, mass,
       &c.) The study of practical methods of preparation is therefore
       one of the paths for the study of chemical mechanics. The reaction
       of iodine on phosphorus and water is a case like that mentioned in
       Note 74, and the matter is here further complicated by the
       possibility of the formation of the compound PH_{3} with HI, as
       well as the production of PI_{2}, PI_{3}, and the affinity of
       hydriodic acid and the acids of phosphorus for water. The
       theoretical interest of equilibria in all their complexity is
       naturally very great, but it falls into the background in presence
       of the primary interest of discovering practical methods for the
       isolation of substances, and the means of employing them for the
       requirements of man. It is only after the satisfaction of these
       requirements that interests of the other order arise, which in
       their turn must exert an influence on the former. For these
       reasons, whilst considering it opportune to point out the
       theoretical interest of chemical equilibria, the chief attention
       of the reader is directed in this work to questions of practical
       importance.

  [76] Hydrobromic acid is also obtained by the action of bromine on
       paraffin heated to 180°. Gustavson proposed to prepare it by the
       action of bromine (best added in drops together with traces of
       aluminium bromide) on anthracene (a solid hydrocarbon from coal
       tar). Balard prepared it by passing bromine vapour over moist
       pieces of common phosphorus. The liquid tribromide of phosphorus,
       directly obtained from phosphorus and bromine, also gives
       hydrobromic acid when treated with water. Bromide of potassium or
       sodium, when treated with sulphuric acid in the presence of
       phosphorus, also gives hydrobromic acid, but hydriodic acid is
       decomposed by this method. In order to free hydrobromic acid from
       bromine vapour it is passed over moist phosphorus and dried either
       by phosphoric anhydride or calcium bromide (calcium chloride
       cannot be used, as hydrochloric acid would be formed). Neither
       hydrobromic nor hydriodic acids can be collected over mercury, on
       which they act, but they may be directly collected in a dry vessel
       by leading the gas-conducting tube to the bottom of the vessel,
       both gases being much heavier than air. Merz and Holtzmann (1889)
       proposed to prepare HBr directly from bromine and hydrogen. For
       this purpose pure dry hydrogen is passed through a flask
       containing boiling bromine. The mixture of gas and vapour then
       passes through a tube provided with one or two bulbs, which is
       heated moderately in the middle. Hydrobromic acid is formed with a
       series of flashes at the part heated. The resultant HBr, together
       with traces of bromine, passes into a Woulfe's bottle into which
       hydrogen is also introduced, and the mixture is then carried
       through another heated tube, after which it is passed through
       water which dissolves the hydrobromic acid. According to the
       method proposed by Newth (1892) a mixture of bromine and hydrogen
       is led through a tube containing a platinum spiral, which is
       heated to redness after the air has been displaced from the tube.
       If the vessel containing the bromine be kept at 60°, the hydrogen
       takes up almost the theoretical amount of bromine required for the
       formation of HBr. Although the flame which appears in the
       neighbourhood of the platinum spiral does not penetrate into the
       vessel containing the bromine, still, for safety, a tube filled
       with cotton wool may be interposed.

       Hydroiodic acid is obtained in the same manner as hydrobromic. The
       iodine is heated in a small flask, and its vapour is carried over
       by hydrogen into a strongly heated tube, The gas passing from the
       tube is found to contain a considerable amount of HI, together
       with some free iodine. At a low red heat about 17 p.c. of the
       iodine vapour enters into combination; at a higher temperature, 78
       p.c. to 79 p.c.; and at a strong heat about 82 p.c.

  [77] But generally more phosphorus is taken than is required for the
       formation of PI_{3}, because otherwise a portion of the iodine
       distils over. If less than one-tenth part of iodine be taken, much
       phosphonium iodide, PH_{4}I, is formed. This proportion was
       established by Gay-Lussac and Kolbe. Hydriodic acid is also
       prepared in many other ways. Bannoff dissolves two parts of iodine
       in one part of a previously prepared strong (sp. gr. 1·67)
       solution of hydriodic acid, and pours it on to red phosphorus in a
       retort. Personne takes a mixture of fifteen parts of water, ten of
       iodine, and one of red phosphorus, which, when heated, disengages
       hydriodic acid mixed with iodine vapour; the latter is removed by
       passing it over moist phosphorus (Note 76). It must be remembered
       however that reverse reaction (Oppenheim) may take place between
       the hydriodic acid and phosphorus, in which the compounds PH_{4}I
       and PI_{2} are formed.

       It should be observed that the reaction between phosphorus, iodine
       and water must be carried out in the above proportions and with
       caution, as they may react with explosion. With red phosphorus the
       reaction proceeds quietly, but nevertheless requires care.

       L. Meyer showed that with an excess of iodine the reaction
       proceeds without the formation of bye-products (PH_{4}I),
       according to the equation P + 5I + 4H_{2}O = PH_{3}O_{4} + 5HI.
       For this purpose 100 grams of iodine and 10 grams of water are
       placed in a retort, and a paste of 5 grams of red phosphorus and
       10 grams of water is added little by little (at first with great
       care). The hydriodic acid may be obtained free from iodine by
       directing the neck of the retort upwards and causing the gas to
       pass through a shallow layer of water (respecting the formation of
       HI, _see_ also Note 75).

In a gaseous form hydrobromic and hydriodic acids are closely analogous
to hydrochloric acid; they are liquefied by pressure and cold, they fume
in the air, form solutions and hydrates, of constant boiling point, and
react on metals, oxides and salts, &c.[78] Only the relatively easy
decomposability of hydrobromic acid and especially of hydriodic acid,
clearly distinguish these acids from hydrochloric acid. For this reason,
hydriodic acid acts in a number of cases as a deoxidiser or reducer, and
frequently even serves as a means for the transference of hydrogen. Thus
Berthelot, Baeyer, Wreden, and others, by heating unsaturated
hydrocarbons in a solution of hydriodic acid, obtained their compounds
with hydrogen nearer to the limit C_{_n_}H_{2_n_ + 2} or even the
saturated compounds. For example, benzene, C_{6}H_{6}, when heated in a
closed tube with a strong solution of hydriodic acid, gives hexylene,
C_{6}H_{12}. The easy decomposability of hydriodic acid accounts for the
fact that iodine does not act by metalepsis on hydrocarbons, for the
hydrogen iodide liberated with the product of metalepsis, RI, formed,
gives iodine and the hydrogen compound, RH, back again. And therefore, to
obtain the products of iodine substitution, either iodic acid, HIO_{3}
(Kekulé), or mercury oxide, HgO (Weselsky), is added, as they immediately
react on the hydrogen iodide, thus: HIO_{3} + 5HI = 3H_{2}O + 3I_{2}, or,
HgO + 2HI = HgI_{2} + H_{2}O. From these considerations it will be
readily understood that iodine acts like chlorine (or bromine) on ammonia
and sodium hydroxide, for in these cases the hydriodic acid produced
forms NH_{4}I and NaI. With tincture of iodine or even the solid element,
a solution of ammonia immediately forms a highly-explosive solid black
product of metalepsis, NHI_{2}, generally known as _iodide_ of
_nitrogen_, although it still contains hydrogen (this was proved beyond
doubt by Szuhay 1893), which may be replaced by silver (with the
formation of NAgI_{2}): 3NH_{3} + 2I_{2} = 2NH_{4}I + NHI_{2}. However,
the composition of the last product is variable, and with an excess of
water NI_{3} seems to be formed. Iodide of nitrogen is just as explosive
as nitrogen chloride.[78 bis] In the action of iodine on sodium
hydroxide no bleaching compound is formed (whilst bromine gives one), but
a direct reaction is always accomplished with the formation of an iodate,
6NaHO + 3I_{2} = 5NaI + 3H_{2}O + NaIO_{3} (Gay-Lussac). Solutions of
other alkalis, and even a mixture of water and oxide of mercury, act in
the same manner.[79] This direct formation of _iodic acid_, HIO_{3} =
IO_{2}(OH), shows the propensity of iodine to give compounds of the type
IX_{5}. Indeed, this capacity of iodine to form compounds of a high type
emphasises itself in many ways. But it is most important to turn
attention to the fact that iodic acid is easily and directly formed by
the action of oxidising substances on iodine. Thus, for instance, strong
nitric acid directly converts iodine into iodic acid, whilst it has no
oxidising action on chlorine.[79 bis] This shows a greater affinity in
iodine for oxygen than in chlorine, and this conclusion is confirmed by
the fact that iodine displaces chlorine from its oxygen acids,[80] and
that in the presence of water chlorine oxidises iodine.[81] Even ozone or
a silent discharge passed through a mixture of oxygen and iodine vapour
is able to directly oxidise iodine[82] into iodic acid. It is disengaged
from solutions as a hydrate, HIO_{3}, which loses water at 170°, and
gives an anhydride, I_{2}O_{5}. Both these substances are crystalline
(sp. gr. I_{2}O_{5} 5·037, HIO_{3} 4·869 at 0°), colourless and soluble
in water;[83] both decompose at a red heat into iodine and oxygen, are in
many cases powerfully oxidising--for instance, they oxidise sulphurous
anhydride, hydrogen sulphide, carbonic oxide, &c.--form chloride of
iodine and water with hydrochloric acid, and with bases form salts, not
only normal MIO_{3}, but also acid; for example, KIO_{3}HIO_{3},
KIO_{3}2HIO_{3}.[83 bis] With hydriodic acid iodic acid immediately
reacts, disengaging iodine, HIO_{3} + 5HI = 3H_{2}O + 3I_{2}.

  [78] The specific gravities of their solutions as deduced by me on the
       basis of Topsöe and Berthelot's determinations for 15°/4° are as
       follows:--

                   10      20      30      40      50   60 p.c.
           HBr   1·071   1·156   1·258   1·374   1·505   1·650
           HI    1·075   1·164   1·267   1·399   1·567   1·769

       Hydrobromic acid forms two hydrates, HBr,2H_{2}O and HBr,H_{2}O,
       which have been studied by Roozeboom with as much completeness as
       the hydrate of hydrochloric acid (Chapter X. Note 37).

       With metallic silver, solutions of hydriodic acid give hydrogen
       with great ease, forming silver iodide. Mercury, lead, and other
       metals act in a similar manner.

  [78 bis] Iodide of nitrogen, NHI_{2} is obtained as a brown pulverulent
       precipitate on adding a solution of iodine (in alcohol, for
       instance) to a solution of ammonia. If it be collected on a
       filter-paper, it does not decompose so long as the precipitate is
       moist; but when dry it explodes violently, so that it can only be
       experimented upon in small quantities. Usually the filter-paper is
       torn into bits while moist, and the pieces laid upon a brick; on
       drying an explosion proceeds not only from friction or a blow, but
       even spontaneously. The more dilute the solution of ammonia, the
       greater is the amount of iodine required for the formation of the
       precipitate of NHI_{2}. A low temperature facilitates its
       formation. NHI_{2} dissolves in ammonia water, and when heated the
       solution forms HIO_{3} and iodine. With KI, iodide of nitrogen
       gives iodine, NH_{3} and KHO. These reactions (Selivanoff) are
       explained by the formation of HIO from NHI_{2} + 2H_{2}O = NH_{3} +
       2HIO--and then KI + HIO = I_{2} + KHO. Selivanoff (_see_ Note 29)
       usually observed a temporary formation of hypoiodous acid, HIO, in
       the reaction of ammonia upon iodine, so that here the formation of
       NHI_{2} is preceded by that of HIO--_i.e._ first I_{2} + H_{2}O =
       HIO + HI, and then not only the HI combines with NH_{3}, but also
       2HIO + NH_{3} = NHI_{2} + 2H_{2}O. With dilute sulphuric acid
       iodide of nitrogen (like NCl_{3}) forms hypoiodous acid, but it
       immediately passes into iodic acid, as is expressed by the
       equation 5HIO = 2I_{2} + HIO_{3} + 2H_{2}O (first 3HIO = HIO_{3} +
       2HI, and then HI + HIO = I_{2} + H_{2}O). Moreover, Selivanoff
       found that iodide of nitrogen, NHI_{2}, dissolves in an excess of
       ammonia water, and that with potassium iodide the solution gives
       the reaction for hypoiodous acid (the evolution of iodine in an
       alkaline solution). This shows that HIO participates in the
       formation and decomposition of NHI_{2}, and therefore the
       condition of the iodine (its metaleptic position) in them is
       analogous, and differs from the condition of the halogens in the
       haloid-anhydrides (for instance, NO_{2}Cl). The latter are
       tolerably stable, while (the haloid being designated by X)
       NHX_{2}, NX_{3}, XOH, RXO (_see_ Chapter XIII. Note 43), &c., are
       unstable, easily decomposed with the evolution of heat, and, under
       the action of water, the haloid is easily replaced by hydrogen
       (Selivanoff), as would be expected in true products of metalepsis.

  [79] Hypoiodous acid, HIO, is not known, but organic compounds, RIO, of
       this type are known. To illustrate the peculiarities of their
       properties we will mention one of these compounds, namely,
       _iodosobenzol_, C_{6}H_{5}IO. This substance was obtained by
       Willgerodt (1892), and also by V. Meyer, Wachter, and Askenasy, by
       the action of caustic alkalis upon phenoldiiodochloride,
       C_{6}H_{5}ICl_{2} (according to the equation, C_{6}H_{5}ICl_{2} +
       2MOH = C_{6}H_{5}IO + 2MCl + H_{2}O). Iodosobenzol is an amorphous
       yellow substance, whose melting point could not be determined
       because it explodes at 210°, decomposing with the evolution of
       iodine vapour. This substance dissolves in hot water and alcohol,
       but is not soluble in the majority of other neutral organic
       solvents. If acids do not oxidise C_{6}H_{5}IO, they give saline
       compounds in which iodosobenzol appears as a basic oxide of a
       diatomic metal, C_{6}H_{5}I. Thus, for instance, when an acetic
       acid solution of iodosobenzol is treated with a solution of nitric
       acid, it gives large monoclinic crystals of a nitric acid salt
       having the composition C_{6}H_{5}I(NO_{3})_{2} (like
       Ca(NO_{3})_{2}). In appearing as the analogue of basic oxides,
       iodosobenzol displaces iodine from potassium iodide (in a solution
       acidulated with acetic or hydrochloric acid)--_i.e._ it acts with
       its oxygen like HClO. The action of peroxide of hydrogen, chromic
       acid, and other similar oxidising agents gives iodoxybenzol,
       C_{6}H_{5}IO_{2}, which is a neutral substance--_i.e._ incapable
       of giving salts with acids (compare Chapter XIII. Note 43).

  [79 bis] The oxidation of iodine by strong nitric acid was discovered
       by Connell; Millon showed that it is effected, although more
       slowly, by the action of the hydrates of nitric acid up to
       HNO_{3},H_{2}O, but that the solution HNO_{3},2H_{2}O, and weaker
       solutions, do not oxidise, but simply dissolve, iodine. The
       participation of water in reactions is seen in this instance. It
       is also seen, for example, in the fact that dry ammonia combines
       directly with iodine--for instance, at 0° forming the compound
       I_{2},4NH_{3}--whilst iodide of nitrogen is only formed in
       presence of water.

  [80] Bromine also displaces chlorine--for instance, from chloric acid,
       directly forming bromic acid. If a solution of potassium chlorate
       be taken (75 parts per 400 parts of water), and iodine be added to
       it (80 parts), and then a small quantity of nitric acid, chlorine
       is disengaged on boiling, and potassium iodate is formed in the
       solution. In this instance the nitric acid first evolves a certain
       portion of the chloric acid, and the latter, with the iodine,
       evolves chlorine. The iodic acid thus formed acts on a further
       quantity of the potassium chlorate, sets a portion of the chloric
       acid free, and in this manner the action is kept up. Potilitzin
       (1887) remarked, however, that not only do bromine and iodine
       displace the chlorine from chloric acid and potassium chlorate,
       but also chlorine displaces bromine from sodium bromate, and,
       furthermore, the reaction does not proceed as a direct
       substitution of the halogens, but is accompanied by the formation
       of free acids; for example, 5NaClO_{3} + 3Br_{2} + 3H_{2}O = 5NaBr
       + 5HClO_{3} + HBrO_{3}.

  [81] If iodine be stirred up in water, and chlorine passed through the
       mixture, the iodine is dissolved; the liquid becomes colourless,
       and contains, according to the relative amounts of water and
       chlorine, either IHCl_{2}, or ICl_{3}, or HIO_{3}. If there be a
       small amount of water, then the iodic acid may separate out
       directly as crystals, but a complete conversion (Bornemann) only
       occurs when not less than ten parts of water are taken to one part
       of iodine--ICl + 3H_{2}O + 2Cl_{2} = IHO_{3} + 5HCl.

  [82] Schönbein and Ogier proved this. Ogier found that at 45° ozone
       immediately oxidises iodine vapour, forming first of all the oxide
       I_{2}O_{3}, which is decomposed by water or on heating into iodic
       anhydride and iodine. Iodic acid is formed at the positive pole
       when a solution of hydriodic acid is decomposed by a galvanic
       current (Riche). It is also formed in the combustion of hydrogen
       mixed with a small quantity of hydriodic acid (Salet).

  [83] Kämmerer showed that a solution of sp. gr. 2·127 at 14°,
       containing 2HIO_{3},9H_{2}O, solidified completely in the cold. On
       comparing solutions HI + _m_H_{2}O with HIO_{3} + _m_H_{2}O, we
       find that the specific gravity increases but the volume decreases,
       whilst in the passage of solutions HCl + _m_H_{2}O to HClO_{3} +
       _m_H_{2}O both the specific gravity and the volume increase, which
       is also observed in certain other cases (for example, H_{3}PO_{3}
       and H_{3}PO_{4}).

  [83 bis] Ditte (1890) obtained many iodates of great variety. A neutral
       salt, 2(LiIO_{3})H_{2}O, is obtained by saturating a solution of
       lithia with iodic acid. There is an analogous ammonium salt,
       2(NH_{4}IO_{3})H_{2}O. He also obtained hydrates of a more complex
       composition, such as 6(NH_{4}IO_{3})H_{2}O and
       6(NH_{4}IO_{3})2H_{2}O. Salts of the alkaline earths,
       Ba(IO_{3})_{2}H_{2}O and Sr(IO_{3})_{2}H_{2}O, may be obtained by
       a reaction of double decomposition from the normal salts of the
       type 2(MeIO_{3})H_{2}O. When evaporated at 70° to 80° with nitric
       acid these salts lose water. A mixture of solutions of nitrate of
       zinc and an alkaline iodate precipitates Zn(IO_{3})_{2}2H_{2}O. An
       anhydrous salt is thrown out if nitric acid be added to the
       solutions. Analogous salts of cadmium, silver, and copper give
       compounds of the type 2Me´IO_{3}4NH_{3} and
       Me´´(IO_{3})_{2}4NH_{3}, with gaseous ammonia (Me´ and Me´´ being
       elements of the first (Ag) and second (Cd, Zn, Cu) groups). With
       an aqueous solution of ammonia the above salts give substances of
       a different composition, such as Zn(IO_{3})_{2}(NH_{4})_{2}O,
       Cd(IO_{3})_{2}(NH_{4})_{2}O. Copper gives
       Cu(IO_{3})_{2}4(NH_{4})_{2}O and Cu(IO_{3})_{2}(NH_{4})_{2}O.
       These salts may be regarded as compounds of I_{2}O_{5}, and MeO
       and (NH_{4})_{2}O; for example, Zn(IO_{3})_{2}(NH_{4})_{2}O may be
       regarded as ZnO(NH_{4})_{2}OI_{2}O_{5}, or, as derived from the
       hydrate, I_{2}O_{5}2H_{2}O = 2(HIO_{3})H_{2}O.

As with chlorine, so with iodine, a _periodic acid_, HIO_{4}, is formed.
This acid is produced in the form of its salts, by the action of chlorine
on alkaline solutions of iodates, and also by the action of iodine on
chloric acid.[84] It crystallises from solutions as a hydrate containing
2H_{2}O (corresponding with HClO_{4},2H_{2}O), but as it forms salts
containing up to 5 atoms of metals, this water must be counted as water
of constitution. Therefore IO(OH)_{5} = HIO_{4},2H_{2}O corresponds with
the highest form of halogen compounds, IX_{7}.[85] In decomposing (at
200°) or acting as an oxidiser, periodic acid first gives iodic acid, but
it may also be ultimately decomposed.

  [84] If sodium iodate be mixed with a solution of sodium hydroxide,
       heated, and chlorine passed through the solution, a sparingly
       soluble salt separates out, which corresponds with periodic acid,
       and has the composition Na_{4}I_{2}O_{9},3H_{2}O.

        6NaHO + 2NaIO_{3} + 4Cl = 4NaCl + Na_{4}I_{2}O_{9} + 3H_{2}O.

       This compound is sparingly soluble in water, but dissolves easily
       in a very dilute solution of nitric acid. If silver nitrate be
       added to this solution a precipitate is formed which contains the
       corresponding compound of silver, Ag_{4}I_{2}O_{9},3H_{2}O. If
       this sparingly soluble silver compound be dissolved in hot nitric
       acid, orange crystals of a salt having the composition AgIO_{4}
       separate on evaporation. This salt is formed from the preceding by
       the nitric acid taking up silver oxide--Ag_{4}I_{2}O_{9} +
       2HNO_{3} = 2AgNO_3 + 2AgIO_{4} + H_{2}O. The silver salt is
       decomposed by water, with the re-formation of the preceding salt,
       whilst iodic acid remains in solution--

              4AgIO_{4} + H_{2}O = Ag_{4}I_{2}O_{9} + 2HIO_{4}.

       The structure of the first of these salts,
       Na_{4}I_{2}O_{9},3H_{2}O, presents itself in a simpler form if the
       water of crystallisation is regarded as an integral portion of the
       salt; the formula is then divided in two, and takes the form of
       IO(OH)_{3}(ONa)_{2}--that is, it answers to the type IOX_{5}, or
       IX_{7}, like AgIO_{4} which is IO_{3}(OAg). The composition of all
       the salts of periodic acids are expressed by this type IX_{7}.
       Kimmins (1889) refers all the salts of periodic acid to four
       types--the meta-salts of HIO_{4} (salts of Ag, Cu, Pb), the
       meso-salts of H_{3}IO_{5} (PbH, Ag_{2}H, CdH), the para-salts of
       H_{5}IO_{6} (Na_{2}H_{3}, Na_{3}H_{2}), and the di-salts of
       H_{4}I_{2}O_{9} (K_{4}, Ag_{4}, Ni_{2}). The three first are
       direct compounds of the type IX_{7}, namely, IO_{3}(OH),
       IO_{2}(OH)_{3}, and IO(OH)_{5}, and the last are types of
       diperiodic salts, which correspond with the type of the
       meso-salts, as pyrophosphoric salts correspond with
       orthophosphoric salts--_i.e._ 2H_{3}IO_{5}-H_{2}O =
       H_{4}I_{2}O_{9}.

  [85] Periodic acid, discovered by Magnus and Ammermüller, and whose
       salts were afterwards studied by Langlois, Rammelsberg, and many
       others, presents an example of hydrates in which it is evident
       that there is not that distinction between the water of hydration
       and of crystallisation which was at first considered to be so
       clear. In HClO,2H_{2}O the water, 2H_{2}O, is not displaced by
       bases, and must be regarded as water of crystallisation, whilst in
       HIO_{4},2H_{2}O it must be regarded as water of hydration. We
       shall afterwards see that the system of the elements obliges us to
       consider the halogens as substances giving a highest saline type,
       _GX__{7}, where _G_ signifies a halogen, and _X_ oxygen (O =
       _X__{2}), OH, and other like elements. The hydrate IO(OH)_{5}
       corresponding with many of the salts of periodic acid (for
       example, the salts of barium, strontium, mercury) does not exhaust
       all the possible forms. It is evident that various other pyro-,
       meta-, &c., forms are possible by the loss of water, as will be
       more fully explained in speaking of phosphoric acid, and as was
       pointed out in the preceding note.

Compounds formed between chlorine and iodine must be classed among the
most interesting halogen bodies.[86] These elements combine together
directly with evolution of heat, and form _iodine monochloride_, ICl, or
_iodine trichloride_, ICl_{3}.[87] As water reacts on these substances,
forming iodic acid and iodine, they have to be prepared from dry iodine
and chlorine.[88] Both substances are formed in a number of reactions;
for example, by the action of aqua regia on iodine, of chlorine on
hydriodic acid, of hydrochloric acid on periodic acid, of iodine on
potassium chlorate (with the aid of heat, &c.) Trapp obtained iodine
monochloride, in beautiful red crystals, by passing a rapid current of
chlorine into molten iodine. The monochloride then distils over and
solidifies, melting at 27°. By passing chlorine over the crystals of the
monochloride, it is easy to obtain iodine trichloride in orange crystals,
which melt at 34° and volatilise at 47°, but in so doing decompose (into
Cl_{2} and ClI). The chemical properties of these chlorides entirely
resemble those of chlorine and iodine, as would be expected, because, in
this instance, a combination of similar substances has taken place as in
the formation of solutions or alloys. Thus, for instance, the unsaturated
hydrocarbons (for example, C_{2}H_{4}), which are capable of directly
combining with chlorine and iodine, also directly combine with iodine
monochloride.

  [86] With respect to hydrogen, oxygen, chlorine, and other elements,
       bromine occupies an intermediate position between chlorine and
       iodine, and therefore there is no particular need for considering
       at length the compounds of bromine. This is the great advantage of
       a natural grouping of the elements.

  [87] They were both obtained by Gay-Lussac and many others. Recent data
       respecting iodine monochloride, ICl, entirely confirm the numerous
       observations of Trapp (1854), and even confirm his statement as to
       the existence of two isomeric (liquid and crystalline) forms
       (Stortenbeker). With a small excess of iodine, iodine monochloride
       remains liquid, but in the presence of traces of iodine
       trichloride it easily crystallises. Tanatar (1893) showed that of
       the two modifications of ICl, one is stable, and melts at 27°;
       while the other, which easily passes into the first, and is formed
       in the absence of ICl_{3}, melts at 14°. Schützenberger amplified
       the data concerning the action of water on the chlorides (Note
       88), and Christomanos gave the fullest data regarding the
       trichloride.

       After being kept for some time, the liquid monochloride of iodine
       yields red deliquescent octahedra, having the composition ICl_{4},
       which are therefore formed from the monochloride with the
       liberation of free iodine, which dissolves in the remaining
       quantity of the monochloride. This substance, however, judging by
       certain observations, is impure iodine trichloride. If 1 part of
       iodine be stirred up in 20 parts of water, and chlorine be passed
       through the liquid, then all the iodine is dissolved, and a
       colourless liquid is ultimately obtained which contains a certain
       proportion of chlorine, because this compound gives a metallic
       chloride and iodate with alkalis without evolving any free iodine:
       ICl_{5} + 6KHO = 5KCl + KIO_{3} + 3H_{2}O. The existence of a
       pentachloride ICl_{5} is, however, denied, because this substance
       has not been obtained in a free state.

       Stortenbeker (1888) investigated the equilibrium of the system
       containing the molecules I_{2}, ICl, ICl_{3}, and Cl_{2}, in the
       same way that Roozeboom (Chapter X. Note 38) examined the
       equilibrium of the molecules HCl, HCl,2H_{2}O, and H_{2}O. He
       found that iodine monochloride appears in two states, one (the
       ordinary) is stable and melts at 27°·2, whilst the other is
       obtained by rapid cooling, and melts at 13°·9, and easily passes
       into the first form. Iodine trichloride melts at 101° only in a
       closed tube under a pressure of 16 atmospheres.

  [88] By the action of water on iodine monochloride and trichloride a
       compound IHCl_{2} is obtained, which does not seem to be altered
       by water. Besides this compound, iodine and iodic acid are always
       formed, 10ICl + 3H_{2}O = HIO_{3} + 5IHCl_{2} + 2I_{2}; and in
       this respect iodine trichloride may be regarded as a mixture, ICl
       + ICl_{5} = 2ICl_{3}, but ICl_{5} + 3H_{2}O = IHO_{3} + 5HCl;
       hence iodic acid, iodine, the compound IHCl_{2}, and hydrochloric
       acid are also formed by the action of water.




                              CHAPTERR XII

                                 SODIUM


The neutral salt, sodium sulphate, Na_{2}SO_{4}, obtained when a mixture
of sulphuric acid and common salt is strongly heated (Chapter X.),[1]
forms a colourless saline mass consisting of fine crystals, soluble in
water. It is the product of many other double decompositions, sometimes
carried out on a large scale; for example, when ammonium sulphate is
heated with common salt, in which case the sal-ammoniac is volatilised,
&c. A similar decomposition also takes place when, for instance, a
mixture of lead sulphate and common salt is heated; this mixture easily
fuses, and if the temperature be further raised heavy vapours of lead
chloride appear. When the disengagement of these vapours ceases, the
remaining mass, on being treated with water, yields a solution of
sodium sulphate mixed with a solution of undecomposed common salt. A
considerable quantity, however, of the lead sulphate remains unchanged
during this reaction, PbSO_{4} + 2NaCl = PbCl_{2} + Na_{2}SO_{4}, the
vapours will contain lead chloride, and the residue will contain the
mixture of the three remaining salts. The cause and nature of the
reaction are just the same as were pointed out when considering the
action of sulphuric acid upon NaCl. Here too it may be shown that the
double decomposition is determined by the removal of PbCl_{2} from the
sphere of the action of the remaining substances. This is seen from the
fact that sodium sulphate, on being dissolved in water and mixed with a
solution of any lead salt (and even with a solution of lead chloride,
although this latter is but sparingly soluble in water), immediately
gives a white precipitate of lead sulphate. In this case the lead takes
up the elements of sulphuric acid from the sodium sulphate in the
solutions. On heating, the reverse phenomenon is observed. The reaction
in the solution depends upon the insolubility of the lead sulphate, and
the decomposition which takes place on heating is due to the volatility
of the lead chloride. Silver sulphate, Ag_{2}SO_{4}, in solution with
common salt, gives silver chloride, because the latter is insoluble in
water, Ag_{2}SO_{4} + 2NaCl = Na_{2}SO_{4} + 2AgCl. Sodium carbonate,
mixed in solution with the sulphates of iron, copper, manganese,
magnesium, &c., gives in solution sodium sulphate, and in the precipitate
a carbonate of the corresponding metal, because these salts of carbonic
acid are insoluble in water; for instance, MgSO_{4} + Na_{2}CO_{3} =
Na_{2}SO_{4} + MgCO_{3}. In precisely the same way sodium hydroxide acts
on solutions of the majority of the salts of sulphuric acid containing
metals, the hydroxides of which are insoluble in water--for instance,
CuSO_{4} + 2NaHO = Cu(HO)_{2} + Na_{2}SO_{4}. Sulphate of magnesium,
MgSO_{4}, on being mixed in solution with common salt, forms, although
not completely, chloride of magnesium, and sodium sulphate. On cooling
the mixture of such (concentrated) solutions sodium sulphate is
deposited, as was shown in Chapter X. This is made use of for preparing
it on the large scale in works where sea-water is treated. In this case,
on cooling, the reaction 2NaCl + MgSO_{4} = MgCl_{2} + Na_{2}SO_{4} takes
place.

  [1] Whilst describing in some detail the properties of sodium
      chloride, hydrochloric acid, and sodium sulphate, I wish to impart,
      by separate examples, an idea of the properties of saline
      substances, but the dimensions of this work and its purpose and aim
      do not permit of entering into particulars concerning every salt,
      acid, or other substance. The fundamental object of this work--an
      account of the characteristics of the elements and an acquaintance
      with the forces acting between atoms--has nothing to gain from the
      multiplication of the number of as yet ungeneralised properties and
      relations.

Thus where sulphates and salts of sodium are in contact, it may be
expected that sodium sulphate will be formed and separated if the
conditions are favourable; for this reason it is not surprising that
sodium sulphate is often found in the native state. Some of the springs
and salt lakes in the steppes beyond the Volga, and in the Caucasus,
contain a considerable quantity of sodium sulphate, and yield it by
simple evaporation of the solutions. Beds of this salt are also met with;
thus at a depth of only 5 feet, about 38 versts to the east of Tiflis, at
the foot of the range of the 'Wolf's mane' (Voltchia griva) mountains, a
deep stratum of very pure Glauber's salt, Na_{2}SO_{4},10H_{2}O, has been
found.[2] A layer two metres thick of the same salt lies at the bottom of
several lakes (an area of about 10 square kilometres) in the Kouban
district near Batalpaschinsk, and here its working has been commenced
(1887). In Spain, near Arangoulz and in many parts of the Western States
of North America, mineral sodium sulphate has likewise been found, and is
already being worked.

  [2] Anhydrous (ignited) sodium sulphate, Na_{2}SO_{4}, is known in
      trade as 'sulphate' or salt-cake, in mineralogy _thenardite_.
      Crystalline decahydrated salt is termed in mineralogy _mirabilite_,
      and in trade Glauber's salt. On fusing it, the monohydrate
      Na_{2}SO_{4}H_{2}O is obtained, together with a supersaturated
      solution.

The methods of obtaining salts by means of double decomposition from
others already prepared are so general, that in describing a given salt
there is no necessity to enumerate the cases hitherto observed of its
being formed through various double decompositions.[3] The possibility
of this occurrence ought to be foreseen according to Berthollet's
doctrine from the properties of the salt in question. On this account it
is important to know the properties of salts; all the more so because
up to the present time those very properties (solubility, formation
of crystallo-hydrates, volatility, &c.) which may be made use of for
separating them from other salts have not been generalised.[4] These
properties as yet remain subjects for investigation, and are rarely
to be foreseen. The crystallo-hydrate of the normal sodium sulphate,
Na_{2}SO_{4},10H_{2}O, very easily parts with water, and may be obtained
in an anhydrous state if it be carefully heated until the weight
remains constant; but if heated further, it partly loses the elements
of sulphuric anhydride. The normal salt fuses at 843° (red heat), and
volatilises to a slight extent when very strongly heated, in which
case it naturally decomposes with the evolution of SO_{3}. At 0° 100
parts of water dissolve 5 parts of the anhydrous salt, at 10° 9 parts,
at 20° 19·4, at 30° 40, and at 34° 55 parts, the same being the case
in the presence of an excess of crystals of Na_{2}SO_{4},10H_{2}O.[5]
At 34° the latter fuses, and the solubility decreases at higher
temperatures.[6] A concentrated solution at 34° has a composition
nearly approaching to Na_{2}SO_{4} + 14H_{2}O, and the decahydrated salt
contains 78·9 of the anhydrous salt combined with 100 parts of water.
From the above figures it is seen that the decahydrated salt cannot
fuse without decomposing,[7] like hydrate of chlorine, Cl_{2},8H_{2}O
(Chapter XI., Note 10). Not only the fused decahydrated salt, but also
the concentrated solution at 34° (not all at once, but gradually),
yields the monohydrated salt, Na_{2}SO_{4},H_{2}O. The heptahydrated
salt, Na_{2}SO_{4},7H_{2}O, also splits up, even at low temperatures,
with the formation of this monohydrated salt, and therefore from 35°
the solubility can be given only for the latter. For 100 parts of water
this is as follows: at 40° 48·8, at 50° 46·7, at 80° 43·7, at 100°
42·5 parts of the anhydrous salt. If the decahydrated salt be fused,
and the solution allowed to cool in the presence of the monohydrated
salt, then at 30° 50·4 parts of anhydrous salt are retained in the
solution, and at 20° 52·8 parts. Hence, with respect to the anhydrous
and monohydrated salts, the solubility is identical, and falls with
increasing temperature, whilst with respect to decahydrated salt, the
solubility rises with increasing temperature. So that if in contact
with a solution of sodium sulphate there are only crystals of that
heptahydrated salt (Chapter I., Note 54), Na_{2}SO_{4},7H_{2}O, which
is formed from saturated solutions, then saturation sets in when the
solution has the following composition per 100 parts of salt: at 0°
19·6, at 10° 30·5, at 20° 44·7, and at 25° 52·9 parts of anhydrous
salt. Above 27° the heptahydrated salt, like the decahydrated
salt at 34°, splits up into the monohydrated salt and a saturated
solution. Thus sodium sulphate has three curves of solubility: one for
Na_{2}SO_{4},7H_{2}O (from 0° to 26°), one for Na_{2}SO_{4},10H_{2}O
(from 0° to 34°), and one for Na_{2}SO_{4},H_{2}O (a descending curve
beginning at 26°), because there are three of these crystallo-hydrates,
and the solubility of a substance only depends upon the particular
condition of that portion of it which has separated from the solution or
is present in excess.[8]

  [3] The salts may be obtained not only by methods of substitution of
      various kinds, but also by many other combinations. Thus sodium
      sulphate may be formed from sodium oxide and sulphuric anhydride,
      by oxidising sodium sulphide, Na_{2}S, or sodium sulphite,
      Na_{2}SO_{3}, &c. When sodium chloride is heated in a mixture of
      the vapours of water, air, and sulphurous anhydride, sodium
      sulphate is formed. According to this method (patented by
      Hargreaves and Robinson), sodium sulphate, Na_{2}SO_{4}, is
      obtained from NaCl without the preliminary manufacture of
      H_{2}SO_{4}. Lumps of NaCl pressed into bricks are loosely packed
      into a cylinder and subjected, at a red heat, to the action of
      steam, air and SO_{2}. Under these conditions, HCl, sulphate, and a
      certain amount of unaltered NaCl are obtained. This mixture is
      converted into soda by Gossage's process (_see_ Note 15) and may
      have some practical value.

  [4] Many observations have been made, but little general information
      has been obtained from particular cases. In addition to which, the
      properties of a given salt are changed by the presence of other
      salts. This takes place not only in virtue of mutual decomposition
      or formation of double salts capable of separate existence, but is
      determined by the influence which some salts exert on others, or by
      forces similar to those which act during solution. Here nothing has
      been generalised to that extent which would render it possible to
      predict without previous investigation, if there be no close
      analogy to help us. Let us state one of these numerous cases: 100
      parts of water at 20° dissolve 34 parts of potassium nitrate but on
      the addition of sodium nitrate the solubility of potassium nitrate
      increases to 48 parts in 10 of water (Carnelley and Thomson). In
      general, in all cases of which there are accurate observations it
      appears that the presence of foreign salts changes the properties
      of any given salt.

  [5] The information concerning solubility (Chapter I.) is given
      according to the determinations of Gay-Lussac, Lovell, and Mulder.

  [6] In Chapter I., Note 24, we have already seen that with many other
      sulphates the solubility also decreases after a certain temperature
      is passed. Gypsum, CaSO_{4},2H_{2}O, lime, and many other compounds
      present such a phenomenon. An observation of Tilden's (1884) is
      most instructive; he showed that on raising the temperature (in
      closed vessels) above 140° the solubility of sodium sulphate again
      begins to increase. At 100° 100 parts of water dissolve about 43
      parts of anhydrous salt, at 140° 42 parts, at 160° 43 parts, at
      180° 44 parts, at 230° 46 parts. According to Étard (1892) the
      solubility of 30 parts of Na{2}SO_{4} in 100 of solution (or 43 per
      100 of water) corresponds to 80°, and above 240° the solubility
      again falls, and very rapidly, so that at 320° the solution
      contains 12 per 100 of solution (about 14 per 100 of water) and a
      further rise of temperature is followed by a further deposition of
      the salt. It is evident that the phenomenon of saturation,
      determined by the presence of an excess of the dissolved substance,
      is very complex, and therefore that for the theory of solutions
      considered as liquid indefinite chemical compounds, many useful
      statements can hardly be given.

  [7] Already referred to in Chapter I., Note 56.

      The example of sodium sulphate is historically very important for
      the theory of solutions. Notwithstanding the number of
      investigations which have been made, it is still insufficiently
      studied, especially from the point of the vapour tension of
      solutions and crystallo-hydrates, so that those processes cannot be
      applied to it which Guldberg, Roozeboom, Van't Hoff, and others
      applied to solutions and crystallo-hydrates. It would also be most
      important to investigate the influence of pressure on the various
      phenomena corresponding with the combinations of water and sodium
      sulphate, because when crystals are separated--for instance, of the
      decahydrated salt--an increase of volume takes place, as can be
      seen from the following data:--the sp. gr. of the anhydrous salt is
      2·66, that of the decahydrated salt = 1·46, but the sp. gr. of
      solutions at 15°/4° = 9,992 + 90·2_p_ + 0·35_p_^2 where p
      represents the percentage of anhydrous salt in the solution, and
      the sp. gr. of water at 4° = 10,000. Hence for solutions containing
      20 p.c. of anhydrous salt the sp. gr. = 1·1936; therefore the
      volume of 100 grams of this solution = 83·8 c.c., and the volume of
      anhydrous salt contained in it is equal to 20/2·66, or = 7·5 c.c.,
      and the volume of water = 80·1 c.c. Therefore, the solution, on
      decomposing into anhydrous salt and water, increases in volume
      (from 83·8 to 87·6); but in the same way 83·8 c.c. of 20 p.c.
      solution are formed from (45·4/1·46 =) 31·1 c.c. of the
      decahydrated salt, and 54·6 c.c. of water--that is to say, that
      during the formation of a solution from 85·7 c.c., 83·8 c.c. are
      formed.

  [8] From this example it is evident the solution remains unaltered
      until from the contact of a solid it becomes either saturated or
      supersaturated, crystallisation being determined by the attraction
      to a solid, as the phenomenon of supersaturation clearly
      demonstrates. This partially explains certain apparently
      contradictory determinations of solubility. The best investigated
      example of such complex relations is cited in Chapter XIV., Note 50
      (for CaCl_{2}).

Thus solutions of sodium sulphate may give crystallo-hydrates of three
kinds on cooling the saturated solution: the unstable heptahydrated salt
is obtained at temperatures below 26°, the decahydrated salt forms under
ordinary conditions at temperatures below 34°, and the monohydrated salt
at temperatures above 34°. Both the latter crystallo-hydrates present a
stable state of equilibrium, and the heptahydrated salt decomposes into
them, probably according to the equation 3Na_{2}SO_{4},7H_{2}0 =
2Na_{2}SO_{4},10H_{2}O + Na_{2}SO_{4},H_{2}O. The ordinary decahydrated
salt is called _Glauber's salt_. All forms of these crystallo-hydrates
lose their water entirely, and give the anhydrous salt when dried over
sulphuric acid.[9]

  [9] According to Pickering's experiments (1886), the molecular weight
      in grams (that is, 142 grams) of anhydrous sodium sulphate, on
      being dissolved in a large mass of water, at 0° absorbs (hence
      the-sign)-1,100 heat units, at 10°-700, at 15°-275, at 20° gives
      out +25, at 25° +300 calories. For the decahydrated salt,
      Na_{2}SO_{4},10H_{2}O, 5°-4,225, 10°-4,000, 15°-3,570, 20°-3,160,
      25°-2,775. Hence (just as in Chapter I., Note 56) the heat of the
      combination Na_{2}SO_{4},10H_{2}O at 5° = +3,125, 10° = +3,250, 20°
      = +3,200, and 25° = +3,050.

      It is evident that the decahydrated salt dissolving in water gives
      a decrease of temperature. Solutions in hydrochloric acid give a
      still greater decrease, because they contain the water of
      crystallisation in a solid state--that is, like ice--and this on
      melting absorbs heat. A mixture of 15 parts of
      Na_{2}SO_{4},10H_{2}O and 12 parts of strong hydrochloric acid
      produces sufficient cold to freeze water. During the treatment with
      hydrochloric acid a certain quantity of sodium chloride is formed.

Sodium sulphate, Na_{2}SO_{4}, only enters into a few reactions of
combination with other salts, and chiefly with salts of the same acid,
forming double sulphates. Thus, for example, if a solution of sodium
sulphate be mixed with a solution of aluminium, magnesium, or ferrous
sulphate, it gives crystals of a double salt when evaporated. Sulphuric
acid itself forms a compound with sodium sulphate, which is exactly
like these double salts. It is formed with great ease when sodium
sulphate is dissolved in sulphuric acid and the solution evaporated.
On evaporation, crystals of the acid salt separate, Na_{2}SO_{4} +
H_{2}SO_{4} = 2NaHSO_{4}. This separates from hot solutions, whilst the
crystallo-hydrate, NaHSO_{4},H_{2}O,[10] separates from cold solutions.
The crystals when exposed to damp air decompose into H_{2}SO_{4}, which
deliquesces, and Na_{2}SO_{4} (Graham, Rose); alcohol also extracts
sulphuric acid from the acid salt. This shows the feeble force which
holds the sulphuric acid to the sodium sulphate.[11] Both acid sodium
sulphate and all mixtures of the normal salt and sulphuric acid lose
water when heated, and are converted into sodium _pyrosulphate_,
Na_{2}S_{2}O_{7}, at a low red heat.[11 bis] This anhydrous salt, at a
bright red heat, parts with the elements of sulphuric anhydride, the
normal sodium sulphate remaining behind--Na_{2}S_{2}O_{7} = Na_{2}SO_{4}
+ SO_{3}. From this it is seen that the normal salt is able to combine
with water, with other sulphates, and with sulphuric anhydride or
acid, &c.

  [10] The very large and well-formed crystals of this salt resemble the
       hydrate H_{2}SO_{4},H_{2}O, or SO(OH)_{4}. In general the
       replacement of hydrogen by sodium modifies many of the properties
       of acids less than its replacement by other metals. This most
       probably depends on the volumes being nearly equal.

  [11] In solution (Berthelot) the acid salt in all probability
       decomposes most in the greatest mass of water. The specific
       gravity (according to the determinations of Marignac) of solutions
       at 15°/4° = 9,992 + 77·92_p_ + 0·231_p_^2 (_see_ Note 7). From
       these figures, and from the specific gravities of sulphuric acid,
       it is evident that on mixing solutions of this acid and sodium
       sulphate _expansion_ will always take place; for instance,
       H_{2}SO_{4} + 25H_{2}O with Na_{2}SO_{4} + 25H_{2}O increases from
       483 volumes to 486. In addition to which, in weak solutions heat
       is absorbed, as shown in Chapter X., Note 27. Nevertheless, even
       more acid salts may be formed and obtained in a crystalline form.
       For instance, on cooling a solution of 1 part of sodium sulphate
       in 7 parts of sulphuric acid, crystals of the composition
       NaHSO_{4},H_{2}SO_{4} are separated (Schultz, 1868). This compound
       fuses at about 100°; the ordinary acid salt, NaHSO_{4}, at 149°.

  [11 bis] On decreasing the pressure, sodium hydrogen sulphate,
       NaHSO_{4}, dissociates much more easily than at the ordinary
       pressure; it loses water and forms the pyrosulphate,
       Na_{2}S_{2}O_{7}; this reaction is utilised in chemical works.

Sodium sulphate may by double decomposition be converted into a sodium
salt of any other acid, by means of heat and taking advantage of the
volatility, or by means of solution and taking advantage of the different
degree of solubility of the different salts. Thus, for instance, owing to
the insolubility of barium sulphate, sodium hydroxide or caustic soda may
be prepared from sodium sulphate, if barium hydroxide be added to its
solution, Na_{2}SO_{4} + Ba(HO)_{2} = BaSO_{4} + 2NaHO. And by taking any
salt of barium, BaX_{2}, the corresponding salt of sodium may be
obtained, Na_{2}SO_{4} + BaX_{2} = BaSO_{4} + 2NaX. Barium sulphate thus
formed, being a very sparingly-soluble salt, is obtained as a
precipitate, whilst the sodium hydroxide, or salt, NaX, is obtained in
solution, because _all salts of sodium are soluble_. Berthollet's
doctrine permits all such cases to be foreseen.

The reactions of _decomposition_ of sodium sulphate are above all
noticeable by the separation of oxygen. Sodium sulphate by itself
is very stable, and it is only at a temperature sufficient to melt
iron that it is possible to separate the elements SO_{3} from it, and
then only partially. However, the oxygen may be separated from sodium
sulphate, as from all other sulphates, by means of many substances
which are able to combine with oxygen, such as charcoal and sulphur,
but hydrogen is not able to produce this action. If sodium sulphate be
heated with charcoal, then carbonic oxide and anhydride are evolved,
and there is produced, according to the circumstances, either the
lower oxygen compound, sodium sulphite, Na_{2}SO_{3} (for instance, in
the formation of glass); or else the decomposition proceeds further,
and sodium sulphide, Na_{2}S, is formed, according to the equation
Na_{2}SO_{4} + 2C = 2CO_{2} + Na_{2}S.

On the basis of this reaction the greater part of the sulphate of sodium
prepared at chemical works is converted into _soda ash_--that is,
_sodium carbonate_, Na_{2}CO_{3}, which is used for many purposes. In
the form of carbonates, the metallic oxides behave in many cases just
as they do in the state of oxides or hydroxides, owing to the feeble
acid properties of carbonic acid. However, the majority of the salts of
carbonic acid are insoluble, whilst sodium carbonate is one of the few
soluble salts of this acid, and therefore reacts with facility. Hence
sodium carbonate is employed for many purposes, in which its alkaline
properties come into play. Thus, even under the action of feeble organic
acids it immediately parts with its carbonic acid, and gives a sodium
salt of the acid taken. Its solutions exhibit an alkaline reaction on
litmus. It aids the passage of certain organic substances (tar, acids)
into solution, and is therefore used, like caustic alkalis and soap
(which latter also acts by virtue of the alkali it contains), for the
removal of certain organic substances, especially in bleaching cotton
and similar fabrics. Besides which a considerable quantity of sodium
carbonate is used for the preparation of sodium hydroxide or caustic
soda, which has also a very wide application. In large chemical works
where sodium carbonate is manufactured from Na_{2}SO_{4}, it is usual
first to manufacture sulphuric acid, and then by its aid to convert
common salt into sodium sulphate, and lastly to convert the sodium
sulphate thus obtained into carbonate and caustic soda. Hence these
works prepare both alkaline substances (soda ash and caustic soda) and
acid substances (sulphuric and hydrochloric acids), the two classes of
chemical products which are distinguished for the greatest energy of
their reactions and are therefore most frequently applied to technical
purposes. Factories manufacturing soda are generally called alkali works.

The process of the conversion of sodium sulphate into sodium carbonate
consists in strongly heating a mixture of the sulphate with charcoal
and calcium carbonate. The following reactions then take place: the
sodium sulphate is first deoxidised by the charcoal, forming sodium
sulphide and carbonic anhydride, Na_{2}SO_{4} + 2C = Na_{2}S + 2CO_{2}.
The sodium sulphide thus formed then enters into double decomposition
with the calcium carbonate taken, and gives calcium sulphide and sodium
carbonate, Na_{2}S + CaCO_{3} = Na_{2}CO_{3} + CaS.

[Illustration: FIG. 68.--Reverberatory furnace for the manufacture
of sodium carbonate. F, grate. A, bridge. M, hearth for the ultimate
calcination of the mixture of sodium sulphate, coal, and calcium
carbonate, which is charged from above into the part of the furnace
furthest removed from the fire F. P, P, doors for stirring and bringing
the mass towards the grate F by means of stirrers R. At the end of the
operation the semifused mass is charged into trucks C.]

Besides which, under the action of the heat, a portion of the excess
of calcium carbonate is decomposed into lime and carbonic anhydride,
CaCO_{3} = CaO + CO_{2}, and the carbonic anhydride with the excess
of charcoal forms carbon monoxide, which towards the end of the
operation shows itself by the appearance of a blue flame. Thus from a
mass containing sodium sulphate we obtain a mass which includes sodium
carbonate, calcium sulphide, and calcium oxide, but none of the sodium
sulphide which was formed on first heating the mixture. The entire
process, which proceeds at a high temperature, may be expressed by a
combination of the three above-mentioned formulæ, if it be considered
that the product contains one equivalent of calcium oxide to two
equivalents of calcium sulphide.[12] The sum of the reactions may then
be expressed thus: 2Na_{2}SO_{4} + 3CaCO_{3} + 9C = 2Na_{2}CO_{3} +
CaO,2CaS + 10CO. Indeed, the quantities in which the substances are
mixed together at chemical works approaches to the proportion required
by this equation. The entire process of decomposition is carried on
in reverberatory furnaces, into which a mixture of 1,000 parts of
sodium sulphate, 1,040 parts of calcium carbonate (as a somewhat porous
limestone), and 500 parts of small coal is introduced from above.
This mixture is first heated in the portion of the furnace which is
furthest removed from the fire-grate; it is then brought to the portion
nearest to the fire-grate, when it is stirred during heating. The
partially fused mass obtained at the end of the process is cooled, and
then subjected to methodical lixiviation[13] to extract the sodium
carbonate, the mixture of calcium oxide and sulphide forming the
so-called 'soda waste' or 'alkali waste.'[14]

  [12] Calcium sulphide, CaS, like many metallic sulphides which are
       soluble in water, is decomposed by it (Chapter X.), CaS + H_{2}O =
       CaO + H_{2}S, because hydrogen sulphide is a very feeble acid. If
       calcium sulphide be acted on by a large mass of water, lime may be
       precipitated, and a state of equilibrium will be reached, when the
       system CaO + 2CaS remains unchanged. Lime, being a product of the
       action of water on CaS, limits this action. Therefore, if in black
       ash the lime were not in excess, a part of the sulphide would be
       in solution (actually there is but very little). In this manner in
       the manufacture of sodium carbonate the conditions of equilibrium
       which enter into double decompositions have been made use of (_see
       above_), and the aim is to form directly the unchangeable product
       CaO,2CaS. This was first regarded as a special insoluble compound,
       but there is no evidence of its independent existence.

  [13] [Illustration: FIG. 69.--Apparatus for the methodical lixiviation
       of black ash, &c. Water flows into the tanks from the pipes _r_,
       _r_, and the saturated liquid is drawn off from _c_, _c_.]

       _Methodical lixiviation_ is the extraction, by means of water, of
       a soluble substance from the mass containing it. It is carried on
       so as not to obtain weak aqueous solutions, and in such a way that
       the residue shall not contain any of the soluble substance. This
       problem is practically of great importance in many industries. It
       is required to extract from the mass all that is soluble in water.
       This is easily effected if water be first poured on the mass, the
       strong solution thus obtained decanted, then water again poured
       on, time being allowed for it to act, then again decanted, and so
       on until fresh water does not take up anything. But then finally
       such weak solutions are obtained that it would be very
       disadvantageous to evaporate them. This is avoided by pouring the
       fresh hot water destined for the lixiviation, not onto the fresh
       mass, but upon a mass which has already been subjected to a first
       lixiviation by weak solutions. In this way the fresh water gives a
       weak solution. The strong solution which goes to the evaporating
       pan flows from those parts of the apparatus which contain the
       fresh, as yet unlixiviated, mass, and thus in the latter parts the
       weak alkali formed in the other parts of the apparatus becomes
       saturated as far as possible with the soluble substance. Generally
       several intercommunicating vessels are constructed (standing at
       the same level) into which in turn the fresh mass is charged which
       is intended for lixiviation; the water is poured in, the alkali
       drawn off, and the lixiviated residue removed. The illustration
       represents such an apparatus, consisting of four communicating
       vessels. The water poured into one of them flows through the two
       nearest and issues from the third. The fresh mass being placed in
       one of these boxes or vessels, the stream of water passing through
       the apparatus is directed in such a manner as to finally issue
       from this vessel containing the fresh unlixiviated mass. The fresh
       water is added to the vessel containing the material which has
       been almost completely exhausted. Passing through this vessel it
       is conveyed by the pipe (syphon passing from the bottom of the
       first box to the top of the second) communicating with the second;
       it finally passes (also through a syphon pipe) into the box (the
       third) containing the fresh material. The water will extract all
       that is soluble in the first vessel, leaving only an insoluble
       residue. This vessel is then ready to be emptied, and refilled
       with fresh material. The levels of the liquids in the various
       vessels will naturally be different, in consequence of the various
       strengths of the solutions which they contain.

       It must not, however, be thought that sodium carbonate alone
       passes into the solution; there is also a good deal of caustic
       soda with it, formed by the action of lime on the carbonate of
       sodium, and there are also certain sodium sulphur compounds with
       which we shall partly become acquainted hereafter. The sodium
       carbonate, therefore, is not obtained in a very pure state. The
       solution is concentrated by evaporation. This is conducted by
       means of the waste heat from the soda furnaces, together with that
       of the gases given off. The process in the soda furnaces can only
       be carried on at a high temperature, and therefore the smoke and
       gases issuing from them are necessarily very hot. If the heat they
       contain was not made use of there would be a great waste of fuel;
       consequently in immediate proximity to these furnaces there is
       generally a series of pans or evaporating boilers, under which the
       gases pass, and into which the alkali solution is poured. On
       evaporating the solution, first of all the undecomposed sodium
       sulphate separates, then the sodium carbonate or soda crystals.
       These crystals as they separate are raked out and placed on
       planks, where the liquid drains away from them. Caustic soda
       remains in the residue, and also any sodium chloride which was not
       decomposed in the foregoing process.

       Part of the sodium carbonate is recrystallised in order to purify
       it more thoroughly. In order to do this a saturated solution is
       left to crystallise at a temperature below 30° in a current of
       air, in order to promote the separation of the water vapour. The
       large transparent crystals (efflorescent in air) of
       Na_{2}CO_{3},10H_{2}O are then formed which have already been
       spoken of (Chapter I.).

  [14] The whole of the sulphur used in the production of the sulphuric
       acid employed in decomposing the common salt is contained in this
       residue. This is the great burden and expense of the soda works
       which use Leblanc's method. As an instructive example from a
       chemical point of view, it is worth while mentioning here two of
       the various methods of recovering the sulphur from the soda waste.
       Chance's process is treated in Chapter XX., Note 6.

       Kynaston (1885) treats the soda waste with a solution (sp. gr.
       l·21) of magnesium chloride, which disengages sulphuretted
       hydrogen: CaS + MgCl_{2} + 2H_{2}O = CaCl_{2} + Mg(OH)_{2} +
       H_{2}S. Sulphurous anhydride is passed through the residue in
       order to form the insoluble calcium sulphite: CaCl_{2} +
       Mg(OH)_{2} + SO_{2} = CaSO_{3} + MgCl_{2} + H_{2}O. The solution
       of magnesium chloride obtained is again used, and the washed
       calcium sulphite is brought into contact at a low temperature with
       hydrochloric acid (a weak aqueous solution) and hydrogen sulphide,
       the whole of the sulphur then separating:

          CaSO_{3} + 2H_{2}S + 2HCl = CaCl_{2} + 3H_{2}O + 3S.

       But most efforts have been directed towards avoiding the formation
       of soda waste.

The above-mentioned process for making soda was discovered in the year
1808 by the French doctor Leblanc, and is known as the Leblanc process.
The particulars of the discovery are somewhat remarkable. Sodium
carbonate, having a considerable application in industry, was for a
long time prepared exclusively from the ash of marine plants (Chapter
XI., page 497). Even up to the present time this process is carried on
in Normandy. In France, where for a long time the manufacture of large
quantities of soap (so-called Marseilles soap) and various fabrics
required a large amount of soda, the quantity prepared at the coast was
insufficient to meet the demand. For this reason during the wars at the
beginning of the century, when the import of foreign goods into France
was interdicted, the want of sodium carbonate was felt. The French
Academy offered a prize for the discovery of a profitable method of
preparing it from common salt. Leblanc then proposed the above-mentioned
process, which is remarkable for its great simplicity.[15]

  [15] Among the drawbacks of the Leblanc process are the accumulation of
       'soda waste' (Note 14) owing to the impossibility at the
       comparatively low price of sulphur (especially in the form of
       pyrites) of finding employment for the sulphur and sulphur
       compounds for which this waste is sometimes treated, and also the
       insufficient purity of the sodium carbonate for many purposes. The
       advantages of the Leblanc process, besides its simplicity and
       cheapness, are that almost the whole of the acids obtained as
       bye-products have a commercial value; for chlorine and bleaching
       powder are produced from the large amount of hydrochloric acid
       which appears as a bye-product; caustic soda also is very easily
       made, and the demand for it increases every year. In those places
       where salt, pyrites, charcoal, and limestone (the materials
       required for alkali works) are found side by side--as, for
       instance, in the Ural or Don districts--conditions are favourable
       to the development of the manufacture of sodium carbonate on an
       enormous scale; and where, as in the Caucasus, sodium sulphate
       occurs naturally, the conditions are still more favourable. A
       large amount, however, of the latter salt, even from soda works,
       is used in making glass. The most important soda works, as regards
       the quantity of products obtained from them, are the English
       works.

       As an example of the other numerous and varied methods of
       manufacturing soda from sodium chloride, the following may be
       mentioned: Sodium chloride is decomposed by oxide of lead, PbO,
       forming lead chloride and sodium oxide, which, with carbonic
       anhydride, yields sodium carbonate (Scheele's process). In Cornu's
       method sodium chloride is treated with lime, and then exposed to
       the air, when it yields a small quantity of sodium carbonate. In
       E. Kopp's process sodium sulphate (125 parts) is mixed with oxide
       of iron (80 parts) and charcoal (55 parts), and the mixture is
       heated in reverberatory furnaces. Here a compound,
       Na_{6}Fe_{4}S_{3}, is formed, which is insoluble in water absorbs
       oxygen and carbonic anhydride, and then forms sodium carbonate and
       ferrous sulphide; this when roasted gives sulphurous anhydride,
       the indispensable material for the manufacture of sulphuric acid,
       and ferric oxide which is again used in the process. In Grant's
       method sodium sulphate is transformed into sodium sulphide, and
       the latter is decomposed by a stream of carbonic anhydride and
       steam, when hydrogen sulphide is disengaged and sodium carbonate
       formed. Gossage prepares Na_{2}S from Na_{2}SO_{4} (by heating it
       with carbon), dissolves it in water and subjects the solution to
       the action of an excess of CO_{2} in coke towers, thus obtaining
       H_{2}S (a gas which gives SO_{2} under perfect combustion, or
       sulphur when incompletely burnt, Chapter XX., Note 6) and
       bicarbonate of sodium; Na_{2}S + 2CO_{2} + 2H_{2}O = H_{2}S +
       2HNaCO_{3}. The latter gives soda and CO_{2} when ignited. This
       process quite eliminates the formation of soda-waste (_see_ Note
       3) and should in my opinion be suitable for the treatment of
       native Na_{2}SO_{4}, like that which is found in the Caucasus, all
       the more since H_{2}S gives sulphur as a bye-product.

       Repeated efforts have been made in recent times to obtain soda
       (and chlorine, _see_ Chapter II., Note 1) from strong solutions of
       salt (Chapter X., Note 23 bis) by the action of an electric
       current, but until now these methods have not been worked out
       sufficiently for practical use, probably partly owing to the
       complicated apparatus needed, and the fact that the chlorine given
       off at the anode corrodes the electrodes and vessels and has but a
       limited industrial application. We may mention that according to
       Hempel (1890) soda in crystals is deposited when an electric
       current and a stream of carbonic acid gas are passed through a
       saturated solution of NaCl.

       Sodium carbonate may likewise be obtained from cryolite (Chapter
       XVII., Note 23) the method of treating this will be mentioned
       under Aluminium.

Of all other industrial processes for manufacturing sodium carbonate,
the _ammonia process_ is the most worthy of mention.[16] In this the
vapours of ammonia, and then an excess of carbonic anhydride, are
directly introduced into a concentrated solution of sodium chloride
in order to form the acid ammonium carbonate, NH_{4}HCO_{3}. Then, by
means of the double saline decomposition of this salt, sodium chloride
is decomposed, and in virtue of its slight solubility acid sodium
carbonate, NaHCO_{3}, is precipitated and ammonium chloride, NH_{4}Cl,
is obtained in solution (with a portion of the sodium chloride and acid
sodium carbonate). The reaction proceeds in the solution owing to the
sparing solubility of the NaHCO_{3} according to the equation NaCl +
NH_{4}HCO_{3} = NH_{4}Cl + NaHCO_{3}. The ammonia is recovered from the
solution by heating with lime or magnesia,[16 bis] and the precipitated
acid sodium carbonate is converted into the normal salt by heating. It
is thus obtained in a very pure state.[17]

  [16] This process (Chapter XVII.) was first pointed out by Turck,
       worked out by Schloesing, and finally applied industrially by
       Solvay. The first (1883) large soda factories erected in Russia
       for working this process are on the banks of the Kama at
       Berezniak, near Ousolia, and belong to Lubimoff. But Russia, which
       still imports from abroad a large quantity of bleaching powder and
       exports a large amount of manganese ore, most of all requires
       works carrying on the Leblanc process. In 1890 a factory of this
       kind was erected by P. K. Oushkoff, on the Kama, near Elagoubi.

  [16 bis] Mond (_see_ Chapter XI., Note 3 bis) separates the NH_{4}Cl
       from the residual solutions by cooling (Chapter X., Note 44);
       ignites the sal-ammoniac and passes the vapour over MgO, and so
       re-obtains the NH_{3}, and forms MgCl_{2}: the former goes back
       for the manufacture of soda, while the latter is employed either
       for making HCl or Cl_{2}.

  [17] Commercial soda ash (calcined, anhydrous) is rarely pure; the
       crystallised soda is generally purer. In order to purify it
       further, it is best to boil a concentrated solution of soda ash
       until two-thirds of the liquid remain, collect the soda which
       settles, wash with cold water, and then shake up with a strong
       solution of ammonia, pour off the residue, and heat. The
       impurities will then remain in the mother liquors, &c.

       Some numerical data may be given for sodium carbonate. The
       specific gravity of the anhydrous salt is 2·48, that of the
       decahydrated salt 1·46. Two varieties are known of the
       heptahydrated salt (Löwel, Marignac, Rammelsberg), which are
       formed together by allowing a saturated solution to cool under a
       layer of alcohol; the one is less stable (like the corresponding
       sulphate) and at 0° has a solubility of 32 parts (of anhydrous
       salt) in 100 water; the other is more stable, and its solubility
       20 parts (of anhydrous salt) per 100 of water. The solubility of
       the decahydrated salt in 100 water = at 0°, 7·0; at 20°, 21·7; at
       30°, 37·2 parts (of anhydrous salt). At 80° the solubility is only
       46·1, at 90° 45·7, at 100°, 45·4 parts (of anhydrous salt). That
       is, it falls as the temperature rises, like Na_{2}SO_{4}. The
       specific gravity (Note 7) of the solutions of sodium carbonate,
       according to the data of Gerlach and Kohlrausch, at 15°/4° is
       expressed by the formula, _s_ = 9,992 + 104·5_p_ + 0·165_p_^2.
       Weak solutions occupy a volume not only less than the sum of the
       volumes of the anhydrous salt and the water, but even less than
       the water contained in them. For instance, 1,000 grams of a 1 p.c.
       solution occupy (at 15°) a volume of 990·4 c.c. (sp. gr. 1·0097),
       but contain 990 grams of water, occupying at 15° a volume of 990·8
       c.c. A similar case, which is comparatively rare occurs also with
       sodium hydroxide, in those dilute solutions for which the factor
       _A_ is greater than 100 if the sp. gr. of water at 4° = 100,000,
       and if the sp. gr. of the solution be expressed by the formula _S_
       = _S__{0} + _Ap_ + _Bp_^2, where _S__{0} is the specific gravity
       of the water. For 5 p.c. the sp. gr. 15°/4° = 1·0520; for 10 p.c.
       1·1057; for 15 p.c. 1·1603. The changes in the sp. gr. with the
       temperature are here almost the same as with solutions of sodium
       chloride with an equal value of _p_.

Sodium carbonate, like sodium sulphate, loses all its water on being
heated, and when anhydrous fuses at a bright-red heat (1098°). A small
quantity of sodium carbonate placed in the loop of a platinum wire
volatilises in the heat of a gas flame, and therefore in the furnaces of
glass works part of the soda is always transformed into the condition
of vapour. Sodium carbonate resembles sodium sulphate in its relation
to water.[18] Here also the greatest solubility is at the temperature
of 37°; both salts, on crystallising at the ordinary temperature,
combine with ten molecules of water, and such crystals of soda, like
crystals of Glauber's salt, fuse at 34°. Sodium carbonate also forms a
supersaturated solution, and, according to the conditions, gives various
combinations with water of crystallisation (mentioned on page 108), &c.

  [18] The resemblance is so great that, notwithstanding the difference
       in the molecular composition of Na_{2}SO_{4} and Na_{2}CO_{3},
       they ought to be classed under the type (NaO)_{2}R, where R =
       SO_{2} or CO. Many other sodium salts also contain 10 mol. H_{2}O.

At a red heat superheated steam liberates carbonic anhydride from
sodium carbonate and forms caustic soda, Na_{2}CO_{3} + H_{2}O = 2NaHO
+ CO_{2}. Here the carbonic anhydride is replaced by water; this
depends on the feebly acid character of carbonic anhydride. By direct
heating, sodium carbonate is only slightly decomposed into sodium oxide
and carbonic anhydride; thus, when sodium carbonate is fused, about 1
per cent. of carbonic anhydride is disengaged.[19] The carbonates of
many other metals--for instance, of calcium, copper, magnesium, iron,
&c.--on being heated lose all their carbonic anhydride. This shows the
considerable basic energy which sodium possesses. With the soluble salts
of most metals, sodium carbonate gives precipitates either of insoluble
carbonates of the metals, or else of the hydroxides (in this latter
case carbonic anhydride is disengaged); for instance, with barium salts
it precipitates an insoluble barium carbonate (BaCl_{2} + Na_{2}CO_{3}
= 2NaCl + BaCO_{3}) and with the aluminium salts it precipitates
aluminium hydroxide, carbonic anhydride being disengaged: 3Na_{2}CO_{3}
+ Al_{2}(SO_{4})_{3} + 3H_{2}O = 3Na_{2}SO_{4} + 2Al(OH)_{3} + 3CO_{2}.
Sodium carbonate, like all the salts of carbonic acid, evolves carbonic
anhydride on treatment with all acids which are to any extent energetic.
But if an acid diluted with water be gradually added to a solution of
sodium carbonate, _at first_ such an evolution does not take place,
because the excess of the carbonic anhydride forms acid sodium carbonate
(sodium bicarbonate), NaHCO_{3}.[20] The acid sodium carbonate is an
unstable salt. Not only when heated alone, but even on being slightly
heated in solution, and also at the ordinary temperature in damp air,
it loses carbonic anhydride and forms the normal salt. And at the
same time it is easy to obtain it in a pure crystalline form, if a
strong solution of sodium carbonate be cooled and a stream of carbonic
anhydride gas passed through it. The acid salt is less soluble in water
than the normal,[21] and therefore a strong solution of the latter
gives crystals of the acid salt if carbonic anhydride be passed through
it. The acid salt may be yet more conveniently formed from effloresced
crystals of sodium carbonate, which, on being considerably heated, very
easily absorb carbonic anhydride.[22] The acid salt crystallises well,
but not, however, in such large crystals as the normal salt; it has a
brackish and not an alkaline taste like that of the normal salt; its
reaction is feebly alkaline, nearly neutral. At 70° its solution begins
to lose carbonic anhydride, and on boiling the evolution becomes very
abundant. From the preceding remarks it is clear that in most reactions
this salt, especially when heated, acts similarly to the normal salt,
but has, naturally, some distinction from it. Thus, for example, if a
solution of sodium carbonate be added to a normal magnesium salt, a
turbidity (precipitate) is formed of magnesium carbonate. MgCO_{3}. No
such precipitate is formed by the acid salt, because magnesium carbonate
is soluble in the presence of an excess of carbonic anhydride.

  [19] According to the observations of Pickering. According to Rose,
       when solutions of sodium carbonate are boiled a certain amount of
       carbonic anhydride is disengaged.

  [20] The composition of this salt, however, may be also represented
       as a combination of carbonic acid, H_{2}CO_{3}, with the normal
       salt, Na_{2}CO_{3}, just as the latter also combines with water.
       Such a combination is all the more likely because (1) there exists
       another salt, Na_{2}CO_{3},2NaHCO_{3},2H_{2}O (sodium
       sesquicarbonate), obtained by cooling a boiling solution of sodium
       bicarbonate, or by mixing this salt with the normal salt; but the
       formula of this salt cannot be derived from that of normal
       carbonic acid, as the formula of the bicarbonate can. At the same
       time the sesqui-salt has all the properties of a definite
       compound; it crystallises in transparent crystals, has a constant
       composition, its solubility (at 0° in 100 of water, 12·6 of
       anhydrous salt) differs from the solubility of the normal and acid
       salts; it is found in nature, and is known by the names of _trona_
       and _urao_. The observations of Watts and Richards showed (1886)
       that on pouring a strong solution of the acid salt into a solution
       of the normal salt saturated by heating, crystals of the salt
       NaHCO_{3},Na_{2}CO_{3},2H_{2}O may be easily obtained, as long as
       the temperature is above 35°. The natural urao (Boussingault) has,
       according to Laurent, the same composition. This salt is very
       stable in air, and may be used for purifying sodium carbonate on
       the large scale. Such compounds have been little studied from a
       theoretical point of view, although particularly interesting,
       since in all probability they correspond with ortho-carbonic acid,
       C(OH)_{4}, and at the same time correspond with double salts like
       astrakhanite (Chapter XIV., Note 25). (2) Water of crystallisation
       does not enter into the composition of the crystals of the acid
       salt, so that on its formation (occurring only at low
       temperatures, as in the formation of crystalline compounds with
       water) the water of crystallisation of the normal salt separates
       and the water is, as it were, replaced by the elements of carbonic
       acid. If anhydrous sodium carbonate be mixed with the amount of
       water requisite for the formation of Na_{2}CO_{3},H_{2}O, this
       salt will, when powdered, absorb CO_{2} as easily at the ordinary
       temperature as it does water.

  [21] 100 parts of water at 0° dissolve 7 parts of the acid salt, which
       corresponds with 4·3 parts of the anhydrous normal salt, but at 0°
       100 parts of water dissolve 7 parts of the latter. The solubility
       of the bi-or acid salt varies with considerable regularity; 100
       parts of water dissolves at 15° 9 parts of the salt, at 30° 11
       parts.

       The ammonium, and more especially the calcium, salt, is much more
       soluble in water. The ammonia process (_see_ p. 524) is founded
       upon this. Ammonium bicarbonate (acid carbonate) at 0° has a
       solubility of 12 parts in 100 water, at 30° of 27 parts. The
       solubility therefore increases very rapidly with the temperature.
       And its saturated solution is more stable than a solution of
       sodium bicarbonate. In fact, saturated solutions of these salts
       have a gaseous tension like that of a mixture of carbonic
       anhydride and water--namely, at 15° and at 50°, for the sodium
       salt 120 and 750 millimetres, for the ammonium salt 120 and 563
       millimetres. These data are of great importance in understanding
       the phenomena connected with the ammonia process. They indicate
       that with an increased pressure the formation of the sodium salt
       ought to increase if there be an excess of ammonium salt.

  [22] Crystalline sodium carbonate (broken into lumps) also absorbs
       carbonic anhydride, but the water contained in the crystals is
       then disengaged: Na_{2}CO_{3},10H_{2}O + CO_{2} =
       Na_{2}CO_{3},H_{2}CO_{3} + 9H_{2}O, and dissolves part of the
       carbonate; therefore part of the sodium carbonate passes into
       solution together with all the impurities. When it is required to
       avoid the formation of this solution, a mixture of ignited and
       crystalline sodium carbonate is taken. Sodium bicarbonate is
       prepared chiefly for medicinal use, and is then often termed
       _carbonate of soda_, also, for instance, in the so-called soda
       powders, for preparing certain artificial mineral waters, for the
       manufacture of digestive lozenges like those made at Essentuki,
       Vichy, &c.

Sodium carbonate is used for the preparation of _caustic soda_[23]--that
is, the hydrate of sodium oxide, or the alkali which corresponds to
sodium. For this purpose the action of lime on a solution of sodium
carbonate is generally made use of. The process is as follows: a weak,
generally 10 per cent., solution of sodium carbonate is taken,[24] and
boiled in a cast-iron, wrought-iron, or silver boiler (sodium hydroxide
does not act on these metals), and lime is added, little by little,
during the boiling. This latter is soluble in water, although but very
slightly. The clear solution becomes turbid on the addition of the lime
because a precipitate is formed; this precipitate consists of calcium
carbonate, almost insoluble in water, whilst caustic soda is formed
and remains in solution. The decomposition is effected according to
the equation: Na_{2}CO_{3} + Ca(HO)_{2} = CaCO_{3} + 2NaHO. On cooling
the solution the calcium carbonate easily settles as a precipitate,
and the clear solution or alkali above it contains the easily soluble
sodium hydroxide formed in the reaction.[25] After the necessary
quantity of lime has been added, the solution is allowed to stand, and
is then decanted off and evaporated in cast or wrought iron boilers,
or in silver pans if a perfectly pure product is required.[26] The
evaporation cannot be conducted in china, glass, or similar vessels,
because caustic soda attacks these materials, although but slightly. The
solution does not crystallise on evaporation, because the solubility
of caustic soda when hot is very great, but crystals containing water
of crystallisation may be obtained by cooling. If the evaporation
of the alkali be conducted until the specific gravity reaches 1·38,
and the liquid is then cooled to 0°, transparent crystals appear
containing 2NaHO,7H_{2}O; they fuse at +6°.[27] If the evaporation be
conducted so long as water is disengaged, which requires a considerable
amount of heat, then, on cooling, the hydroxide, NaHO, solidifies
in a semi-transparent crystalline mass,[28] which eagerly absorbs
moisture and carbonic anhydride from the air.[29] Its specific gravity
is 2·13;[30] it is easily soluble in water, with disengagement of a
considerable quantity of heat.[31] A saturated solution at the ordinary
temperature has a specific gravity of about 1·5, contains about 45 per
cent. of sodium hydroxide, and boils at 130°; at 55° water dissolves
an equal weight of it.[32] Caustic soda is not only soluble in water
but in alcohol, and even in ether. Dilute solutions of sodium hydroxide
produce a soapy feeling on the skin because the active base of soap
consists of caustic soda.[33] Strong solutions have a corroding action.

  [23] In chemistry, sodium oxide is termed 'soda,' which word must be
       carefully distinguished from the word sodium, meaning the metal.

  [24] With a small quantity of water, the reaction either does not
       take place, or even proceeds in the reverse way--that is, sodium
       and potassium hydroxides remove carbonic anhydride from calcium
       carbonate (Liebig, Watson, Mitscherlich, and others). The
       influence of the mass of water is evident. According to Gerberts,
       however, strong solutions of sodium carbonate are decomposed by
       lime, which is very interesting if confirmed by further
       investigation.

  [25] As long as any undecomposed sodium carbonate remains in solution,
       excess of acid added to the solution disengages carbonic
       anhydride, and the solution after dilution gives a white
       precipitate with a barium salt soluble in acids, showing the
       presence of a carbonate in solution (if there be sulphate present,
       it also forms a white precipitate, but this is insoluble in
       acids). For the decomposition of sodium carbonate, milk of
       lime--that is, slaked slime suspended in water--is employed.
       Formerly pure sodium hydroxide was prepared (according to
       Berthollet's process) by dissolving the impure substance in
       alcohol (sodium carbonate and sulphate are not soluble), but now
       that metallic sodium has become cheap and is purified by
       distillation, _pure caustic soda_ is prepared by acting on a small
       quantity of water with sodium. Perfectly pure sodium hydroxide may
       also be obtained by allowing strong solutions to crystallise (in
       the cold) (Note 27).

       In alkali works where the Leblanc process is used, caustic soda is
       prepared directly from the alkali remaining in the mother liquors
       after the separation of the sodium carbonate by evaporation (Note
       14). If excess of lime and charcoal have been used, much sodium
       hydroxide maybe obtained. After the removal as much as possible of
       the sodium carbonate, a red liquid (from iron oxide) is left,
       containing sodium hydroxide mixed with compounds of sulphur and of
       cyanogen (_see_ Chapter IX.) and also containing iron. This red
       alkali is evaporated and air is blown through it, which oxidises
       the impurities (for this purpose sometimes sodium nitrate is
       added, or bleaching powder, &c.) and leaves fused caustic soda.
       The fused mass is allowed to settle in order to separate the
       ferruginous precipitate, and poured into iron drums, where the
       sodium hydroxide solidifies. Such caustic soda contains about 10
       p.c. of water in excess and some saline impurities, but when
       properly manufactured is almost free from carbonate and from iron.
       The greater part of the caustic soda, which forms so important an
       article of commerce, is manufactured in this manner.

  [26] Löwig gave a method of preparing sodium hydroxide from sodium
       carbonate by heating it to a dull red heat with an excess of
       ferric oxide. Carbonic anhydride is given off, and warm water
       extracts the caustic soda from the remaining mass. This reaction,
       as experiment shows, proceeds very easily, and is an example of
       contact action similar to that of ferric oxide on the
       decomposition of potassium chlorate. The reason of this may be
       that a small quantity of the sodium carbonate enters into double
       decomposition with the ferric oxide, and the ferric carbonate
       produced is decomposed into carbonic anhydride and ferric oxide,
       the action of which is renewed. Similar explanations expressing
       the _reason_ for a reaction really adds but little to that
       elementary conception of contact which, according to my opinion,
       consists in the change of motion of the atoms in the molecules
       under the influence of the substance in contact. In order to
       represent this clearly it is sufficient, for instance, to imagine
       that in the sodium carbonate the elements CO_{2} move in a circle
       round the elements Na_{2}O, but at the points of contact with
       Fe_{2}O_{3} the motion becomes elliptic with a long axis, and at
       some distance from Na_{2}O the elements of CO_{2} are parted, not
       having the faculty of attaching themselves to Fe_{2}O_{3}.

  [27] By allowing strong solutions of sodium hydroxide to crystallise in
       the cold, impurities--such as, for instance, sodium sulphate--may
       be separated from them. The fused crystallo-hydrate 2NaHO,7H_{2}O
       forms a solution having a specific gravity of 1·405 (Hermes). The
       crystals on dissolving in water produce cold, while NaHO produces
       heat. Besides which Pickering obtained hydrates with 1, 2, 4, 5,
       and 7 H_{2}O.

  [28] In solid caustic soda there is generally an excess of water beyond
       that required by the formula NaHO. The caustic soda used in
       laboratories is generally cast in sticks, which are broken into
       pieces. It must be preserved in carefully closed vessels, because
       it absorbs water and carbonic anhydride from the air.

  [29] By the way it changes in air it is easy to distinguish caustic
       soda from caustic potash, which in general resembles it. Both
       alkalis absorb water and carbonic anhydride from the air, but
       caustic potash forms a deliquescent mass of potassium carbonate,
       whilst caustic soda forms a dry powder of efflorescent salt.

  [30] As the molecular weight of NaHO = 40, the volume of its molecule =
       40/2·13 = 18·5, which very nearly approaches the volume of a
       molecule of water. The same rule applies to the compounds of
       sodium in general--for instance, its salts have a molecular volume
       approaching the volume of the acids from which they are derived.

  [31] The molecular quantity of sodium hydroxide (40 grams), on being
       dissolved in a large mass (200 gram molecules) of water, develops,
       according to Berthelot 9,780, and according to Thomsen 9,940,
       heat-units, but at 100° about 13,000 (Berthelot). Solutions of
       NaHO + _n_H_{2}O, on being mixed with water, evolve heat if they
       contain less than 6H_{2}O, but if more they absorb beat.

  [32] The specific gravity of solutions of sodium hydroxide at 15°/4° is
       given in the short table below:--

          NaHO, p.c.    5      10     15     20     30     40
          Sp. gr.     1·057  1·113  1·169  1·224  1·331  1·436


       1,000 grams of a 5 p.c. solution occupies a volume of 946 c.c.;
       that is, less than the water serving to make the solution (_see_
       Note 18).

  [33] Sodium hydroxide and some other alkalis are capable of
       hydrolysing--saponifying, as it is termed--the compounds of acids
       with alcohols. If RHO (or R(HO)_{_n_}) represent the composition
       of an alcohol--that is, of the hydroxide of a hydrocarbon
       radicle--and QHO an acid, then the compound of the acid with the
       alcohol or ethereal salt of the given acid will have the
       composition RQO. Ethereal salts, therefore, present a likeness to
       metallic salts, just as alcohols resemble basic hydroxides. Sodium
       hydroxide acts on ethereal salts in the same way that it acts on
       the majority of metallic salts--namely, it liberates alcohol, and
       forms the sodium salt of that acid which was in the ethereal salt.
       The reaction takes place in the following way:--

                    RQO    +  NaHO   =  NaQO  +   RHO
                  Ethereal   Caustic   Sodium   Alcohol
                    salt      soda      salt

       Such a decomposition is termed saponification; similar reactions
       were known very long ago for the ethereal salts corresponding with
       glycerin, C_{3}H_{5}(OH)_{3} (Chapter IX.), found in animals and
       plants, and composing what are called fats or oils. Caustic soda,
       acting on fat and oil, forms glycerin, and sodium salts of those
       acids which were in union with the glycerin in the fat, as
       Chevreul showed at the beginning of this century. The sodium salts
       of the fatty acids are commonly known as soaps. That is to say,
       soap is made from fat and caustic soda, glycerin being separated
       and a sodium salt or soap formed. As glycerin is usually found in
       union with certain acids, so also are the sodium salts of the same
       acids found in soap. The greater part of the acids found in
       conjunction with glycerin in fats are the solid palmitic and
       stearic acids, C_{16}H_{32}O_{2} and C_{18}H_{38}O_{2}, and the
       liquid oleic acid, C_{18}H_{34}O_{2}. In preparing soap the fatty
       substances are mixed with a solution of caustic soda until an
       emulsion is formed; the proper quantity of caustic soda is then
       added in order to produce saponification on heating, the soap
       being separated from the solution either by means of an excess of
       caustic soda or else by common salt, which displaces the soap from
       the aqueous solution (salt water does not dissolve soap, neither
       does it form a lather). Water acting on soap partly decomposes it
       (because the acids of the soap are feeble), and the alkali set
       free acts during the application of soap. Hence it may be replaced
       by a very feeble alkali. Strong solutions of alkali corrode the
       skin and tissues. They are not formed from soap, because the
       reaction is reversible, and the alkali is only set free by the
       excess of water. Thus we see how the teaching of Berthollet
       renders it possible to understand many phenomena which occur in
       every-day experience (_see_ Chapter IX., Note 15).

The chemical _reactions of sodium hydroxide_ serve as a type for those
of a whole class of alkalis--that is, of soluble basic hydroxides,
MOH. The solution of sodium hydroxide is a very caustic liquid--that
is to say, it acts in a destructive way on most substances, for
instance on most organic tissues--hence caustic soda, like all soluble
alkalis, is a poisonous substance; acids, for example hydrochloric,
serve as antidotes. The action of caustic soda on bones, fat, starch,
and similar vegetable and animal substances explains its action on
organisms. Thus bones, when plunged into a weak solution of caustic
soda, fall to powder,[34] and evolve a smell of ammonia, owing to the
caustic soda changing the gelatinous organic substance of the bones
(which contains carbon, hydrogen, nitrogen, oxygen, and sulphur, like
albumin), dissolving it and in part destroying it, whence ammonia is
disengaged. Fats, tallow, and oils become saponified by a solution of
caustic soda--that is to say, they form with it _soaps_ soluble in
water, or sodium salts of the organic acids contained in the fats.[35]
The most characteristic reactions of sodium hydroxide are determined by
the fact that it _saturates all acids, forming salts with them_, which
are almost all soluble in water, and in this respect caustic soda is as
characteristic amongst the bases as nitric acid is among the acids. It
is impossible to detect sodium by means of the formation of precipitates
of insoluble sodium salts, as may be done with other metals, many of
whose salts are but slightly soluble. The powerful alkaline properties
of caustic soda determine its capacity for combining with even the
feeblest acids, its property of disengaging ammonia from ammonium salts,
its faculty of forming precipitates from solutions of salts whose bases
are insoluble in water, &c. If a solution of the salt of almost any
metal be mixed with caustic soda, then a soluble sodium salt will be
formed, and an insoluble hydroxide of the metal will be separated--for
instance, copper nitrate yields copper hydroxide, Cu(NO_{3})_{2} + 2NaHO
= Cu(HO)_{2} + 2NaNO_{3}. Even many _basic oxides_ precipitated by
caustic soda _are capable_ of _combining_ with it and forming soluble
compounds, and therefore caustic soda in the presence of salts of such
metals first forms a precipitate of hydroxide, and then, employed in
excess, dissolves this precipitate. This phenomenon occurs, for example,
when caustic soda is added to the salts of aluminium. This shows the
property of such an alkali as caustic soda of combining not only with
acids, but also with feeble basic oxides. For this reason caustic soda
_acts on most elements_ which are capable of forming acids or oxides
similar to them; thus the metal aluminium gives hydrogen with caustic
soda in consequence of the formation of alumina, which combines with
the caustic soda--that is, in this case, the caustic alkali acts on the
metal just as sulphuric acid does on Fe or Zn. If caustic soda acts
in this manner on a metalloid capable of combining with the hydrogen
evolved (aluminium does not give a compound with hydrogen), then it
forms such a hydrogen compound. Thus, for instance, phosphorus acts in
this way on caustic soda, yielding hydrogen phosphide. When the hydrogen
compound disengaged is capable of combining with the alkali, then,
naturally, a salt of the corresponding acid is formed. For example,
chlorine and sulphur act in this way on caustic soda. Chlorine, with the
hydrogen of the caustic soda, forms hydrochloric acid, and the latter
forms common salt with the sodium hydroxide, whilst the other atom in
the molecule of chlorine, Cl_{2}, takes the place of the hydrogen, and
forms the hypochlorite, NaClO. In the same way, by the action of sodium
hydroxide on sulphur, hydrogen sulphide is formed, which acts on the
soda forming sodium _sulphide_, in addition to which sodium thiosulphate
is formed (_see_ Chapter XX.) By virtue of such reactions, sodium
hydroxide acts on many metals and non-metals. Such action is often
accelerated by the presence of the oxygen of the air, as by this means
the formation of acids and oxides rich in oxygen is facilitated. Thus
many metals and their lower oxides, in the presence of an alkali, absorb
oxygen and form acids. Even manganese peroxide, when mixed with caustic
soda, is capable of absorbing the oxygen of the air, and forming sodium
manganate. Organic acids when heated with caustic soda give up to it the
elements of carbonic anhydride, forming sodium carbonate, and separating
that hydrocarbon group which exists, in combination with carbonic
anhydride, in the organic acid.

  [34] On this is founded the process of Henkoff and Engelhardt for
       treating bones. The bones are mixed with ashes, lime, and water;
       it is true that in this case more potassium hydroxide than sodium
       hydroxide is formed, but their action is almost identical.

  [35] As explained in Note 33.

Thus sodium hydroxide, like the soluble alkalis in general, ranks
amongst the most active substances in the chemical sense of the term,
and but few substances are capable of resisting it. Even siliceous
rocks, as we shall see further on, are transformed by it, forming when
fused with it vitreous slags. Sodium hydroxide (like ammonium and
potassium hydroxides), as a typical example of the basic hydrates, in
distinction from many other basic oxides, easily _forms acid salts_
with acids (for instance, NaHSO_{4}, NaHCO_{3}), and does not form any
basic salts at all; whilst many less energetic bases, such as the oxides
of copper and lead, easily form basic salts, but acid salts only with
difficulty. This capability of forming acid salts, particularly with
polybasic acids, may be explained by the energetic basic properties
of sodium hydroxide, contrasted with the small development of these
properties in the bases which easily form basic salts. An energetic
base is capable of retaining a considerable quantity of acid, which a
slightly energetic base would not have the power of doing. Also, as will
be shown in the subsequent chapters, sodium belongs to the univalent
metals, being exchangeable for hydrogen atom for atom--that is, amongst
metals sodium may, like chlorine amongst the non-metals, serve as the
representative of the univalent properties. Most of the elements which
are not capable of forming acid salts are bivalent. Whence it may be
understood that in a bibasic acid--for instance, carbonic, H_{2}CO_{3},
or sulphuric, H_{2}SO_{4}--the hydrogen may be exchanged, atom for atom,
for sodium, and yield an acid salt by means of the first substitution,
and a normal salt by means of the second--for instance, NaHSO_{4}, and
Na_{2}SO_{4}, whilst such bivalent metals as calcium and barium do not
form acid salts because one of their atoms at once takes the place of
both hydrogen atoms, forming, for example, CaCO_{3} and CaSO_{4}.[35 bis]

  [35 bis] It might be expected, from what has been mentioned above, that
       bivalent metals would easily form acid salts with acids containing
       more than two atoms of hydrogen--for instance, with tribasic
       acids, such as phosphoric acid, H_{3}PO_{4}--and actually such
       salts do exist; but all such relations are complicated by the fact
       that the character of the base very often changes and becomes
       weakened with the increase of valency and the change of atomic
       weight; the feebler bases (like silver oxide), although
       corresponding with univalent metals, do not form acid salts, while
       the feeblest bases (CuO, PbO, &c.) easily form basic salts, and
       notwithstanding their valency do not form acid salts which are in
       any degree stable--that is, which are undecomposable by water.
       Basic and acid salts ought to be regarded rather as compounds
       similar to crystallo-hydrates, because such acids as sulphuric
       form with sodium not only an acid and a normal salt, as might be
       expected from the valency of sodium, but also salts containing a
       greater quantity of acid. In sodium sesquicarbonate we saw an
       example of such compounds. Taking all this into consideration, we
       must say that the property of more or less easily forming acid
       salts depends more upon the energy of the base than upon its
       valency, and the best statement is that _the capacity of a base
       for forming acid and basic salts is characteristic_, just as the
       faculty of forming compounds with hydrogen is characteristic of
       elements.

We have seen the transformation of common salt into sodium sulphate,
of this latter into sodium carbonate, and of sodium carbonate into
caustic soda. Lavoisier still regarded sodium hydroxide as an element,
because he was unacquainted with its decomposition with the formation
of metallic sodium, which separates the hydrogen from water, reforming
caustic soda.

The preparation of _metallic sodium_ was one of the greatest discoveries
in chemistry, not only because through it the conception of elements
became broader and more correct, but especially because in sodium,
chemical properties were observed which were but feebly shown in the
other metals more familiarly known. This discovery was made in 1807
by the English chemist _Davy_ by means of the galvanic current. By
connecting with the positive pole (of copper or carbon) a piece of
caustic soda (moistened in order to obtain electrical conductivity),
and boring a hole in it filled with mercury connected with the negative
pole of a strong Volta's pile, Davy observed that on passing the current
a peculiar metal dissolved in the mercury, less volatile than mercury,
and capable of decomposing water, again forming caustic soda. In this
way (by analysis and synthesis) Davy demonstrated the compound nature
of alkalis. On being decomposed by the galvanic current, caustic soda
disengages hydrogen and sodium at the negative pole and oxygen at the
positive pole. Davy showed that the metal formed volatilises at a red
heat, and this is its most important physical property in relation to
its extraction, all later methods being founded on it. Besides this Davy
observed that sodium easily oxidises, its vapour taking fire in air,
and the latter circumstance was for a long time an obstacle to the easy
preparation of this metal. The properties of sodium were subsequently
more thoroughly investigated by Gay-Lussac and Thénard, who observed
that metallic iron at a high temperature was capable of reducing caustic
soda to sodium.[36] Brunner latterly discovered that not only iron,
but also charcoal, has this property, although hydrogen has not.[37]
But still the methods of extracting sodium were very troublesome,
and consequently it was a great rarity. The principal obstacle
to its production was that an endeavour was made to condense the
easily-oxidising vapours of sodium in vacuo in complicated apparatus.
For this reason, when Donny and Maresca, having thoroughly studied the
matter, constructed a specially simple condenser, the production of
sodium was much facilitated. Furthermore, in practice the most important
epoch in the history of the production of sodium is comprised in the
investigation of Sainte-Claire Deville, who avoided the complex methods
in vogue up to that time, and furnished those simple means by which the
production of sodium is now rendered feasible in chemical works.

  [36] Deville supposes that such a decomposition of sodium hydroxide
       by metallic iron depends solely on the dissociation of the alkali
       at a white heat into sodium, hydrogen, and oxygen. Here the part
       played by the iron is only that it retains the oxygen formed,
       otherwise the decomposed elements would again reunite upon
       cooling, as in other cases of dissociation. If it be supposed that
       the temperature at the commencement of the dissociation of the
       iron oxides is higher than that of sodium oxide, then the
       decomposition may be explained by Deville's hypothesis. Deville
       demonstrates his views by the following experiment:--An iron
       bottle, filled with iron borings, was heated in such a way that
       the upper part became red hot, the lower part remaining cooler;
       sodium hydroxide was introduced into the upper part. The
       decomposition was then effected--that is, sodium vapours were
       produced (this experiment was really performed with potassium
       hydroxide). On opening the bottle it was found that the iron in
       the upper part was not oxidised, but only that in the lower part.
       This may be explained by the decomposition of the alkali into
       sodium, hydrogen, and oxygen taking place in the upper part,
       whilst the iron in the lower part absorbed the oxygen set free. If
       the whole bottle be subjected to the same moderate heat as the
       lower extremity, no metallic vapours are formed. In that case,
       according to the hypothesis, the temperature is insufficient for
       the dissociation of the sodium hydroxide.

  [37] It has been previously remarked (Chapter II. Note 9) that Beketoff
       showed the displacement of sodium by hydrogen, not from sodium
       hydroxide but from the oxide Na_{2}O; then, however, only one half
       is displaced, with the formation of NaHO.

For the production of sodium according to Deville's method, a mixture
of anhydrous sodium carbonate (7 parts), charcoal (two parts), and lime
or chalk (7 parts) is heated. This latter ingredient is only added in
order that the sodium carbonate, on fusing, shall not separate from
the charcoal.[38] The chalk on being heated loses carbonic anhydride,
leaving infusible lime, which is permeated by the sodium carbonate and
forms a thick mass, in which the charcoal is intimately mixed with the
sodium carbonate. When the charcoal is heated with the sodium carbonate,
at a white heat, carbonic oxide and vapours of sodium are disengaged,
according to the equation:

                    Na_{2}CO_{3} + 2C = Na_{2} + 3CO

  [38] Since the close of the eighties in England, where the preparation
       of sodium is at present carried out on a large commercial scale
       (from 1860 to 1870 it was only manufactured in a few works in
       France), it has been the practice to add to Deville's mixture
       iron, or iron oxide which with the charcoal gives metallic and
       carburetted iron, which still further facilitates the
       decomposition. At present a kilogram of sodium may be purchased
       for about the same sum (2/-) as a gram cost thirty years ago.
       Castner, in England, greatly improved the manufacture of sodium in
       large quantities, and so cheapened it as a reducing agent in the
       preparation of metallic aluminium. He heated a mixture of 44 parts
       of NaHO, and 7 parts of carbide of iron in large iron retorts at
       1,000° and obtained about 6-1/2 parts of metallic sodium. The
       reaction proceeds more easily than with carbon or iron alone, and
       the decomposition of the NaHO proceeds according to the equation:
       3NaHO + C = Na_{2}CO_{3} + 3H + Na. Subsequently, in 1891,
       aluminium was prepared by electrolysis (_see_ Chapter XVII.), and
       metallic sodium found two new uses; (1) for the manufacture of
       peroxide of sodium (see later on) which is used in bleaching
       works, and (2) in the manufacture of potassium and sodium cyanide
       from yellow prussiate (Chapter XIII., Note 12).

[Illustration: FIG. 70.--Manufacture of sodium by Deville's process.
A C, iron tube containing a mixture of soda, charcoal, and chalk. B,
condenser.]

On cooling the vapours and gases disengaged, the vapours condense into
molten metal (in this form sodium does not easily oxidise, whilst in
vapour it burns) and the carbonic oxide remains as gas.

[Illustration: FIG. 71.--Donny and Maresca's sodium condenser, consisting
of two cast-iron plates screwed together.]

In sodium works an iron tube, about a metre long and a decimeter in
diameter, is made out of boiler plate. The pipe is luted into a furnace
having a strong draught, capable of giving a high temperature, and the
tube is charged with the mixture required for the preparation of sodium.
One end of the tube is closed with a cast-iron stopper A with clay
luting, and the other with the cast-iron stopper C provided with an
aperture. On heating, first of all the moisture contained in the various
substances is given off, then carbonic anhydride and the products of the
dry distillation of the charcoal, then the latter begins to act on the
sodium carbonate, and carbonic oxide and vapours of sodium appear. It is
easy to observe the appearance of the latter, because on issuing from
the aperture in the stopper C they take fire spontaneously and burn with
a very bright yellow flame. A pipe is then introduced into the aperture
C, compelling the vapours and gases formed to pass through the condenser
B. This condenser consists of two square cast-iron trays, A and A´, fig.
71, with wide edges firmly screwed together. Between these two trays
there is a space in which the condensation of the vapours of sodium is
effected, the thin metallic walls of the condenser being cooled by the
air but remaining hot enough to preserve the sodium in a liquid state,
so that it does not choke the apparatus, but continually flows from
it. The vapours of sodium, condensing in the cooler, flow in the shape
of liquid metal into a vessel containing some non-volatile naphtha or
hydrocarbon. This is used in order to prevent the sodium oxidising as
it issues from the condenser at a somewhat high temperature. In order
to obtain sodium of a pure quality it is necessary to distil it once
more, which may even be done in porcelain retorts, but the distillation
must be conducted in a stream of some gas on which sodium does not
act, for instance in a stream of nitrogen; carbonic anhydride is not
applicable, because sodium partially decomposes it, absorbing oxygen
from it. Although the above described methods of preparing sodium by
chemical means have proved very convenient in practice, still it is now
(since 1893) found profitable in England to obtain it (to the amount of
several tons a week) by Davy's classical method, _i.e._ by the action of
an electric current at a moderately high temperature, because the means
for producing an electric current (by motors and dynamos) now render
this quite feasible. This may be regarded as a sign that in process of
time many other technical methods for producing various substances by
_decomposition_ may be profitably carried on by electrolysis.

Pure sodium is a lustrous metal, white as silver, soft as wax; it
becomes brittle in the cold. In ordinary moist air it quickly tarnishes
and becomes covered with a film of hydroxide, NaHO, formed at the
expense of the water in the air. In perfectly dry air sodium retains its
lustre for an indefinite time. Its density at the ordinary temperature
is equal to 0·98, so that it is lighter than water; it fuses very
easily at a temperature of 95°, and distils at a bright red heat (742°
according to Perman, 1889). Scott (1887) determined the density of
sodium vapour and found it to be nearly 12 (if H = 1). This shows that
its molecule contains one atom (like mercury and cadmium) Na.[38 bis] It
forms alloys with most metals, combining with them, heat being sometimes
evolved and sometimes absorbed. Thus, if sodium (having a clean surface)
be thrown into mercury, especially when heated, there is a flash, and
such a considerable amount of heat is evolved that part of the mercury
is transformed into vapour.[39] Compounds or solutions of sodium in
mercury, or _amalgams_ of sodium, even when containing 2 parts of sodium
to 100 parts of mercury, are solids. Only those amalgams which are the
very poorest in sodium are liquid. Such alloys of sodium with mercury
are often used instead of sodium in chemical investigations, because in
combination with mercury sodium is not easily acted on by air, and is
heavier than water, and therefore more convenient to handle, whilst at
the same time it retains the principal properties of sodium,[40] for
instance it decomposes water, forming NaHO.

  [38 bis] This is also shown by the fall in the temperature of
       solidification of tin produced by the addition of sodium (and also
       Al and Zn). Heycock and Neville (1889).

  [39] By dissolving sodium amalgams in water and acids, and deducting
       the heat of solution of the sodium, Berthelot found that _for each
       atom of the sodium_ in amalgams containing a larger amount of
       mercury than NaHg_{5}, the amount of heat evolved increases, after
       which the heat of formation falls, and the heat evolved decreases.
       In the formation of NaHg_{5} about 18,500 calories are evolved;
       when NaHg_{3} is formed, about 14,000; and for NaHg about 10,000
       calories. Kraft regarded the definite crystalline amalgam as
       having the composition of NaHg_{6}, but at the present time, in
       accordance with Grimaldi's results, it is thought to be NaHg_{5}.
       A similar amalgam is very easily obtained if a 3 p.c. amalgam be
       left several days in a solution of sodium hydroxide until a
       crystalline mass is formed, from which the mercury may be removed
       by strongly pressing in chamois leather. This amalgam with a
       solution of potassium hydroxide forms a potassium amalgam,
       KHg_{10}. It may be mentioned here that the latent heat of fusion
       (of atomic quantities) of Hg = 360 (Personne), Na = 730 (Joannis),
       and K = 610 calories (Joannis).

  [40] Alloys are so similar to solutions (exhibiting such complete
       parallelism in properties) that they are included in the same
       class of so-called indefinite compounds. But in alloys, as
       substances passing from the liquid to the solid state, it is
       easier to discover the formation of definite chemical compounds.
       Besides the alloys of Na with Hg, those with tin (Bailey 1892
       found Na_{2}Sn), lead (NaPb), bismuth (Na_{3}Bi), &c. (Joannis
       1892 and others) have been investigated.

It is easy to form an alloy of mercury and sodium having a crystalline
structure, and a definite atomic composition, NaHg_{5}. The alloy
of sodium with hydrogen or _sodium hydride_, Na_{2}H, which has the
external appearance of a metal,[41] is a most instructive example of
the characteristics of alloys. At the ordinary temperature sodium does
not absorb hydrogen, but from 300° to 421° the absorption takes place
at the ordinary pressure (and at an increased pressure even at higher
temperatures), as shown by Troost and Hautefeuille (1874). One volume of
sodium absorbs as much as 238 volumes of hydrogen. The metal increases
in volume, and when once formed the alloy can be preserved for some time
without change at the ordinary temperature. The appearance of sodium
hydride resembles that of sodium itself; it is as soft as this latter,
when heated it becomes brittle, and decomposes above 300°, evolving
hydrogen. In this decomposition all the phenomena of dissociation are
very clearly shown--that is, the hydrogen gas evolved has a definite
tension[42] corresponding with each definite temperature. This confirms
the fact that the formation of substances capable of dissociation can
only be accomplished within the dissociation limits. Sodium hydride
melts more easily than sodium itself, and then does not undergo
decomposition if it is in an atmosphere of hydrogen. It oxidises easily
in air, but not so easily as potassium hydride. The chemical reactions
of sodium are retained in its hydride, and, if we may so express it,
they are even increased by the addition of hydrogen. At all events,
in the properties of sodium hydride[43] we see other properties than
in such hydrogen compounds as HCl, H_{2}O, H_{3}N, H_{4}C, or even in
the gaseous metallic hydrides AsH_{3}, TeH_{2}. Platinum, palladium,
nickel, and iron, in absorbing hydrogen form compounds in which hydrogen
is in a similar state. In them, as in sodium hydride, the hydrogen is
compressed, absorbed, occluded (Chapter II.)[43 bis]

  [41] Potassium forms a similar compound, but lithium, under the same
       circumstances, does not.

  [42] The tension of dissociation of hydrogen _p_, in millimetres of
       mercury, is:--

                             _t_ = 330°  350°  400°   430°
                for Na_{2}H  _p_ =  28    57   447    910
                for K_{2}H          45    72   548   1100


  [43] In general, during the formation of alloys the volumes change
       very slightly, and therefore from the volume of Na_{2}H some idea
       may be formed of the volume of hydrogen in a solid or liquid
       state. Even Archimedes concluded that there was gold in an alloy
       of copper and gold by reason of its volume and density. From the
       fact that the density of Na_{2}H is equal to 0·959, it may be seen
       that the volume of 47 grams (the gram molecule) of this compound =
       49·0 c.c. The volume of 46 grams of sodium contained in the
       Na_{2}H (the density under the same conditions being 0·97) is
       equal to 47·4 c.c. Therefore the volume of 1 gram of hydrogen in
       Na_{2}H is equal to 1·6 c.c., and consequently the density of
       metallic hydrogen, or the weight of 1 c.c., approaches 0·6 gram.
       This density is also proper to the hydrogen alloyed with potassium
       and palladium. Judging from the scanty information which is at
       present available, liquid hydrogen near its absolute boiling point
       (Chapter II.) has a much lower density.

  [43 bis] We may remark that at low temperatures Na absorbs NH_{3} and
       forms (NH_{3}Na)_{2} (_see_ Chapter VI., Note 14); this substance
       absorbs CO and gives (NaCO)n (Chapter IX., Note 31), although by
       itself Na does not combine directly with CO (but K does).

The most important chemical property of sodium is its power of easily
decomposing water and _evolving hydrogen_ from the majority of the
hydrogen compounds, and especially from all acids, and hydrates in which
hydroxyl must be recognised. This depends on its power of combining
with the elements which are in combination with the hydrogen. We
already know that sodium disengages hydrogen, not only from water,
hydrochloric acid,[44] and all other acids, but also from ammonia,[44 bis]
with the formation of sodamide NH_{2}Na, although it does not
displace hydrogen from the hydrocarbons.[45] Sodium burns both in
chlorine and in oxygen, evolving much heat. These properties are closely
connected with its power of taking up oxygen, chlorine, and similar
elements from most of their compounds. Just as it removes the oxygen
from the oxides of nitrogen and from carbonic anhydride, so also does
it decompose the majority of oxides at definite temperatures. Here the
action is essentially the same as in the decomposition of water. Thus,
for instance, when acting on magnesium chloride the sodium displaces the
magnesium, and when acting on aluminium chloride it displaces metallic
aluminium. Sulphur, phosphorus, arsenic and a whole series of other
elements, also combine with sodium.[46]

  [44] H. A. Schmidt remarked that perfectly dry hydrogen chloride is
       decomposed with great difficulty by sodium, although the
       decomposition proceeds easily with potassium and with sodium in
       moist hydrogen chloride. Wanklyn also remarked that sodium burns
       with great difficulty in dry chlorine. Probably these facts are
       related to other phenomena observed by Dixon, who found that
       perfectly dry carbonic oxide does not explode with oxygen on
       passing an electric spark.

  [44 bis] Sodamide, NH_{2}Na, (Chapter IV., Note 14), discovered by
       Gay-Lussac and Thénard, has formed the object of repeated
       research, but has been most fully investigated by A. W. Titherley
       (1894). Until recently the following was all that was known about
       this compound:--

       By heating sodium in dry ammonia, Gay-Lussac and Thénard obtained
       an olive-green, easily-fusible mass, _sodamide_, NH_{2}Na,
       hydrogen being separated. This substance with water forms sodium
       hydroxide and ammonia; with carbonic oxide, CO, it forms sodium
       cyanide, NaCN, and water, H_{2}O; and with dry hydrogen chloride
       it forms sodium and ammonium chlorides. These and other reactions
       of sodamide show that the metal in it preserves its energetic
       properties in reaction, and that this compound of sodium is more
       stable than the corresponding chlorine amide. When heated,
       sodamide, NH_{2}Na, only partially decomposes, with evolution of
       hydrogen, the principal part of it giving ammonia and sodium
       nitride, Na_{3}N, according to the equation 3NH_{2}Na = 2NH_{3} +
       NNa_{3}. The latter is an almost black powdery mass, decomposed by
       water into ammonia and sodium hydroxide.

       Titherley's researches added the following data:--

       Iron or silver vessels should be used in preparing this body,
       because glass and porcelain are corroded at 300°-400°, at which
       temperature ammonia gas acts upon sodium and forms the amide with
       the evolution of hydrogen. The reaction proceeds slowly, but is
       complete if there be an excess of NH_{3}. Pure NH_{2}Na is
       colourless (its colouration is due to various impurities),
       semi-transparent, shows traces of crystallisation, has a
       conchoidal fracture, and melts at 145°. Judging from the increase
       in weight of the sodium and the quantity of hydrogen which is
       disengaged, the composition of the amide is exactly NH_{2}Na. It
       partially volatilises (sublimes) in vacuo at 200°, and breaks up
       into 2Na + N_{2} + 2H_{2} at 500°. The same amide is formed when
       oxide of sodium is heated in NH_{3}: Na_{2}O + 2NH_{3} = 2NaH_{2}N
       + H_{2}O. NaHO is also formed to some extent by the resultant
       H_{2}O. Potassium and lithium form similar amides. With water,
       alcohol, and acids, NH_{2}Na gives NH_{3} and NaHO, which react
       further. Anhydrous CaO absorbs NH_{2}Na when heated without
       decomposing it. When sodamide is heated with SiO_{2}, NH_{3} is
       disengaged, and silicon nitride formed. It acts still more readily
       upon boric anhydride when heated with it: 2NH_{2}Na + B_{2}O_{3} =
       2BN + 2NaHO + H_{2}O. When slightly heated, NH_{2}Na + NOCl = NaCl
       + N_{2} + H_{2}O (NHNa_{2} and NNa_{3} are apparently not formed
       at a higher temperature). The halogen organic compounds react with
       the aid of heat, but with so much energy that the reaction
       frequently leads to the ultimate destruction of the organic groups
       and production of carbon.

  [45] As sodium does not displace hydrogen from the hydrocarbons, _it
       may be preserved_ in liquid hydrocarbons. Naphtha is generally
       used for this purpose, as it consists of a mixture of various
       liquid hydrocarbons. However, in naphtha sodium usually becomes
       coated with a crust composed of matter produced by the action of
       the sodium on certain of the substances contained in the mixture
       composing naphtha. In order that sodium may retain its lustre in
       naphtha, secondary octyl alcohol is added. (This alcohol is
       obtained by distilling castor oil with caustic potash.) Sodium
       keeps well in a mixture of pure benzene and paraffin.

  [46] If sodium does not directly displace the hydrogen in hydrocarbons,
       still by indirect means compounds may be obtained which contain
       sodium and hydrocarbon groups. Some of these compounds have been
       produced, although not in a pure state. Thus, for instance, zinc
       ethyl, Zn(C_{2}H_{5})_{2}, when treated with sodium, loses zinc
       and forms sodium ethyl, C_{2}H_{5}Na, but this decomposition is
       not complete, and the compound formed cannot be separated by
       distillation from the remaining zinc ethyl. In this compound the
       energy of the sodium is clearly manifest, for it reacts with
       substances containing haloids, oxygen, &c., and directly absorbs
       carbonic anhydride, forming a salt of a carboxylic acid
       (propionic).

With _oxygen_ sodium unites in three degrees of combination, forming
a suboxide Na_{4}O,[46 bis] an oxide, Na_{2}O, and a peroxide, NaO.
They are thus termed because Na_{2}O is a stable basic oxide (with
water it forms a basic hydroxide), whilst Na_{4}O and NaO do not
form corresponding saline hydrates and salts. The suboxide is a grey
inflammable substance which easily decomposes water, disengaging
hydrogen; it is formed by the slow oxidation of sodium at the ordinary
temperature. The peroxide is a greenish yellow substance, fusing at
a bright red heat; it is produced by burning sodium in an excess of
oxygen, and it yields oxygen when treated with water:

             Suboxide: Na_{4}O + 3H_{2}O = 4NaHO + H_{2}[47]
                Oxide: Na_{2}O + H_{2}O  = 2NaHO[48]
             Peroxide: Na_{2}O_{2} + H_{2}O = 2NaHO + O[49]

All three oxides form sodium hydroxide with water, but only the oxide
Na_{2}O is directly transformed into a hydrate. The other oxides liberate
either hydrogen or oxygen; they also present a similar distinction with
reference to many other agents. Thus carbonic anhydride combines directly
with the oxide Na_{2}O, which when heated in the gas burns, forming
sodium carbonate, whilst the peroxide yields oxygen in addition. When
treated with acids, sodium and all its oxides only form the salts
corresponding with sodium oxide--that is, of the formula or type NaX.
Thus the oxide of sodium, Na_{2}O, is _the only salt-forming oxide_ of
this metal, as water is in the case of hydrogen. Although the peroxide
H_{2}O_{2} is derived from hydrogen, and Na_{2}O_{2} from sodium, yet
there are no corresponding salts known, and if they are formed they are
probably as unstable as hydrogen peroxide. Although carbon forms carbonic
oxide, CO, still it has only one salt-forming oxide--carbonic anhydride,
CO_{2}. Nitrogen and chlorine both give several salt-forming oxides and
types of salts. But of the oxides of nitrogen, NO and NO_{2} do not form
salts, as do N_{2}O_{3}, N_{2}O_{4}, and N_{2}O_{5}, although N_{2}O_{4}
does not form special salts, and N_{2}O_{5} corresponds with the highest
form of the saline compounds of nitrogen. Such distinctions between the
elements, according to their power of giving one or several saline forms,
is a radical property of no less importance than the basic or acid
properties of their oxides. Sodium as a typical metal does not form any
acid oxides, whilst chlorine, as a typical non-metal, does not form bases
with oxygen. Therefore sodium _as an element_ may be thus characterised:
it forms one very stable salt-forming oxide, Na_{2}O, having powerful
basic properties, and its salts are of the general formula, NaX,
therefore in its compounds it is, like hydrogen, a basic and univalent
element.

  [46 bis] It is even doubtful whether the suboxide exists
           (_see_ Note 47).

  [47] A compound, Na_{2}Cl, which corresponds with the suboxide, is
       apparently formed when a galvanic current is passed through fused
       common salt; the sodium liberated dissolves in the common salt,
       and does not separate from the compound either on cooling or on
       treatment with mercury. It is therefore supposed to be Na_{2}Cl;
       the more so as the mass obtained gives hydrogen when treated with
       water: Na_{2}Cl + H_{2}O = H + NaHO + NaCl, that is, it acts like
       suboxide of sodium. If Na_{2}Cl really exists as a salt, then the
       corresponding base Na_{4}O, according to the rule with other bases
       of the composition M_{4}O, ought to be called a quaternary oxide.
       According to certain evidence, a suboxide is formed when thin
       sheets or fine drops of sodium slowly oxidise in moist air.

  [48] According to observations easily made, sodium when fused in air
       oxidises but does not burn, the combustion only commencing with
       the formation of vapour--that is, when considerably heated. Davy
       and Karsten obtained the oxides of potassium, K_{2}O, and of
       sodium, Na_{2}O, by heating the metals with their hydroxides,
       whence NaHO + Na = Na_{2}O + H, but N. N. Beketoff failed to
       obtain oxides by this means. He prepared them by directly igniting
       the metals in dry air, and afterwards heating with the metal in
       order to destroy any peroxide. The oxide produced, Na_{2}O, when
       heated in an atmosphere of hydrogen, gave a mixture of sodium and
       its hydroxide: Na_{2}O + H = NaHO + Na (_see_ Chapter II., Note
       9). If both the observations mentioned are accurate, then the
       reaction is reversible. Sodium oxide ought to be formed during the
       decomposition of sodium carbonate by oxide of iron (_see_ Note
       26), and during the decomposition of sodium nitrite. According to
       Karsten, its specific gravity is 2·8, according to Beketoff 2·3.
       The difficulty in obtaining it is owing to an excess of sodium
       forming the suboxide, and an excess of oxygen the peroxide. The
       grey colour peculiar to the suboxide and oxide perhaps shows that
       they contain metallic sodium. In addition to this, in the presence
       of water it may contain sodium hydride and NaHO.

  [49] Of the oxides of sodium, that easiest to form is the peroxide, NaO
       or Na_{2}O_{2}; this is obtained when sodium is burnt in an excess
       of oxygen. If NaNO_{3} be melted, it gives Na_{2}O_{2} with
       metallic Na. In a fused state the peroxide is reddish yellow, but
       it becomes almost colourless when cold. When heated with iodine
       vapour, it loses oxygen: Na_{2}O_{2} +I_{2} = Na_{2}OI_{2} + O.
       The compound Na_{2}OI_{2} is akin to the compound Cu_{2}OCl_{2}
       obtained by oxidising CuCl. This reaction is one of the few in
       which iodine directly displaces oxygen. The substance Na_{2}OI_{2}
       is soluble in water, and when acidified gives free iodine and a
       sodium salt. Carbonic oxide is absorbed by heated sodium peroxide
       with formation of sodium carbonate: Na_{2}CO_{3} = Na_{2}O_{2} +
       CO, whilst carbonic anhydride liberates oxygen from it. With
       nitrous oxide it reacts thus: Na_{2}O_{2} +2N_{2}O = 2NaNO_{2}
       +N_{2}; with nitric oxide it combines directly, forming sodium
       nitrite, NaO + NO = NaNO_{2}. Sodium peroxide, when treated with
       water, does not give hydrogen peroxide, because the latter in the
       presence of the alkali formed (Na_{2}O_{2}+ 2H_{2}O = 2NaHO +
       H_{2}O_{2}) decomposes into water and oxygen. In the presence of
       dilute sulphuric acid it forms H_{2}O_{2} (Na_{2}O_{2} +
       H_{2}SO_{4} = Na_{2}SO_{4} + H_{2}O_{2}). Peroxide of sodium is
       now prepared on a large scale (by the action of air upon Na at
       300°) for bleaching wool, silk &c. (when it acts in virtue of the
       H_{2}O_{2} formed). The oxidising properties of Na_{2}O_{2} under
       the action of heat are seen, for instance, in the fact that when
       heated with I it forms sodium iodate; with PbO, Na_{2}PbO_{3};
       with pyrites, sulphates, &c. When peroxide of sodium comes into
       contact with water, it evolves much heat, forming H_{2}O_{2}, and
       decomposing with the disengagement of oxygen; but, as a rule,
       there is no explosion. But if Na_{2}O_{2} be placed in contact
       with organic matter, such as sawdust, cotton, &c., it gives a
       violent explosion when heated, ignited, or acted on by water.
       Peroxide of sodium forms an excellent oxidising agent for the
       preparation of the higher product of oxidation of Mn, Cr, W, &c.,
       and also for oxidising the metallic sulphides. It should therefore
       find many applications in chemical analysis. To prepare
       Na_{2}O_{2} on a large scale, Castner melts Na in an aluminium
       vessel, and at 300° passes first air deprived of a portion of its
       oxygen (having been already once used), and then ordinary dry air
       over it.

On comparing sodium and its analogues, which will be described later
with other metallic elements, it will be seen that these properties,
together with the relative lightness of the metal itself and its
compounds, and the magnitude of its atomic weight comprise the most
essential properties of this element, clearly distinguishing it from
others, and enabling us easily to recognise its analogues.




                              CHAPTER XIII

       POTASSIUM, RUBIDIUM, CÆSIUM, AND LITHIUM. SPECTRUM ANALYSIS


Just as the series of halogens, fluorine, bromine and iodine correspond
with the chlorine contained in common salt, so also there exists a
corresponding series of elements: lithium, Li = 7, potassium, K = 39,
rubidium, Rb = 85, and cæsium, Cs = 133, which are analogous to the
sodium in common salt. These elements bear as great a resemblance to
sodium, Na = 23, as fluorine, F = 19, bromine, Br = 80, and iodine, I =
127, do to chlorine, Cl = 35·5. Indeed, in a free state, these elements,
like sodium, are soft metals which rapidly oxidise in moist air and
decompose water at the ordinary temperature, forming soluble hydroxides
having clearly-defined basic properties and the composition RHO, like
that of caustic soda. The resemblance between these metals is sometimes
seen with striking clearness, especially in compounds such as salts.[1]
The corresponding salts of nitric, sulphuric, carbonic, and nearly all
acids with these metals have many points in common. The metals which
resemble sodium so much in their reactions are termed the _metals of the
alkalis_.

  [1] Tutton's researches (1894) upon the analogy of the crystalline
       forms of K_{2}SO_{4}, Rb_{2}SO_{4} and Cs_{2}SO_{4} may be taken
      as a typical example of the comparison of analogous compounds. We
      cite the following data from these excellent researches: the sp.
      gr. at 20°/4° of K_{2}SO_{4} is 2·6633 of Rb_{2}SO_{4}, 3·6113, and
      of Cs_{2}SO_{4}, 4·2434. The coefficient of cubical expansion (the
      mean between 20° and 60°) for the K salt is 0·0053, for the Rb salt
      0·0052, for the Cs salt 0·0051. The linear expansion (the maximum
      for the vertical axis) along the axis of crystallisation is the
      same for all three salts, within the limits of experimental error.
      The replacement of potassium by rubidium causes the distance
      between the centres of the molecules in the direction of the three
      axes of crystallisation to increase equally, and less than with the
      replacement of rubidium by cæsium. The index of refraction for all
      rays and for every crystalline path (direction) is greater for the
      rubidium salt than for the potassium salt, and less than for the
      cæsium salt, and the differences are nearly in the ratio 2 : 5. The
      lengths of the rhombic crystalline axes for K_{2}SO_{4} are in the
      ratio 0·5727 : 1 : 0·7418, for Rb_{2}SO_{4}, 0·5723 : 1 : 0·7485,
      and for Cs_{2}SO_{4}, 0·5712 : 1 : 0·7521. The development of the
      basic and brachy-pinacoids gradually increases in passing from K to
      Rb and Cs. The optical properties also follow the same order both
      at the ordinary and at a higher temperature. Tutton draws the
      general conclusion that the crystallographic properties of the
      isomorphic rhombic sulphates R_{2}SO_{4} are a function of the
      atomic weight of the metals contained in them (_see_ Chapter XV.)
      Such researches as these should do much towards hastening the
      establishment of a true molecular mechanics of physico-chemical
      phenomena.

Among the metals of the alkalis, the most widely distributed in nature,
after sodium, is _potassium_. Like sodium, it does not appear either in
a free state or as oxide or hydroxide, but in the form of salts, which
present much in common with the salts of sodium in the manner of their
occurrence. The compounds of potassium and sodium in the earth's crust
occur as mineral compounds of silica. With silica, SiO_{2}, potassium
oxide, like sodium oxide, forms saline mineral substances resembling
glass. If other oxides, such as lime, CaO, and alumina, Al_{2}O_{3},
combine with these compounds, glass is formed, a vitreous stony mass,
distinguished by its great stability, and its very slight variation
under the action of water. It is such complex silicious compounds as
these which contain potash (potassium oxide), K_{2}O, or soda (sodium
oxide), Na_{2}O, and sometimes both together, silica, SiO_{2}, lime,
CaO, alumina, Al_{2}O_{3}, and other oxides, that form the chief mass
of rocks, out of which, judging by the direction of the strata, the
chief mass of the accessible crust (envelope) of the earth is made up.
The primary rocks, like granite, porphyry, &c.,[1 bis] are formed of
such crystalline silicious rocks as these. The oxides entering into the
composition of these rocks do not form a homogeneous amorphous mass
like glass, but are distributed in a series of peculiar, and in the
majority of cases crystalline, compounds, into which the primary rocks
may be divided. Thus a felspar (orthoclase) in granite contains from
8 to 15 per cent. of potassium, whilst another variety (plagioclase)
which also occurs in granite contains 1·2 to 6 per cent. of potassium,
and 6 to 12 per cent. of sodium. The mica in granite contains 3 to 10
per cent. of potassium. As already mentioned, and further explained in
Chapter XVII., the friable, crumbling, and stratified formations which
in our times cover a large part of the earth's surface have been formed
from these primary rocks by the action of the atmosphere and of water
containing carbonic acid. It is evident that in the chemical alteration
of the primary rocks by the action of water, the compounds of potassium,
as well as the compounds of sodium, must have been dissolved by the
water (as they are soluble in water), and that therefore the compounds
of potassium must be accumulated together with those of sodium in sea
water. And indeed compounds of potassium are always found in _sea
water_, as we have already pointed out (Chapters I. and X.). This forms
one of the sources from which they are extracted. After the evaporation
of sea water, there remains a mother liquor, which contains potassium
chloride and a large proportion of magnesium chloride. On cooling this
solution crystals separate out which contain chlorides of magnesium
and potassium. A double salt of this kind, called _carnallite_,
KMgCl_{3},6H_{2}O, occurs at Stassfurt. This carnallite[2] is now
employed as a material for the extraction of potassium chloride, and
of all the compounds of this element.[3] Besides which, potassium
chloride itself is sometimes found at Stassfurt as _sylvine_.[3 bis]
By a method of double saline decomposition, the chloride of potassium
may be converted into all the other potassium salts,[4] some of which
are of practical use. The potassium salts have, however, their greatest
importance as an indispensable component of the food of plants.[5]

  [1 bis] The origin of the primary rocks has been mentioned in
      Chapter X., Note 2.

  [2] Carnallite belongs to the number of double salts which are directly
      decomposed by water, and it only crystallises from solutions which
      contain an excess of magnesium chloride. It may be prepared
      artificially by mixing strong solutions of potassium and magnesium
      chlorides, when colourless crystals of sp. gr. 1·60 separate,
      whilst the Stassfurt salt is usually of a reddish tint, owing to
      traces of iron. At the ordinary temperature sixty-five parts of
      carnallite are soluble in one hundred parts of water in the
      presence of an excess of the salt. It deliquesces in the air,
      forming a solution of magnesium chloride and leaving potassium
      chloride. The quantity of carnallite produced at Stassfurt is now
      as much as 100,000 tons a year.

  [3] The method of separating sodium chloride from potassium chloride
      has been described in Chapter I. On evaporation of a mixture of the
      saturated solutions, sodium chloride separates; and then, on
      cooling, potassium chloride separates, owing to the difference of
      rate of variation of their solubilities with the temperature. The
      following are the most trustworthy figures for the solubility of
      _potassium chloride_ in one hundred parts of water (for sodium
      chloride, _see_ Chapter X., Note 13):--

                      10°   20°   40°   60°   100°
                      32    35    40    46     57

      When mixed with solutions of other salts the solubility of
      potassium chloride naturally varies, but not to any great extent.

  [3 bis] The specific gravity of the solid salt is 1·99--that is, less
      than that of sodium chloride. All the salts of sodium are
      specifically heavier than the corresponding salts of potassium, as
      are also their solutions for equal percentage compositions. If the
      specific gravity of water at 4° = 10,000, then at 15° the specific
      gravity of a solution of _p_ p.c. potassium chloride = 9,992 +
      63·29_p_ + 0·226_p_^2, and therefore for 10 p.c. = 1·0647, 20 p.c.
      = 1·1348, &c.

      Potassium chloride combines with iodine trichloride to form a
      compound KCl + ICl_{3} = KICl_{4}, which has a yellow colour, is
      fusible, loses iodine trichloride at a red heat, and gives
      potassium iodate and hydrochloric acid with water. It is not only
      formed by direct combination, but also by many other methods; for
      instance, by passing chlorine into a solution of potassium iodide
      so long as the gas is absorbed, KI + 2Cl_{2} = KCl,ICl_{3}.
      Potassium iodide, when treated with potassium chlorate and strong
      hydrochloric acid, also gives this compound; another method for its
      formation is given by the equation KClO_{3} + I + 6HCl =
      KCl,ICl_{3} + 3Cl + 3H_{2}O. This is a kind of salt corresponding
      with KIO_{2} (unknown) in which the oxygen is replaced by chlorine.
      If valency be taken as the starting-point in the study of chemical
      compounds, and the elements considered as having a constant
      atomicity (number of bonds)--that is, if K, Cl, and I be taken as
      univalent elements--then it is impossible to explain the formation
      of such a compound because, according to this view, univalent
      elements are only able to form dual compounds with each other; such
      as, KCl, ClI, KI, &c., whilst here they are grouped together in the
      molecule KICl_{4}. Wells, Wheeler, and Penfield (1892) obtained a
      large number of such poly-haloid salts. They may all be divided
      into two large classes: the tri-haloid and the penta-haloid salts.
      They have been obtained not only for K but also for Rb and Cs, and
      partially also for Na and Li. The general method of their formation
      consists in dissolving the ordinary halogen salt of the metal in
      water, and treating it with the requisite amount of free halogen.
      The poly-haloid salt separates out after evaporating the solution
      at a more or less low temperature. In this manner, among the
      tri-haloid salts, may be obtained: KI_{3}, KBr_{2}I, KCl_{2}I, and
      the corresponding salts of rubidium and cæsium, for instance,
      CsI_{3}, CsBrI_{2}, CsBr_{2}I, CsClBrI,CsCl_{2}I, CsBr_{3},
      CsClBr_{2}, CsCl_{2}Br, and in general MX_{3} where X is a halogen.
      The colour of the crystals varies according to the halogen, thus
      CsI_{3} is black, CrBr_{3} yellowish red, CrBrI_{2} reddish brown,
      CsBr_{2}I red, CsCl_{2}Br yellow. The cæsium salts are the most
      stable, and those of potassium least so, as also those which
      contain Br and I separately or together; for cæsium no compounds
      containing Cl and I were obtained. The penta-haloid salts form a
      smaller class; among these salts potassium forms KCl_{4}I, rubidium
      RbCl_{4}I, cæsium CsI_{5}, CsBr, CsCl_{4}I, lithium LiCl_{4}I (with
      4H_{2}O) and sodium NaCl_{4}I (with 2H_{2}O). The most stable are
      those salts containing the metal with the greatest atomic
      weight--cæsium (_see_ Chapter XI., Note 63).

  [4] It is possible to extract the compounds of potassium directly from
      the primary rocks which are so widely distributed over the earth's
      surface and so abundant in some localities. From a chemical point
      of view this problem presents no difficulty; for instance, by
      fusing powdered orthoclase with lime and fluor spar (Ward's method)
      and then extracting the alkali with water (on fusion the silica
      gives an insoluble compound with lime), or by treating the
      orthoclase with hydrofluoric acid (in which case silicon fluoride
      is evolved as a gas) it is possible to transfer the alkali of the
      orthoclase to an aqueous solution, and to separate it in this
      manner from the other insoluble oxides. However, as yet there is no
      profit in, nor necessity for, recourse to this treatment, as
      carnallite and potash form abundant materials for the extraction of
      potassium compounds by cheaper methods. Furthermore, the salts of
      potassium are now in the majority of chemical reactions replaced by
      salts of sodium, especially since the preparation of sodium
      carbonate has been facilitated by the Leblanc process. The
      replacement of potassium compounds by sodium compounds not only has
      the advantage that the salts of sodium are in general cheaper than
      those of potassium, but also that a smaller quantity of a sodium
      salt is needed for a given reaction than of a potassium salt,
      because the combining weight of sodium (23) is less than that of
      potassium (39).

 [5] It has been shown by direct experiment on the cultivation of plants
      in artificial soils and in solutions that under conditions
      (physical, chemical, and physiological) otherwise identical plants
      are able to thrive and become fully developed in the entire absence
      of sodium salts, but that their development is impossible without
      potassium salts.

The primary rocks contain an almost equal proportion of potassium and
sodium. But in sea water the compounds of the latter metal predominate.
It may be asked, what became of the compounds of potassium in the
disintegration of the primary rocks, if so small a quantity went to the
sea water? They remained with the other products of the decomposition
of the primary rocks. When granite or any other similar rock formation
is disintegrated, there are formed, besides the soluble substances,
also insoluble substances--sand and finely-divided clay, containing
water, alumina, and silica. This clay is carried away by the water,
and is then deposited in strata. It, and especially its admixture with
vegetable remains, retain compounds of potassium in a greater quantity
than those of sodium. This has been proved with absolute certainty
to be the case, and is due to the _absorptive power of the soil_. If
a dilute solution of a potassium compound be filtered through common
mould used for growing plants, containing clay and the remains of
vegetable decomposition, this mould will be found to have retained a
somewhat considerable percentage of the potassium compounds. If a salt
of potassium be taken, then during the filtration an equivalent quantity
of a salt of calcium--which is also found, as a rule, in soils--is
set free. Such a process of filtration through finely divided earthy
substances proceeds in nature, and the compounds of potassium are
everywhere retained by the friable earth in considerable quantity. This
explains the presence of so small an amount of potassium salts in the
water of rivers, lakes, streams, and oceans, where the lime and soda
have accumulated. The compounds of potassium retained by the friable
mass of the earth are absorbed as an aqueous solution by the roots
of _plants_. Plants, as everyone knows, when burnt leave an ash, and
this ash, besides various other substances, without exception contains
compounds of potassium. Many land plants contain a very small amount of
sodium compounds,[6] whilst potassium and its compounds occur in all
kinds of vegetable ash. Among the generally cultivated plants, grass,
potatoes, the turnip, and buckwheat are particularly rich in potassium
compounds. The ash of plants, and especially of herbaceous plants,
buckwheat straw, sunflower and potato leaves are used in practice for
the extraction of potassium compounds. There is no doubt that potassium
occurs in the plants themselves in the form of complex compounds,
and often as salts of organic acids. In certain cases such salts of
potassium are even extracted from the juice of plants. Thus, sorrel
and oxalis, for example, contain in their juices the acid oxalate of
potassium, C_{2}HKO_{4}, which is employed for removing ink stains.
Grape juice contains the so-called cream of tartar, which is the acid
tartrate of potassium, C_{4}H_{5}KO_{6}.[7] This salt also separates
as a sediment from wine. When the plants, containing one or more of the
salts of potassium, are burnt, the carbonaceous matter is oxidised,
and in consequence the potassium is obtained in the ash as carbonate,
K_{2}CO_{3}, which is generally known as _potashes_. Hence potashes
occur ready prepared in the ash of plants, and therefore the ash of
land plants is employed as a source for the extraction of potassium
compounds. Potassium carbonate is extracted by lixiviating the ash with
water.[8] Potassium carbonate may also be obtained from the chloride
by a method similar to that by which sodium carbonate is prepared from
sodium chloride.[8 bis] There is no difficulty in obtaining any salt of
potassium--for example, the sulphate,[9] bromide, and iodide[10]--by
the action of the corresponding acid on KCl and especially on the
carbonate, whilst the hydroxide, _caustic potash_, KHO, which is in many
respects analogous to caustic soda, is easily obtained by means of lime
in exactly the same manner in which sodium hydroxide is prepared from
sodium carbonate.[11] Therefore, in order to complete our knowledge of
the alkali metals, we will only describe two salts of potassium which
are of practical importance, and whose analogues have not been described
in the preceding chapter, potassium cyanide and potassium nitrate.

  [6] If herbaceous plants contain much sodium salts, it is evident that
      these salts mainly come from the sodium compounds in the water
      absorbed by the plants.

  [7] As plants always contain mineral substances and cannot thrive in
      a medium which does not contain them, more especially in one which
      is free from the salts of the four basic oxides, K_{2}O, CaO, MgO,
      and Fe_{2}O_{3}, and of the four acid oxides, CO_{2}, N_{2}O_{5},
      P_{2}O_{5}, and SO_{3}, and as the amount of ash-forming substances
      in plants is small, the question inevitably arises as to what part
      these play in the development of plants. With the existing chemical
      data only one answer is possible to this question, and it is still
      only a hypothesis. This answer was particularly clearly expressed
      by Professor Gustavson of the Petroffsky Agricultural Academy.
      Starting from the fact (Chapter XI., Note 55) that a small quantity
      of aluminium renders possible or facilitates the reaction of
      bromine on hydrocarbons at the ordinary temperature, it is easy to
      arrive at the conclusion, which is very probable and in accordance
      with many data respecting the reactions of organic compounds, that
      the addition of mineral substances to organic compounds lowers the
      temperature of reaction and in general facilitates chemical
      reactions in plants, and thus aids the conversion of the most
      simple nourishing substances into the complex component parts of
      the plant organism. The province of chemical reactions proceeding
      in organic substances in the presence of a small quantity of
      mineral substances has as yet been but little investigated,
      although there are already several disconnected data concerning
      reactions of this kind, and although a great deal is known with
      regard to such reactions among inorganic compounds. The essence of
      the matter may be expressed thus--two substances, A and B, do not
      react on each other of their own accord, but the addition of a
      small quantity of a third particularly active substance, C,
      produces the reaction of A on B, because A combines with C, forming
      AC, and B reacts on this new compound, which has a different store
      of chemical energy, forming the compound AB or its products, and
      setting C free again or retaining it.

      It may here be remarked that all the mineral substances necessary
      for plants (those enumerated at the beginning of the note) are the
      highest saline compounds of their elements, that they enter into
      the plants as salts, that the lower forms of oxidation of the same
      elements (for instance, sulphites and phosphites) are harmful to
      plants (poisonous), and that strong solutions of the salts
      assimilated by plants (their osmotic pressure being great and
      contracting the cells, as De Vries showed, (_see_ Chapter I., Note
      19)) not only do not enter into the plants but kill them (poison
      them).

  [8] Besides which, it will be understood from the preceding paragraph
      that the salts of potassium may become exhausted from the soil by
      long cultivation, and that there may therefore be cases when the
      direct fertilisation by salts of potassium may be profitable. But
      manure and animal excrements, ashes, and, in general, nearly all
      refuse which may serve for fertilising the soil, contain a
      considerable quantity of potassium salts, and therefore, as regards
      the natural salts of potassium (Stassfurt), and especially
      potassium sulphate, if they often improve the crops, it is in all
      probability due to their action on the properties of the soil. The
      agriculturist cannot therefore be advised to add potassium salts,
      without making special experiments showing the advantage of such a
      fertiliser on a given kind of soil and plant.

      The animal body also contains potassium compounds, which is
      natural, since animals consume plants. For example, milk, and
      especially human milk, contains a somewhat considerable quantity of
      potassium compounds. Cow's milk, however, does not contain much
      potassium salt. Sodium compounds generally predominate in the
      bodies of animals. The excrement of animals, and especially of
      herbivorous animals, on the contrary, often contains a large
      proportion of potassium salts. Thus sheep's dung is rich in them,
      and in washing sheep's wool salts of potassium pass into the water.

      The ash of tree stems, as the already dormant portion of the plant
      (Chapter VIII., Note 1), contains little potash. For the extraction
      of potash, which was formerly carried on extensively in the east of
      Russia (before the discovery of the Stassfurt salt), the ash of
      grasses, and the green portions of potatoes, buckwheat, &c., are
      taken and treated with water (lixiviated), the solution is
      evaporated, and the residue ignited in order to destroy the organic
      matter present in the extract. The residue thus obtained is
      composed of raw potash. It is refined by a second dissolution in a
      small quantity of water, for the potash itself is very soluble in
      water, whilst the impurities are sparingly soluble. The solution
      thus obtained is again evaporated, and the residue ignited, and
      this potash is then called refined potash, or pearlash. This method
      of treatment cannot give chemically pure potassium carbonate. A
      certain amount of impurities remain. To obtain chemically pure
      potassium carbonate, some other salt of potassium is generally
      taken and purified by crystallisation. Potassium carbonate
      crystallises with difficulty, and it cannot therefore be purified
      by this means, whilst other salts, such as the tartrate, acid
      carbonate, sulphate, or nitrate, &c., crystallise easily and may
      thus be directly purified. The tartrate is most frequently
      employed, since it is prepared in large quantities (as a sediment
      from wine) for medicinal use under the name of cream of tartar.
      When ignited without access of air, it leaves a mixture of charcoal
      and potassium carbonate. The charcoal so obtained being in a
      finely-divided condition, the mixture (called 'black flux'), is
      sometimes used for reducing metals from their oxides with the aid
      of heat. A certain quantity of nitre is added to burn the charcoal
      formed by heating the cream of tartar. Potassium carbonate thus
      prepared is further purified by converting it into the acid salt,
      by passing a current of carbonic anhydride through a strong
      solution. KHCO_{3} is then formed, which is less soluble than the
      normal salt (as is also the case with the corresponding sodium
      salts), and therefore crystals of the acid salt separate from the
      solution on cooling. When ignited, they part with their water and
      carbonic anhydride, and pure potassium carbonate remains behind.
      The physical properties of potassium carbonate distinguish it
      sufficiently from sodium carbonate; it is obtained from solutions
      as a powdery white mass, having an alkaline taste and reaction,
      and, as a rule, shows only traces of crystallisation. It also
      attracts the moisture of the air with great energy. The crystals do
      not contain water, but absorb it from the air, deliquescing into a
      saturated solution. It melts at a red heat (1045°), and at a still
      higher temperature is even converted into vapour, as has been
      observed at glass works where it is employed. It is very soluble.
      At the ordinary temperature, water dissolves an equal weight of the
      salt. Crystals containing two equivalents of water separate from
      such a saturated solution when strongly cooled (Morel obtained
      K_{2}CO_{3}3H_{2}O in well-formed crystals at +10°). There is no
      necessity to describe its reactions, because they are all analogous
      to those of sodium carbonate. When manufactured sodium carbonate
      was but little known, the consumption of potassium carbonate was
      very considerable, and even now washing soda is frequently replaced
      for household purposes by 'lye'--_i.e._ an aqueous solution
      obtained from ashes. It contains potassium carbonate, which acts
      like the sodium salt in washing tissues, linen, &c.

      A mixture of potassium and sodium carbonates fuses with much
      greater ease than the separate salts, and a mixture of their
      solutions gives well-crystallised salts--for instance (Marguerite's
      salt), K_{2}CO_{3},6H_{2}O,2Na_{2}CO_{3},6H_{2}O. Crystallisation
      also occurs in other multiple proportions of K and Na (in the above
      case 1 : 2, but 1 : 1 and 1 : 3 are known), and always with 6 mol.
      H_{2}O. This is evidently a combination _by similarity_, as in
      alloys, solutions, &c.

  [8 bis] About 25,000 tons of potash annually are now prepared from KCl
      by this method at Stassfurt.

  [9] _Potassium sulphate_, K_{2}SO_{4}, crystallises from its solutions
      in an anhydrous condition, in which respect it differs from the
      corresponding sodium salt, just as potassium carbonate differs from
      sodium carbonate. In general, it must be observed that the majority
      of sodium salts combine more easily with water of crystallisation
      than the potassium salts. The solubility of _potassium sulphate_
      does not show the same peculiarities as that of sodium sulphate,
      because it does not combine with water of crystallisation; at the
      ordinary temperature 100 parts of water dissolve about 10 parts of
      the salt, at 0° 8·3 parts, and at 100° about 26 parts. _The acid
      sulphate_, KHSO_{4}, obtained easily by heating crystals of the
      normal salt with sulphuric acid, is frequently employed in chemical
      practice. On heating the mixture of acid and salt, fumes of
      sulphuric acid are at first given off; when they cease to be
      evolved, the acid salt is contained in the residue. At a higher
      temperature (of above 600°) the acid salt parts with all the acid
      contained in it, the normal salt being re-formed. The definite
      composition of this acid salt, and the ease with which it
      decomposes, render it exceedingly valuable for certain chemical
      transformations accomplished by means of sulphuric acid at a high
      temperature, because it is possible to take, in the form of this
      salt, a strictly definite quantity of sulphuric acid, and to cause
      it to act on a given substance at a high temperature, which it is
      often necessary to do, more especially in chemical analysis. In
      this case, the acid salt acts in exactly the same manner as
      sulphuric acid itself, but the latter is inefficient at
      temperatures above 400°, because it all evaporates, while at that
      temperature the acid salt still remains in a fused state, and acts
      with the elements of sulphuric acid on the substance taken. Hence
      by its means the boiling-point of sulphuric acid is raised. Thus
      the acid potassium sulphate is employed, where for conversion of
      certain oxides, such as those of iron, aluminium, and chromium,
      into salts, a high temperature is required.

      Weber, by heating potassium sulphate with an excess of sulphuric
      acid at 100°, observed the formation of a lower stratum, which was
      found to contain a definite compound containing eight equivalents
      of SO_{3} per equivalent of K_{2}O. The salts of rubidium, cæsium,
      and thallium give a similar result, but those of sodium and lithium
      do not. (_See_ Note 1.)

  [10] The _bromide_ and _iodide_ of potassium are used, like the
       corresponding sodium compounds, in medicine and photography.
       Potassium iodide is easily obtained in a pure state by saturating
       a solution of hydriodic acid with caustic potash. In practice,
       however, this method is rarely had recourse to, other more simple
       processes being employed although they do not give so pure a
       product. They aim at the direct formation of hydriodic acid in the
       liquid in the presence of potassium hydroxide or carbonate. Thus
       iodine is thrown into a solution of pure potash, and hydrogen
       sulphide passed through the mixture, the iodine being thus
       converted into hydriodic acid. Or a solution is prepared from
       phosphorus, iodine, and water, containing hydriodic and phosphoric
       acid; lime is then added to this solution, when calcium iodide is
       obtained in solution, and calcium phosphate as a precipitate. The
       solution of calcium iodide gives, with potassium carbonate,
       insoluble calcium carbonate and a solution of potassium iodide. If
       iodine is added to a slightly-heated solution of caustic potash
       (free from carbonate--that is, freshly prepared), so long as the
       solution is not  from the presence of an excess of iodine,
       there is formed (as in the action of chlorine on a solution of
       caustic potash) a mixture of potassium iodide and iodate. On
       evaporating the solution thus obtained and igniting the residue,
       the iodate is destroyed and converted into iodide, the oxygen
       being disengaged, and potassium iodide only is left behind. On
       dissolving the residue in water and then evaporating, cubical
       crystals of the anhydrous salt are obtained, which are soluble in
       water and alcohol, and on fusion give an alkaline reaction, owing
       to the fact that when ignited a portion of the salt decomposes,
       forming potassium oxide. The neutral salt may be obtained by
       adding hydriodic acid to this alkaline salt until it gives an acid
       reaction. It is best to add some finely-divided charcoal to the
       mixture of iodate and iodide before igniting it, as this
       facilitates the evolution of the oxygen from the iodate. The
       iodate may also be converted into iodide by the action of certain
       reducing agents, such as zinc amalgam, which when boiled with a
       solution containing an iodate converts it into iodide. Potassium
       iodide may also be prepared by mixing a solution of ferrous iodide
       (it is best if the solution contain an excess of iodine) and
       potassium carbonate, in which case ferrous carbonate FeCO_{3}, is
       precipitated (with an excess of iodine the precipitate is
       granular, and contains a compound of the suboxide and oxide of
       iron), while potassium iodide remains in solution. Ferrous iodide,
       FeI_{2}, is obtained by the direct action of iodine on iron in
       water. Potassium iodide considerably lowers the temperature (by
       24°), when it dissolves in water, 100 parts of the salt dissolve
       in 73·5 parts of water at 12·5°, in 70 parts at 18°, whilst the
       saturated solution which boils at 120° contains 100 parts of salt
       per 45 parts of water. Solutions of potassium iodide dissolve a
       considerable amount of iodine; strong solutions even dissolving as
       much or more iodine than they contain as potassium iodide (_see_
       Note 3 bis and Chapter XI., Note 64).

  [11] Caustic potash is not only formed by the action of lime on dilute
       solutions of potassium carbonate (as sodium hydroxide is prepared
       from sodium carbonate), but by igniting potassium nitrate with
       finely-divided copper (_see_ Note 15), and also by mixing
       solutions of potassium sulphate (or even of alum, KAlS_{2}O_{8})
       and barium hydroxide, BaH_{2}O_{2}. It is sometimes purified by
       dissolving it in alcohol (the impurities, for example, potassium
       sulphate and carbonate, are not dissolved) and then evaporating
       the alcohol.

       The specific gravity of potassium hydroxide is 2·04, but that of
       its solutions (see Chapter XII., Note 18) at 15° S = 9,992 +
       90·4_p_ + 0·28_p_^2 (here _p_^2 is +, and for sodium hydroxide it
       is-). Strong solutions, when cooled, yield a crystallo-hydrate,
       KHO,4H_{2}O, which dissolves in water, producing cold (like
       2NaHO,7H_{2}O), whilst potassium hydroxide in solution develops a
       considerable amount of heat.

_Potassium cyanide_, which presents in its chemical relations a certain
analogy with the halogen salts of potassium, is not only formed
according to the equation, KHO + HCN = H_{2}O + KCN, but also whenever
a nitrogenous carbon compound--for instance, animal matter--is heated
in the presence of metallic potassium, or of a compound of potassium,
and even when a mixture of potash and carbon is heated in a stream of
nitrogen. Potassium cyanide is obtained from yellow prussiate, which
has been already mentioned in Chapter IX., and whose preparation
on a large scale will be described in Chapter XXII. If the yellow
prussiate be ground to a powder and dried, so that it loses its water
of crystallisation, it then melts at a red heat, and decomposes into
carbide of iron, nitrogen, and potassium cyanide, FeK_{4}C_{6}N_{6} =
4KCN + FeC_{2} + N_{2}. After the decomposition it is found that the
yellow salt has been converted into a white mass of potassium cyanide.
The carbide of iron formed collects at the bottom of the vessel. If
the mass thus obtained be treated with water, the potassium cyanide is
partially decomposed by the water, but if it be treated with alcohol,
then the cyanide is dissolved, and on cooling separates in a crystalline
form.[12] A solution of potassium cyanide has a powerfully alkaline
reaction, a smell like that of bitter almonds, peculiar to prussic acid,
and acts as a most powerful poison. Although exceedingly stable in a
fused state, potassium cyanide easily changes when in solution. Prussic
acid is so very feebly energetic that even water decomposes potassium
cyanide. A solution of the salt, even without access of air, easily
turns brown and decomposes, and when heated evolves ammonia and forms
potassium formate; this is easily comprehensible from the representation
of the cyanogen compounds which was developed in Chapter IX., KCN +
2H_{2}O = CHKO_{2} + NH_{3}. Furthermore, as carbonic anhydride acts on
potassium cyanide with evolution of prussic acid, and as potassium
cyanate, which is also unstable, is formed by the action of air, it will
be easily seen that solutions of potassium cyanide are very unstable.
Potassium cyanide, containing as it does carbon and potassium, is a
substance which can act in a very vigorously reducing manner, especially
when fused; it is therefore used as a powerful reducing agent at a red
heat.[13] The property of potassium cyanide of giving double salts
with other cyanides is very clearly shown by the fact that many metals
dissolve in a solution of potassium cyanide, with the evolution of
hydrogen. For example, iron, copper, and zinc act in this manner. Thus--

          4KCN + 2H_{2}O + Zn = K_{2}ZnC_{4}N_{4} + 2KHO + H_{2}

  [12] When the yellow prussiate is heated to redness, all the cyanogen
       which was in combination with the iron is decomposed into
       nitrogen, which is evolved as gas, and carbon, which combines with
       the iron. In order to avoid this, potassium carbonate is added to
       the yellow prussiate while it is being fused. A mixture of 8 parts
       of anhydrous yellow prussiate and 3 parts of pure potassium
       carbonate is generally taken. Double decomposition then takes
       place, resulting in the formation of ferrous carbonate and
       potassium cyanide. But by this method, as by the first, a pure
       salt is not obtained, because a portion of the potassium cyanide
       is oxidised at the expense of the iron carbonate and forms
       potassium cyanate, FeCO_{3} + KCN = CO_{2} + Fe + KCNO; and the
       potassium cyanide very easily forms oxide, which acts on the sides
       of the vessel in which the mixture is heated (to avoid this iron
       vessels should be used). By adding one part of charcoal powder to
       the mixture of 8 parts of anhydrous yellow prussiate and 3 parts
       of potassium carbonate a mass is obtained which is free from
       cyanate, because the carbon absorbs the oxygen, but in that case
       it is impossible to obtain a colourless potassium cyanide by
       simple fusion, although this may be easily done by dissolving it
       in alcohol. Cyanide of potassium may also be obtained from
       potassium thiocyanate, which is formed from ammonium thiocyanate
       obtained by the action of ammonia upon bisulphide of carbon (_see_
       works upon Organic Chemistry). Potassium cyanide is now prepared
       in large quantities from yellow prussiate for gilding and
       silvering. When fused in large quantities the action of the oxygen
       of the air is limited, and with great care the operation may be
       successfully conducted, and therefore, on a large scale, very pure
       salt is sometimes obtained. When slowly cooled, the fused salt
       separates in cubical crystals like potassium chloride.

       Pure KCN is obtained by passing CNH gas into an alcoholic solution
       of KHO. The large amount of potassium cyanide which is now
       required for the extraction of gold from its ores, is being
       replaced by a mixture (Rossler and Gasslaker, 1892) of KCN and
       NaCN, prepared by heating powdered and dried yellow prussiate with
       metallic sodium: K_{4}Fe(CN)_{6} + 2Na = 4KCN + 2NaCN + Fe. This
       method offers two advantages over the above methods: (1) the whole
       of the cyanide is obtained, and does not decompose with the
       formation of N_{2}; and (2) no cyanates are formed, as is the case
       when carbonate of potash is heated with the prussiate.

  [13] A considerable quantity of potassium cyanide is used in the arts,
       more particularly for the preparation of metallic solutions which
       are decomposed by the action of a galvanic current; thus it is
       very frequently employed in electro-silvering and gilding. An
       alkaline solution is prepared, which is moderately stable owing to
       the fact that potassium cyanide in the form of certain double
       salts--that is, combined with other cyanides--is far more stable
       than when alone (yellow prussiate, which contains potassium
       cyanide in combination with ferrous cyanide, is an example of
       this).

Gold and silver are soluble in potassium cyanide in the presence of air,
in which case the hydrogen, which would otherwise be evolved in the
reaction, combines with the oxygen of the air, forming water (Eissler,
MacLaurin, 1893), for example, 4Au + 4KCN + O + H_{2}O = 2AuKC_{2}N_{2}
+ 2KHO, which is taken advantage of for extracting gold from its ores
(Chapter XXIV.).[13 bis] Platinum, mercury, and tin are not dissolved in
a solution of potassium cyanide, even with access of air.

  [13 bis] A dilute solution of KCN is taken, not containing more than
       1 per cent. KCN. MacLaurin explains this by the fact that strong
       solutions dissolve gold less rapidly, owing to their dissolving
       less air, whose oxygen is necessary for the reaction.

_Potassium nitrate_, or common _nitre_ or _saltpetre_, KNO_{3}, is
chiefly used as a component part of gunpowder, in which it cannot be
replaced by the sodium salt, because the latter is deliquescent. It
is necessary that the nitre in gunpowder should be perfectly pure, as
even small traces of sodium, magnesium, and calcium salts, especially
chlorides, render the nitre and the gunpowder capable of attracting
moisture. Nitre may easily be obtained pure, owing to its great
disposition to form crystals both large and small, which aids its
separation from other salts. The considerable differences between the
solubility of nitre at different temperatures aids this crystallisation.
A solution of nitre saturated at its boiling point (116°) contains
335 parts of nitre to 100 parts of water, whilst at the ordinary
temperature--for instance, 20°--the solution is only able to retain 32
parts of the salt. Therefore, in the preparation and refining of nitre,
its solution, saturated at the boiling point, is cooled, and nearly
all the nitre is obtained in the form of crystals. If the solution be
quietly and slowly cooled in large quantities then large crystals are
formed, but if it be rapidly cooled and agitated then small crystals are
obtained. In this manner, if not all, at all events the majority, of
the impurities present in small quantities remain in the mother liquor.
If an unsaturated solution of nitre be rapidly cooled, so as to prevent
the formation of large crystals (in whose crevices the mother liquor,
together with the impurities, would remain), the very minute crystals of
nitre known as saltpetre flour are obtained.

Common nitre occurs in nature, but only in small quantities in admixture
with other nitrates, and especially with sodium, magnesium, and calcium
nitrates. Such a mixture of salts of nitric acid is formed in nature
in fertile earth, and in those localities where, as in _the soil_,
nitrogenous organic remains are decomposed in the presence of alkalis or
alkaline bases with free access of air. This method of the formation of
nitrates requires moisture, besides the free access of air, and takes
place principally during warm weather.[14] In warm countries, and in
temperate climates during the summer months, fertile soils produce a
small quantity of nitre. In this respect India is especially known as
affording a considerable supply of nitre extracted from the soil. The
nitre-bearing soil after the rainy season sometimes becomes covered
during the summer with crystals of nitre, formed by the evaporation of
the water in which it was previously dissolved. This soil is collected,
subjected to repeated lixiviations, and treated for nitre as will be
presently described. In temperate climates nitrates are obtained from
the lime rubbish of demolished buildings which have stood for many
years, and especially from those portions which have been in contact
with the ground. The conditions there are very favourable for the
formation of nitre, because the lime used as a cement in buildings
contains the base necessary for the formation of nitrates, while the
excrement, urine, and animal refuse are sources of nitrogen. By the
methodical lixiviation of this kind of rubbish a solution of nitrogenous
salts is formed similar to that obtained by the lixiviation of fertile
soil. A similar solution is also obtained by the lixiviation of the
so-called _nitre plantations_. They are composed of manure interlaid
with brushwood, and strewn over with ashes, lime, and other alkaline
rubbish. These nitre plantations are set up in those localities where
the manure is not required for the fertilisation of the soil, as, for
example, in the south-eastern 'black earth' Governments of Russia.
The same process of oxidation of nitrogenous matter freely exposed to
air and moisture during the warm season in the presence of alkalis
takes place in nitre plantations as in fertile soil and in the walls
of buildings. From all these sources there is obtained a solution
containing various salts of nitric acid mixed with soluble organic
matter. The simplest method of treating this impure solution of nitre
is to add a solution of potassium carbonate, or to simply treat it with
ashes containing this substance. The potassium carbonate enters into
double decomposition with the calcium and magnesium salts, forming
insoluble carbonates of these bases and leaving the nitre in solution.
Thus, for instance, K_{2}CO_{3} + Ca(NO_{3})_{2} = 2KNO_3 + CaCO_{3}.
Both calcium and magnesium carbonates are insoluble, and therefore after
treatment with potassium carbonate the solution no longer contains salts
of these metals but only the salts of sodium and potassium together with
organic matter. The latter partially separates on heating in an insoluble
form, and is entirely destroyed by heating the nitre to a low red heat.
The nitre thus obtained is easily purified by repeated crystallisation.
The greater part of the nitre used for making gunpowder is now obtained
from the sodium salt _Chili saltpetre_ or _cubic nitre_, which occurs in
nature, as already mentioned. The conversion of this salt into common
nitre is also carried on by means of a double decomposition. This is done
either by adding potassium carbonate (when, on mixing the strong and hot
solutions, sodium carbonate is directly obtained as a precipitate), or,
as is now most frequent, potassium chloride. When a mixture of strong
solutions of potassium chloride and sodium nitrate is evaporated, sodium
chloride first separates, because this salt, which is formed by the
double decomposition KCl + NaNO_{3} = KNO_{3} + NaCl, is almost equally
soluble in hot and cold water; on cooling, therefore, a large amount of
potassium nitrate separates from the saturated solution, while the sodium
chloride remains dissolved. The nitre is ultimately purified by
recrystallisation and by washing with a saturated solution of nitre,
which cannot dissolve a further quantity of nitre but only the
impurities.

  [14] Besides which Schloesing and Müntz, by employing similar methods
       to Pasteur, showed that the formation of nitre in the
       decomposition of nitrogenous substances is accomplished by the aid
       of peculiar micro-organisms (ferments), without which the
       simultaneous action of the other necessary conditions (alkalis,
       moisture, a temperature of 37°, air, and nitrogenous substances)
       cannot give nitre.

Nitre is a colourless salt having a peculiar cool taste. It crystallises
easily in long striated six-sided rhombic prisms terminating in rhombic
pyramids. Its crystals (sp. gr. 1·93) do not contain water, but their
cavities generally contain a certain quantity of the solution from
which they have crystallised. For this reason in refining nitre,
the production of large crystals is prevented, _saltpetre flour_
being prepared. At a low red heat (339°) nitre melts to a colourless
liquid.[14 bis] Potassium nitrate at the ordinary temperature and in a
solid form is inactive and stable, but _at a high temperature_ it acts
as a powerful _oxidising agent_, giving up a considerable amount of
oxygen to substances mixed with it.[15] When thrown on to incandescent
charcoal it brings about its rapid combustion, and a mechanical mixture
of powdered charcoal and nitre ignites when brought into contact
with a red-hot substance, and continues to burn by itself. In this
action, nitrogen is evolved, and the oxygen oxidises the charcoal, in
consequence of which potassium carbonate and carbonic anhydride are
formed: 4KNO_{3} + 5C = 2K_{2}CO_{3} + 3CO_{2} + 2N_{2}. This phenomenon
depends on the fact that oxygen in combining with carbon evolves more
heat than it does in combining with nitrogen. Hence, when once the
combustion has been started at the expense of the nitre, it is able to
go on without requiring the aid of external heat. A similar oxidation
or combustion at the expense of the contained oxygen takes place when
nitre is heated with different combustible substances. If a mixture of
sulphur and nitre be thrown upon a red-hot surface, the sulphur burns,
forming potassium sulphate and sulphurous anhydride. In this case, also,
the nitrogen of the nitre is evolved as gas: 2KNO_{3} + 2S = K_{2}SO_{4}
+ N_{2} + SO_{2}. A similar phenomenon occurs when nitre is heated with
many metals. The oxidation of those metals which are able to form acid
oxides with an excess of oxygen is especially remarkable. In this case
they remain in combination with potassium oxide as potassium salts.
Manganese, antimony, arsenic, iron, chromium, &c. are instances of this
kind. These elements, like carbon and sulphur, displace free nitrogen.
The lower oxides of these metals when fused with nitre pass into the
higher oxides. Organic substances are also oxidised when heated with
nitre--that is, they burn at the expense of the nitre. It will be
readily understood from this that nitre is frequently used in practical
chemistry and the arts as an oxidising agent at high temperatures.
Its application in _gunpowder_ is based on this property; gunpowder
consists of a mechanical mixture of finely-ground sulphur, nitre, and
charcoal. The relative proportion of these substances varies according
to the destination of the powder and to the kind of charcoal employed
(a friable, incompletely-burnt charcoal, containing therefore hydrogen
and oxygen, is employed). Gases are formed in its combustion, chiefly
nitrogen and carbonic anhydride, which create a considerable pressure
if their escape be in any way impeded. This action of gunpowder may be
expressed by the equation: 2KNO_{3} + 3C + S = K_{2}S + 3CO_{2} + N_{2}.

  [14 bis] Before fusing, the crystals of potassium nitrate change their
       form, and take the same form as sodium nitrate--that is, they
       change into rhombohedra. Nitre crystallises from hot solutions,
       and in general under the influence of a rise of temperature, in a
       different form from that given at the ordinary or lower
       temperatures. Fused nitre solidifies to a radiated crystalline
       mass; but it does not exhibit this structure if metallic chlorides
       be present, so that this method may be taken advantage of to
       determine the degree of purity of nitre.

       Carnelley and Thomson (1888) determined the fusing point of
       mixtures of potassium and sodium nitrates. The first salt fuses at
       339° and the second at 316°, and if _p_ be the percentage amount
       of potassium nitrate, then the results obtained were--

         _p_ = 10   20    30    40    50    60    70    80    90
            298°   283°  268°  242°  231°  231°  242°  284°  306°

       which confirms Shaffgotsch's observation (1857) that the lowest
       fusing point (about 231°) is given by mixing molecular quantities
       (_p_ = 54·3) of the salts--that is, in the formation of the alloy,
       KNO_{3},NaNO_{3}.

       A somewhat similar result was discovered by the same observers for
       the solubility of mixtures of these salts at 20° in 100 parts of
       water. Thus, if _p_ be the weight of potassium nitrate mixed with
       100-_p_ parts by weight of sodium nitrate taken for solution, and
       _c_ be the quantity of the mixed salts which dissolves in 100, the
       solubility of sodium nitrate being 85, and of potassium nitrate
       34, parts in 100 parts of water, then--

              _p_ =  10   20   30   40   50  60  70  80  90
              _c_ = 110  136  136  138  106  81  73  54  41

       The maximum solubility proved not to correspond with the most
       fusible mixture, but to one much richer in sodium nitrate.

       Both these phenomena show that in homogeneous liquid mixtures the
       chemical forces that act between substances are the same as those
       that determine the molecular weights of substances, even when the
       mixture consists of such analogous substances as potassium and
       sodium nitrates, between which there is no direct chemical
       interchange. It is instructive to note also that the maximum
       solubility does not correspond with the minimum fusing point,
       which naturally depends on the fact that in solution a third
       substance, namely water, plays a part, although an attraction
       between the salts, like that which exists between sodium and
       potassium carbonates (Note 8), also partially acts.

  [15] Fused nitre, with a further rise of temperature, disengages oxygen
       and then nitrogen. The nitrite KNO_{2} is first formed and then
       potassium oxide. The admixture of certain metals--for example, of
       finely-divided copper--aids the last decomposition. The oxygen in
       this case naturally passes over to the metal.

It is found by this equation that gunpowder should contain thirty-six
parts of charcoal (13·3 p.c.), and thirty-two parts (11·9 p.c.) of
sulphur, to 202 parts (74·8 p.c.) of nitre, which is very near to its
actual composition.[16]

  [16] In China, where the manufacture of gunpowder has long been carried
       on, 75·7 parts of nitre, 14·4 of charcoal, and 9·9 of sulphur are
       used. Ordinary powder for sporting purposes contains 80 parts of
       nitre, 12 of charcoal, and 8 of sulphur, whilst the gunpowder used
       in heavy ordnance contains 75 of nitre, 15 of charcoal, and 10 of
       sulphur. Gunpowder explodes when heated to 300°, when struck, or
       by contact with a spark. A compact or finely-divided mass of
       gunpowder burns slowly and has but little disruptive action,
       because it burns gradually. To act properly the gunpowder must
       have a definite rate of combustion, so that the pressure should
       increase during the passage of the projectile along the barrel of
       the fire-arm. This is done by making the powder in large granules
       or in the shape of six-sided prisms with holes through them
       (prismatic powder).

       The products of combustion are of two kinds: (1) gases which
       produce the pressure and are the cause of the dynamical action of
       gunpowder, and (2) a solid residue, usually of a black colour
       owing to its containing unburnt particles of charcoal. Besides
       charcoal, the residue generally contains potassium sulphide,
       K_{2}S, and a whole series of other salts--for instance, carbonate
       and sulphate. It is apparent from this that the combustion of
       gunpowder is not so simple as it appears to be from the above
       formula, and hence the weight of the residue is also greater than
       indicated by that formula. According to the formula, 270 parts of
       gunpowder give 110 parts of residue--that is, 100 parts of powder
       give 37·4 parts of residue, K_{2}S, whilst in reality the weight
       of the residue varies from 40 p.c. to 70 p.c. (generally 52 p.c.).
       This difference depends on the fact that so much oxygen (of the
       nitre) remains in the residue, and it is evident that if the
       residue varies the composition of the gases evolved by the powder
       will vary also, and therefore the entire process will be different
       in different cases. The difference in the composition of the gases
       and residue depends, as the researches of Gay-Lussac, Shishkoff
       and Bunsen, Nobel and Abel, Federoff, Debus, &c., show, on the
       conditions under which the combustion of the powder proceeds. When
       gunpowder burns in an open space, the gaseous products which are
       formed do not remain in contact with the residue, and then a
       considerable portion of the charcoal entering into the composition
       of the powder remains unburnt, because the charcoal burns after
       the sulphur at the expense of the oxygen of the nitre. In this
       extreme case the commencement of the combustion of the gunpowder
       may be expressed by the equation, 2KNO_{3} + 3C + S = 2C +
       K_{2}SO_{4} + CO_{2} + N_{2}. The residue in a blank cartridge
       often consists of a mixture of C, K_{2}SO_{4}, K_{2}CO_{3}, and
       K_{2}S_{2}O_{3}. If the combustion of the gunpowder be impeded--if
       it take place in a cartridge in the barrel of a gun--the quantity
       of potassium sulphate will first be diminished, then the amount of
       sulphite, whilst the amount of carbonic anhydride in the gases and
       the amount of potassium sulphide in the residue will increase. The
       quantity of charcoal entering into the action will then be also
       increased, and hence the amount in the residue will decrease.
       Under these circumstances the weight of the residue will be
       less--for example, 4K_{2}CO_{3} + 4S = K_{2}SO_{4} + 3K_{2}S +
       4CO_{2}. Besides which, carbonic oxide has been found in the
       gases, and potassium bisulphide, K_{2}S_{2}, in the residue of
       gunpowder. The amount of potassium sulphide, K_{2}S, increases
       with the completeness of the combustion, and is formed in the
       residue at the expense of the potassium sulphite. In recent times
       the knowledge of the action of gunpowder and other explosives has
       made much progress, and has developed into a vast province of
       artillery science, which, guided by the discoveries of chemistry,

       has worked out a 'smokeless powder' which burns without leaving a
       residue, and does not therefore give any 'powder smoke' (to hinder
       the rapidity of firing and aiming), and at the same time
       disengages a greater volume of gas and consequently gives (under
       proper conditions of combustion) the possibility of communicating
       to the charge a greater initial velocity, and therefore greater
       distance, force, and accuracy of aim. Such 'smokeless powder' is
       prepared either from the varieties of nitrocellulose (Chapter VI.,
       Note 37) or from a mixture of them with nitro-glycerine (_ibid_).
       In burning they give, besides steam and nitrogen, generally a
       large amount of oxide of carbon (this is a very serious drawback
       in all the present forms of smokeless powder, because carbonic
       oxide is poisonous), and also CO_{2}, H_{2}, &c.

_Metallic potassium_ was obtained like sodium; first by the action of
a galvanic current, then by reduction of the hydroxide by means of
metallic iron, and lastly, by the action of charcoal on the carbonate
at a high temperature. The behaviour of metallic potassium differs,
however, from that of sodium, because it easily combines with carbonic
oxide, forming an explosive and inflammable mass.[17]

  [17] The substances obtained in this case are mentioned in Chapter IX.,
       Note 31.

Potassium is quite as volatile as sodium, if not more so. At the
ordinary temperature potassium is even softer than sodium; its
freshly-cut surfaces present a whiter colour than sodium, but, like
the latter, and with even greater ease, it oxidises in moist air. It
is brittle at low temperatures, but is quite soft at 25°, and melts at
58°. At a low red heat (667°, Perkin) it distils without change, forming
a green vapour, whose density,[18] according to A. Scott (1887), is
equal to 19 (if that of hydrogen = 1). This shows that the molecule
of potassium (like that of sodium, mercury, and zinc) contains but one
atom. This is also the case with many other metals, judging by recent
researches.[19] The specific gravity of potassium at 15° is 0·87, and
is therefore less than that of sodium, as is also the case with all
its compounds.[20] Potassium decomposes water with great ease at the
ordinary temperature, evolving 45,000 heat units per atomic weight in
grams. The heat evolved is sufficient to inflame the hydrogen, the flame
being  violet from the presence of particles of potassium.[21]

  [18] A. Scott (1887) determined the vapour densities of many of the
       alkali elements and their compounds in a platinum vessel heated in
       a furnace and previously filled with nitrogen. But these, the
       first data concerning a subject of great importance, have not yet
       been sufficiently fully described, nor have they received as much
       attention as could be desired. Taking the density of hydrogen as
       unity, Scott found the vapour densities of the following
       substances to be--

              Na         12·75 (11·5).      KI      92 (84).
              K          19    (19·5).      RbCl    70 (60).
              CsCl       89·5  (84·2).      CsI    133 (130).
              FeCl_{3}   68.                AgCl    80 (71·7).

       In brackets are given the densities corresponding with the
       formulæ, according to Avogadro-Gerhardt's law. This figure is not
       given for FeCl_{3}, because in all probability under these
       conditions (the temperature at which it was determined) a portion
       of the FeCl_{3} was decomposed. If it was not decomposed, then a
       density 81 would correspond with the formula FeCl_{3}, and if the
       decomposition were Fe_{2}Cl_{6} = 2FeCl_{2} + Cl_{2}, then the
       density should be 54. With regard to the silver chloride, there is
       reason to think that the platinum decomposed this salt. The
       majority of Scott's results so closely correspond with the formulæ
       that a better concord cannot be expected in such determinations.
       V. Meyer (1887) gives 93 as the density of KI.

  [19] The molecules of non-metals are more complex--for instance, H_{2},
       O_{3}, Cl_{2}, &c. But arsenic, whose superficial appearance
       recalls that of metals, but whose chemical properties approach
       more nearly to the non-metals, has a complex molecule containing
       As_{4}.

  [20] As the atomic weight of potassium is greater than that of sodium,
       the volumes of the molecules, or the quotients of the molecular
       weight by the specific gravity, for potassium compounds are
       greater than those of sodium compounds, because both the
       denominator and numerator of the fraction increase. We cite for
       comparison the volumes of the corresponding compounds--

       Na  24    NaHO  18    NaCl  28    NaNO_{3}  37    Na_{2}SO_{4}  54
       K   45    KHO   27    KCl   39    KNO_{3}   48    K_{2}SO_{4}   66

  [21] The same precautions must be taken in decomposing water by
       potassium as have to be observed with sodium (Chapter II., Note
       8).

       It must be observed that potassium decomposes carbonic anhydride
       and carbonic oxide when heated, the carbon being liberated and the
       oxygen taken up by the metal, whilst on the other hand charcoal
       takes up oxygen from potassium, as is seen from the preparation of
       potassium by heating potash with charcoal, hence the reaction
       K_{2}O + C = K_{2} + CO is reversible and the relation is the same
       in this case as between hydrogen and zinc.

With regard to the relation of potassium to hydrogen and oxygen, it
is closely analogous to sodium in this respect. Thus, with hydrogen
it forms potassium hydride, K_{2}H (between 200° and 411°), and with
oxygen it gives a suboxide K_{4}O, oxide K_{2}O, and peroxide, only
more oxygen enters into the composition of the latter than in sodium
peroxide; potassium peroxide contains KO_{2}, but it is probable that
in the combustion of potassium an oxide KO is also formed. Potassium,
like sodium, is soluble in mercury.[22] In a word, the relation between
sodium and potassium is as close as that between chlorine and bromine,
or, better still, between fluorine and chlorine, as the atomic weight
of sodium, 23, is as much greater than that of fluorine, 19, as that of
potassium, 39, is greater than that of chlorine, 35·5.

  [22] _Potassium_ forms _alloys with sodium_ in all proportions. The
       alloys containing 1 and 3 equivalents of potassium to one
       equivalent of sodium are _liquids_, like mercury at the ordinary
       temperature. Joannis, by determining the amount of heat developed
       by these alloys in decomposing water, found the evolution for
       Na_{2}K, NaK, NaK_{2} and NaK_{3} to be 44·5, 44·1, 43·8 and 44·4
       thousand heat units respectively (for Na 42·6 and for K 45·4). The
       formation of the alloy NaK_{2} is therefore accompanied by the
       development of heat, whilst the other alloys may be regarded as
       solutions of potassium or sodium in this alloy. In any case a fall
       of the temperature of fusion is evident in this instance as in the
       alloys of nitre (Note 14). The liquid alloy NaK_{2} is now used
       for filling thermometers employed for temperatures above 360°,
       when mercury boils.

The resemblance between _potassium_ and _sodium_ is so great that _their
compounds_ can only be easily _distinguished_ in the form of certain of
their salts. For instance, the acid potassium tartrate, C_{4}H_{5}KO_{6}
(cream of tartar), is distinguished by its sparing solubility in
water and in alcohol, and in a solution of tartaric acid, whilst the
corresponding sodium salt is easily soluble. Therefore, if a solution of
tartaric acid be added in considerable excess to the solutions of the
majority of potassium salts, a precipitate of the sparingly-soluble acid
salt is formed, which does not occur with salts of sodium. The chlorides
KCl and NaCl in solutions easily give double salts K_{2}PtCl_{6} and
Na_{2}PtCl_{6}, with platinic chloride, PtCl_{4}, and the solubility
of these salts is very different, especially in a mixture of alcohol
and ether. The sodium salt is easily soluble, whilst the potassium salt
is insoluble or almost so, and therefore the reaction with platinic
chloride is that most often used for the separation of potassium from
sodium, as is more fully described in works on analytical chemistry.

It is possible to discover the least traces of these metals in
admixture together, by means of their property of imparting different
colours to _a flame_. The presence of a salt of sodium in a flame is
recognised by a brilliant yellow coloration, and a pure potassium
salt colours a colourless flame violet. However, in the presence of
a sodium salt the pale violet coloration given by a potassium salt
is quite undistinguishable, and it is at first sight impossible in
this case to discover the potassium salt in the presence of that of
sodium. But by decomposing the light given by a flame  by these
metals or a mixture of them, by means of a prism, they are both easily
distinguishable, because the yellow light emitted by the sodium salt
depends on a group of light rays having a definite index of refraction
which corresponds with the yellow portion of the solar spectrum, having
the index of refraction of the Fraunhofer line (strictly speaking,
group of lines) D, whilst the salts of potassium give a light from
which these rays are entirely absent, but which contain rays of a red
and violet colour. Therefore, if a potassium salt occur in a flame, on
decomposing the light (after passing it through a narrow slit) by means
of a prism, there will be seen red and violet bands of light situated
at a considerable distance from each other; whilst if a sodium salt be
present a yellow line will also appear. If both metals simultaneously
occur in a flame and emit light, the spectrum lines corresponding to the
potassium and the sodium will appear simultaneously.

[Illustration: FIG. 72.--Spectroscope. The prism and table are covered
with an opaque cover. The spectrum obtained from the flame 
by a substance introduced on the wire is viewed through B. A light is
placed before the scale D in order to illuminate the image of the scale
reflected through B by the side of the prism.]

For convenience in carrying on this kind of testing, _spectroscopes_
(fig. 72) are constructed,[23] consisting of a refracting prism and
three tubes placed in the plane of the refracting angle of the prism.
One of the tubes, C, has a vertical slit at the end, giving access to
the light to be tested, which then passes into the tube (collimator),
containing a lens which gives the rays a parallel direction. The rays
of light having passed through the slit, and having become parallel,
are refracted and dispersed in the prism, and the spectrum formed is
observed through the eye-piece of the other telescope B. The third
tube D contains a horizontal transparent scale (at the outer end)
which is divided into equal divisions. The light from a source such
as a gas burner or candle placed before this tube, passes through the
scale, and is reflected on that face of the prism which stands before
the telescope B, so that the image of the scale is seen through this
telescope simultaneously with the spectrum given by the rays passing
through the slit of the tube C. In this manner the image of the scale
and the spectrum given by the source of light under investigation are
seen simultaneously. If the sun's rays be directed through the slit
of the tube C, then the observer looking through the eye-piece of B
will see the solar spectrum, and (if the aperture of the slit be narrow
and the apparatus correctly adjusted) the dark Fraunhofer lines in
it.[24] Small-sized spectroscopes are usually so adjusted that (looking
through B) the violet portion of the spectrum is seen to the right and
the red portion to the left, and the Fraunhofer line D (in the bright
yellow portion of the spectrum) is situated on the 50th division of the
scale.[25] If the light emitted by an incandescent solid--for example,
the Drummond light--be passed through the spectroscope, then all the
colours of the solar spectrum are seen, but not the Fraunhofer lines. To
observe the result given by a flame  by various salts a Bunsen
gas burner (or the pale flame of hydrogen gas issuing from a platinum
orifice) giving so pale a flame that its spectrum will be practically
invisible is placed before the slit. If any compound of sodium be placed
in the flame of the gas burner (for which purpose a platinum wire on
whose end sodium chloride is fused is fixed to the stand), then the
flame is  yellow, and on looking through the spectroscope the
observer will see a bright _yellow_ line falling upon the 50th division
of the scale, which is seen together with the spectrum in the telescope.
No yellow lines of other refractive index, nor any rays of any other
colour, will be seen, and, therefore, the spectrum corresponding with
sodium compounds consists of yellow rays of that index of refraction
which belong to the Fraunhofer (black) line D of the solar spectrum.
If a potassium salt be introduced into the flame instead of a sodium
salt, then two bands will be seen which are much feebler than the
bright sodium band--namely, one red line near the Fraunhofer line A and
another violet line. Besides which, a pale, almost continuous, spectrum
will be observed in the central portions of the scale. If a mixture
of sodium and potassium salts be now introduced into the flame, three
lines will be seen simultaneously--namely, the red and pale violet
lines of potassium and the yellow line of sodium. In this manner it is
possible, by the aid of the spectroscope, to determine the relation
between the spectra of metals and known portions of the solar spectrum.
The continuity of the latter is interrupted by dark lines (that is,
by an absence of light of a definite index of refraction), termed the
Fraunhofer lines of the solar spectrum. It has been shown by careful
observations (by Fraunhofer, Brewster, Foucault, Ångstrom, Kirchhoff,
Cornu, Lockyer, Dewar, and others) that there exists an exact _agreement
between the spectra_ of certain _metals_ and certain of the _Fraunhofer_
lines. Thus the bright yellow sodium line exactly corresponds with the
dark Fraunhofer line D of the solar spectrum. A similar agreement is
observed in the case of many other metals. This is not an approximate or
chance correlation. In fact, if a spectroscope having a large number of
refracting prisms and a high magnifying power be used, it is seen that
the dark line D of the solar spectrum consists of an entire system of
closely adjacent but definitely situated fine and wide (sharp, distinct)
dark lines,[26] and an exactly similar group of bright lines is obtained
when the yellow sodium line is examined through the same apparatus, so
that each bright sodium line exactly corresponds with a dark line in the
solar spectrum.[26 bis] This conformity of the bright lines formed by
sodium with the dark lines of the solar spectrum cannot be accidental.
This conclusion is further confirmed by the fact that the bright lines
of other metals correspond with dark lines of the solar spectrum. Thus,
for example, a series of sparks passing between the iron electrodes
of a Ruhmkorff coil gives 450 very distinct lines characterising this
metal. All these 450 bright lines, constituting the whole spectrum
corresponding with iron, are repeated, as Kirchhoff showed, in the
solar spectrum as dark Fraunhofer lines which occur in exactly the
same situations as the bright lines in the iron spectrum, just as the
sodium lines correspond with the band D in the solar spectrum. Many
observers have in this manner studied the solar spectrum and the spectra
of different metals simultaneously, and discovered in the former lines
which correspond not only with sodium and iron, but also with many
other metals.[27] The spectra of such elements as hydrogen, oxygen,
nitrogen, and other gases may be observed in the so-called Geissler's
tubes--that is, in glass tubes containing rarefied gases, through which
the discharge of a Ruhmkorff's coil is passed. Thus hydrogen gives a
spectrum composed of three lines--a red line corresponding with the
Fraunhofer line C, a green line corresponding with the line F, and a
violet line corresponding with one of the lines between G and H. Of
these rays the red is the brightest, and therefore the general colour of
luminous hydrogen (with an electric discharge through a Geissler tube)
is reddish.

  [23] For accurate measurements and comparative researches more
       complicated spectroscopes are required which give a greater
       dispersion, and are furnished for this purpose with several
       prisms--for example, in Browning's spectroscope the light passes
       through six prisms, and then, having undergone an internal total
       reflection, passes through the upper portion of the same six
       prisms, and again by an internal total reflection passes into the
       ocular tube. With such a powerful dispersion the relative position
       of the spectral lines may be determined with accuracy. For the
       absolute and exact determination of the wave lengths it is
       particularly important that the spectroscope should be furnished
       with diffraction gratings. The construction of spectroscopes
       destined for special purposes (for example, for investigating the
       light of stars, or for determining the absorption spectra in
       microscopic preparations, &c.) is exceedingly varied. Details of
       the subject must be looked for in works on physics and on spectrum
       analysis. Among the latter the best known for their completeness
       and merit are those of Roscoe, Kayser, Vogel, and Lecoq de
       Boisbaudran.

  [24] The arrangement of all the parts of the apparatus so as to give
       the clearest possible vision and accuracy of observation must
       evidently precede every kind of spectroscopic determination.
       Details concerning the practical use of the spectroscope must be
       looked for in special works on the subject. In this treatise the
       reader is supposed to have a certain knowledge of the physical
       data respecting the refraction of light, and its dispersion and
       diffraction, and the theory of light, which allows of the
       determination of the length of the waves of light in absolute
       measure on the basis of observations with diffraction gratings,
       the distance between whose divisions may be easily measured in
       fractions of a millimetre; by such means it is possible to
       determine the wave-length of any given ray of light.

  [25] In order to give an idea of the size of the scale, we may observe
       that the ordinary spectrum extends from the zero of the scale
       (where the red portion is situated) to the 170th division (where
       the end of the visible violet portion of the spectrum is
       situated), and that the Fraunhofer line A (the extreme prominent
       line in the red) corresponds with the 17th division of the scale;
       the Fraunhofer line F (at the beginning of the blue, near the
       green colour) is situated on the 90th division, and the line G,
       which is clearly seen in the beginning of the violet portion of
       the spectrum, corresponds with the 127th division of the scale.

  [26] The two most distinct lines of D, or of sodium, have wave-lengths
       of 589·5 and 588·9 millionths of a millimeter, besides which
       fainter and fainter lines are seen whose wave-lengths in
       millionths of a millimeter are 588·7 and 588·1, 616·0 and 615·4,
       515·5 and 515·2, 498·3 and 498·2, &c., according to Liveing and
       Dewar.

  [26 bis] In the ordinary spectroscopes which are usually employed in
       chemical research, one yellow band, which does not split up into
       thinner lines, is seen instead of the system of sodium lines,
       owing to the small dispersive power of the prism and the width of
       the slit of the object tube.

  [27] The most accurate investigations made in this respect are carried
       on with spectra obtained by diffraction, because in this case the
       position of the dark and bright lines does not depend on the index
       of refraction of the material of the prism, nor on the dispersive
       power of the apparatus. The best--that is, the most general and
       accurate--method of expressing the results of such determinations
       consists in determining the lengths of the waves corresponding to
       the rays of a definite index of refraction. (Sometimes instead of
       this the fraction of 1 divided by the square of the wave-length is
       given.) We will express this _wave-length_ in _millionth parts of
       a millimetre_ (the ten-millionth parts are already doubtful, and
       fall within the limits of error). In order to illustrate the
       relation between the wave-lengths and the positions of the lines
       of the spectrum, we will cite the wave-lengths corresponding with
       the chief Fraunhofer lines and colours of the spectrum.

       Fraunhofer line   A      B      C         D         E       b
       Wave-length     761·0  687·5  656·6  589·5-588·9  527·3   518·7
                       +-----------------+  +---------+  +----+  +---+
       Colour                 red              orange    yellow  green

       Fraunhofer line   F      G      H
       Wave-length     486·5  431·0  397·2
                       +----------+  +----+
       Colour              blue      violet

       In the following table are given the _wave-lengths_ of the light
       rays (the longest and most distinct, _see_ later) for certain
       elements, those in black type being the most clearly defined and
       distinct lines, which are easily obtained either in the flame of a
       Bunsen's burner, or in Geissler's tubes, or in general, by an
       electric discharge. These lines refer to the elements (the lines
       of compounds are different, as will be afterwards explained, but
       many compounds are decomposed by the flame or by an electric
       discharge), and moreover to the elements in an incandescent and
       rarefied gaseous state, for the spectra sometimes vary
       considerably with a variation of temperature and pressure.

       It may be mentioned that the _red_ colour corresponds with lines
       having a wave-length of from 780 (with a greater wave-length the
       lines are hardly visible, and are ultra red) to 650, the _orange_
       from 650 to 590, the _yellow_ from 590 to 520, the _green_ from
       520 to 490, the _blue_ from 490 to 420, and the _violet_ from 420
       to 380 millionth parts of a millimetre. Beyond 380 the lines are
       scarcely visible, and belong to the ultra-violet. For fluorine
       Moissan found as many as 13 bright lines from 744 to 623.

       In the table (p. 565) which is arranged in conformity with the
       image of the spectrum as it is seen (the red lines on the
       left-hand and the violet on the right-hand side), the figures in
       black type correspond with lines which are so bright and
       distinctly visible that they may easily be made use of, both in
       determining the relation between the divisions of the scale and
       the wave-lengths, and in determining the admixture of a given
       element with another. Brackets join those lines between which
       several other lines are clearly visible if the dispersive power of
       the spectroscope permits distinguishing the neighbouring lines. In
       the ordinary laboratory spectroscopes with one prism, even with
       all possible precision of arrangement and with a brilliancy of
       light permitting the observations being made with a very narrow
       aperture, the lines whose wave-lengths only differ by 2-3
       millionths of a millimetre, are blurred together; and with a wide
       aperture a series of lines differing by even as much as 20
       millionths of a millimetre appear as one wide line. With a faint
       light (that is, with a small quantity of light entering into the
       spectroscope) only the most _brilliant_ lines are clearly visible.
       The _length_ of the lines does not always correspond with their
       brilliancy. According to Lockyer this length is determined by
       placing the carbon electrodes (between which the incandescent
       vapours of the metals are formed), not horizontally to the slit
       (as they are generally placed, to give more light), but vertically
       to it. Then certain lines appear long and others short. As a rule
       (Lockyer, Dewar, Cornu), the longest lines are those with which it
       is easiest to obtain _reversed_ spectra (_see_ later).
       Consequently, these lines are the most characteristic. Only the
       longest and most brilliant are given in our table, which is
       composed on the basis of a collection of the data at our disposal
       for _bright_ spectra of the _incandescent and rarefied vapours of
       the elements_. As the spectra change with great variations of
       temperature and vapour density (the faint lines become brilliant
       whilst the bright lines sometimes disappear), which is
       particularly clear from Ciamician's researches on the halogens,
       until the method of observation and the theory of the subject are
       enlarged, particular theoretical importance should not be given to
       the wave-lengths showing the maximum brilliancy, which only
       possess a practical significance in the common methods of
       spectroscopic observations. In general the spectra of metals are
       simpler than those of the halogens, and the latter are variable;
       at an increased pressure all spectral lines become broader.

       +-------+-------+--------+--------+-------+-------+------+------+
       | N_{2} | O_{2} | Cl_{2} | Br_{2} | I_{2} |  Pb   |  Sn  |  Tl  |
       +-------+-------+--------+--------+-------+-------+------+------+
       |   --  |   --  |   --   |   --   |   --  |  --   |  --  |  --  |
       |   --  |   --  |   --   |   --   |   --  |  --   |  --  |  --  |
       |  662  |   --  |   --   |   --   |   --  |  --   |  --  |  --  |
       |  632} |   --  |   --   |  636   |   --  |  --   | 645  |  --  |
       |  620} |  615  |   --   |   --   |  621  |  --   |  --  |  --  |
       |  585} |   --  |   --   |   --   |  613  | 605·7 | 580} |  --  |
       |  574} |   --  |  546}  |   --   |  579} | 560·7 | 556} |  --  |
       |  544} |  543  |  539}  |  544}  |  560} | 554·7 |  --  | 549  |
       |  535} |  533  |  528}  |  523}  |  545} | 537   |  --  | 535  |
       |  527} |  516  |  519}  |  517   |  506} |  --   |  --  |  --  |
       |  516} |  495} |  494}  |  479   |   --  |  --   |  --  | 489  |
       |   --  |  494} |  480}  |  470}  |   --  |  --   |  --  |  --  |
       |  457  |  470} |   --   |  462}  |   --  |  --   | 452  |  --  |
       |  442  |  465} |   --   |  454   |  445  |  --   |  --  |  --  |
       |  436} |  447} |  436}  |  437   |   --  |  --   |  --  |  --  |
       |  426} |  432} |  431}  |   --   |   --  |  --   |  --  |  --  |
       |   --  |   --  |   --   |   --   |  421  |  --   |  --  |  --  |
       |   --  |   --  |   --   |   --   |   --  |  --   |  --  |  --  |
       |  409  |   --  |   --   |   --   |   --  | 406   |  --  |  --  |
       |   --  |   --  |   --   |   --   |   --  |  --   |  --  |  --  |
       +-------+-------+--------+--------+-------+-------+------+------+

       +-----+-----+-----+-------+-----+------+------+-----+-------+-----+
       |  In |  Ga |  Al |  Ba   |  Sr |  Ca  |  Mg  |  Zn |  Cd   |  Hg |
       |-----+-----+-----+-------+-----|------+------|-----+-------+-----+
       |  -- |  -- |  -- |  --   |  -- |  --  |  --  |  -- |  --   |  -- |
       |  -- |  -- |  -- |  --   |  -- |  --  |  --  |  -- |  --   |  -- |
       |  -- |  -- |  -- |  --   |  -- | 646  |  --  |  -- |  --   |  -- |
       |  -- |  -- | 624 | 649·7 | 641 | 644  |  --  | 636 | 643·8 |  -- |
       | 619 |  -- | 623 | 614   | 606 |  --  |  --  |  -- |  --   | 615 |
       |  -- |  -- |  -- |  --   |  -- | 612  |  --  |  -- |  --   | 579 |
       |  -- |  -- | 572 | 553·5 |  -- |  --  |  --  |  -- |  --   | 577 |
       |  -- |  -- | 570 | 549   | 548 | 559  |  --  |  -- | 537·7 | 546 |
       | 525 |  -- |  -- |  --   | 524 |  --  | 518  |  -- | 533·6 |  -- |
       |  -- |  -- |  -- |  --   |  -- |  --  | 516  | 492 | 508·5 |  -- |
       |  -- |  -- |  -- | 493·3 |  -- |  --  |  --  | 481 |  --   |  -- |
       |  -- |  -- |  -- |  --   |  -- |  --  | 471  | 472 | 479·9 |  -- |
       | 451 |  -- | 466 | 455   | 460 |  --  |  --  | 468 | 467·7 |  -- |
       |  -- |  -- |  -- |  --   |  -- | 445} | 448  |  -- |  --   |  -- |
       |  -- |  -- |  -- |  --   |  -- | 442} |  --  |  -- |  --   | 436 |
       |  -- |  -- |  -- |  --   | 430 |  --  |  --  |  -- |  --   |  -- |
       |  -- | 417 |  -- |  --   | 421 | 423  |  --  |  -- |  --   |  -- |
       |  -- |  -- |  -- | 413   |  -- |  --  |  --  |  -- |  --   |  -- |
       | 410 | 403 | 396 |  --   | 408 | 397  | 384} |  -- |  --   | 404 |
       |  -- |  -- | 394 |  --   |  -- | 393  | 383} |  - -|  --   |  -- |
       +-----+-----+-----+-------+-----+------+------+-----+-------+-----+

       +------+------+-----+-------+-----+-------+------+-------+-------+
       |  Mn  |  Fe  |  Cu |  Ag   |  Cs |  Rb   |  K   |  Na   |  Li   |
       |------+------+-----+-------+-----+-------+------+-------+-------+
       |  --  |  --  |  -- |  --   |  -- |  780  | 770  |   --  |  --   |
       |  --  |  --  |  -- |  --   |  -- |   --  | 766  |   --  |  --   |
       |  --  |  --  |  -- |  --   |  -- |   --  |  --  |   --  | 670·6 |
       |  --  | 640  |  -- |  --   |  -- |   --  |  --  |   --  |  --   |
       | 602} |  --  |  -- |  --   | 622 | 629·6 |  --  |   --  | 610   |
       | 601} | 561} | 578 |  --   | 600 |   --  | 583} | 589·5 |  --   |
       | 551} | 544} | 570 | 546·4 |  -- |   --  | 578} | 588·9 |  --   |
       | 534} | 537  | 522 |  --   |  -- |   --  |  --  |   --  |  --   |
       |  --  | 532} | 515 | 520·8 |  -- |   --  | 535  |   --  |  --   |
       |  --  | 521} | 511 |  --   |  -- |   --  | 532} |   --  |  --   |
       | 482} | 496} |  -- |  --   |  -- |   --  |  --  |   --  | 497   |
       | 471} | 489} |  -- |  --   |  -- |   --  |  --  |   --  |  --   |
       |  --  |  --  |  -- |  --   | 459 |   --  |  --  |   --  | 460·3 |
       |  --  | 441} |  -- |  --   | 456 |   --  |  --  |   --  |  --   |
       |  --  | 430} |  -- |  --   |  -- |   --  |  --  |   --  |  --   |
       |  --  | 427  |  -- |  --   |  -- |   --  |  --  |   --  |  --   |
       | 424} |  --  |  -- | 421   |  -- |  420  |  --  |   --  |  --   |
       | 403} | 407} |  -- |  --   |  -- |   --  |  --  |   --  |  --   |
       |  --  | 404} |  -- |  --   |  -- |   --  | 404  |   --  |  --   |
       |  --  |  --  |  -- |  --   |  -- |   --  |  --  |   --  |  --   |
       +------+------+-----+-------+-----+-------+------+-------+-------+

       +-------+
       | H_{2} |
       +-------+
       |  --   |
       |  --   |
       |  --   |
       | 656·2 |
       |  --   |
       |  --   |
       |  --   |
       |  --   |
       |  --   |
       |  --   |
       | 486·1 |
       |  --   |
       |  --   |
       |  --   |
       | 434   |
       |  --   |
       |  --   |
       |  --   |
       |  --   |
       |  --   |
       +-------+

[Illustration: FIG. 73.--Absorption spectrum (Lecoq de Boisbaudran) of
salts of didymium in concentrated and dilute solutions.]

The correlation of the Fraunhofer lines with the spectra of metals
depends on the phenomenon of the so-called _reversal of the spectrum_.
This phenomenon consists in this, that instead of the bright spectrum
corresponding with a metal, under certain circumstances a similar dark
spectrum in the form of Fraunhofer lines may be obtained, as will be
explained directly. In order to clearly understand the phenomenon of
reversed spectra, it must be known that when light passes through
certain transparent substances these substances retain rays of a certain
refrangibility. The colour of solutions is a proof of this. Light
which has passed through a yellow solution of a uranium salt contains
no violet rays, and after having passed through a red solution of a
permanganate, does not contain many rays in the yellow, blue, and green
portions of the spectrum. Solutions of copper salts absorb nearly all
red rays. Sometimes colourless solutions also absorb rays of certain
definite refractive indexes, and give _absorption spectra_. Thus
solutions of salts of didymium absorb rays of a certain refrangibility,
and therefore an impression of black lines is received,[28] as shown in
fig. 73. Many vapours (iodine) and gases (nitric peroxide) give similar
spectra. Light which has passed through a deep layer of aqueous vapour,
oxygen, or nitrogen also gives an absorption spectrum. For this reason
the peculiar (winter) dark lines discovered by Brewster are observed in
sunlight, especially in the evening and morning, when the sun's rays
pass through the atmosphere (containing these substances) by a longer
path than at mid-day. It is evident that the Fraunhofer lines may be
ascribed to the absorption of certain rays of light in its passage from
the luminous mass of the sun to the earth. The remarkable progress made
in all spectroscopic research dates from the investigations made by
_Kirchhoff_ (1859) on the relation between absorption spectra and the
spectra of luminous incandescent gases. It had already been observed
long before (by Fraunhofer, Foucault, Ångstrom) that the bright spectrum
of the sodium flame gives two bright lines which are in exactly the
same position as two black lines known as D in the solar spectrum,
which evidently belong to an absorption spectrum. When Kirchhoff caused
diffused sunlight to fall upon the slit of a spectroscope, and placed
a sodium flame before it, a perfect superposition was observed--the
bright sodium lines completely covered the black lines D of the solar
spectrum. When further the continuous spectrum of a Drummond light
showed the black line D on placing a sodium flame between it and the
slit of the spectroscope--that is, when the Fraunhofer line of the solar
spectrum was artificially produced--then there was no doubt that its
appearance in the solar spectrum was due to the light passing somewhere
through incandescent vapours of sodium. Hence a new theory of _reversed
spectra_[29] arose--that is, of the relation between the waves of
light emitted and absorbed by a substance under given conditions of
temperature; this is expressed by Kirchhoff's law, discovered by a
careful analysis of the phenomena. This law may be formulated in an
elementary way as follows: At a given temperature the relation between
the intensity of the light emitted (of a definite wave-length) and
the absorptive capacity with respect to the same colour (of the same
wave-length) is a constant quantity.[30] As a black dull surface emits
and also absorbs a considerable quantity of heat rays whilst a polished
metallic surface both absorbs and emits but few, so a flame 
by sodium emits a considerable quantity of yellow rays of a definite
refrangibility, and has the property of absorbing a considerable
quantity of the rays of the same refractive index. In general, the
medium which emits definite rays also absorbs them.

  [28] The method of observing absorption spectra consists in taking a
       continuous spectrum of white light (one which does not show either
       dark lines or particularly bright luminous bands--for instance,
       the light of a candle, lamp, or other source). The collimator
       (that is, the tube with the slit) is directed towards this light,
       and then all the colours of the spectrum are visible in the ocular
       tube. A transparent absorptive medium--for instance, a solution or
       tube containing a gas--is then placed between the source of light
       and the apparatus (or anywhere inside the apparatus itself in the
       path of the rays). In this case either the entire spectrum is
       uniformly fainter, or absorption bands appear on the bright field
       of the continuous spectrum in definite positions along it. These
       bands have different lengths and positions, and distinctness and
       intensity of absorption, according to the properties of the
       absorptive medium. Like the luminous spectra given by incandescent
       gases and vapours, the absorption spectra of a number of
       substances have already been studied, and some with great
       precision--as, for example, the spectrum of the brown vapours of
       nitrogen dioxide by Hasselberg (at Pulkowa), the spectra of
       colouring matters (Eder and others), especially of those applied
       to orthochromatic photography, the spectra of blood, chlorophyll
       (the green constituent of leaves), and other similar substances,
       all the more carefully as by the aid of their spectra the presence
       of these substances may be discovered in small quantities (even in
       microscopical quantities, by the aid of special appliances on the
       microscope), and the changes they undergo investigated.

       [Illustration: FIG. 74.--Absorption spectra of nitrogen dioxide
       and iodine.]

       The absorption spectra, obtained at the ordinary temperature and
       proper to substances in all physical states, offer a most
       extensive but as yet little studied field, both for the general
       theory of spectroscopy, and for gaining an insight into the
       structure of substances. The investigation of colouring matters
       has already shown that in certain cases a definite change of
       composition and structure entails not only a definite change of
       the colours but also a displacement of the absorption bands by a
       definite number of wave-lengths.

  [29] A number of methods have been invented to demonstrate the
       reversibility of spectra; among these methods we will cite two
       which are very easily carried out. In Bunsen's method sodium
       chloride is put into an apparatus for evolving hydrogen (the spray
       of the salt is then carried off by the hydrogen and colours the
       flame with the yellow sodium colour), and the hydrogen is ignited
       in two burners--in one large one with a wide flame giving a bright
       yellow sodium light, and in another with a small fine orifice
       whose flame is pale: this flame will throw a dark patch on the
       large bright flame. In Ladoffsky's method the front tube (p. 561)
       is unscrewed from a spectroscope directed towards the light of a
       lamp (a continuous spectrum), and the flame of a spirit lamp
        by a small quantity of NaCl is placed between the tube
       and the prism; a black band corresponding to sodium will then be
       seen on looking through the ocular tube. This experiment is always
       successful if only there be the requisite relation between the
       strength of light of the two lamps.

  [30] The absorptive capacity is the relation between the intensity
       of the light (of a given wave-length) falling upon and retained by
       a substance. Bunsen and Roscoe showed by direct experiment that
       this ratio is a constant quantity for every substance. If _A_
       stand for this ratio for a given substance at a given
       temperature--for instance, for a flame  by sodium--and _E_
       be the intensity of the light of the same wave-length emitted at
       the same temperature by the same substance, then Kirchhoff's law,
       the explanation and deduction of which must be looked for in
       text-books of physics, states that the fraction _A/E_ is a
       constant quantity depending on the nature of a substance (as _A_
       depends on it) and determined by the temperature and wave-length.

[Illustration: FIG. 75.--Bright spectra of copper compounds.]

Thus the bright spectral rays characteristic of a given metal may
be reversed--that is, converted into dark lines--by passing light
which gives a continuous spectrum through a space containing the
heated vapours of the given metal. A similar phenomenon to that thus
artificially produced is observed in sunlight, which shows dark lines
characteristic of known metals--that is, the Fraunhofer lines form
an absorption spectrum or depend on a reversed spectrum; it being
presupposed that the sun itself, like all known sources of artificial
light, gives a continuous spectrum without Fraunhofer lines.[31] We
must imagine that the sun, owing to the high temperature which is
proper to it, emits a brilliant light which gives a continuous spectrum,
and that this light, before reaching our eyes, passes through a space
full of the vapours of different metals and their compounds. As the
earth's atmosphere[32] contains very little, or no, metallic vapours,
and as they cannot be supposed to exist in the celestial space,[32 bis]
the only place in which the existence of such vapours can be admitted
is in the _atmosphere surrounding the sun itself_. As the cause of
the sun's luminosity must be looked for in its high temperature, the
existence of an atmosphere containing metallic vapours is readily
understood, because at that high temperature such metals as sodium,
and even iron, are separated from their compounds and converted into
vapour. The sun must be imagined as surrounded by an atmosphere of
incandescent vaporous and gaseous matter,[33] including those elements
whose reversed spectra correspond with the Fraunhofer lines--namely,
sodium, iron, hydrogen, lithium, calcium, magnesium, &c. Thus in
spectrum analysis we find a means of determining the composition of
the inaccessible heavenly luminaries, and much has been done in this
respect since Kirchhoff's theory was formulated. By observations on the
spectra of many heavenly bodies, changes have been discovered going on
in them,[34] and many of the elements known to us have been found with
certainty in them.[35] From this it must be concluded that the same
elements which exist on the earth occur throughout the whole universe,
and that at that degree of heat which is proper to the sun those simple
substances which we accept as the elements in chemistry are still
undecomposed and remain unchanged. A high temperature forms one of those
conditions under which compounds most easily decompose; and if sodium
or a similar element were a compound, in all probability it would be
decomposed into component parts at the high temperature of the sun. This
may indeed be concluded from the fact that in ordinary spectroscopic
experiments the spectra obtained often belong to the metals and not
to the compounds taken; this depends on the decomposition of these
compounds in the heat of the flame. If common salt be introduced
into the flame of a gas-burner, a portion of it is decomposed, first
forming, in all probability, with water, hydrochloric acid and sodium
hydroxide, and the latter then becoming partially decomposed by the
hydrocarbons, giving metallic sodium, whose incandescent vapour emits
light of a definite refrangibility. This conclusion is arrived at from
the following experiment:--If hydrochloric acid gas be introduced into
a flame  by sodium it is observed that the sodium spectrum
disappears, owing to the fact that metallic sodium cannot remain in the
flame in the presence of an excess of hydrochloric acid. The same thing
takes place on the addition of sal-ammoniac, which in the heat of the
flame gives hydrochloric acid. If a porcelain tube containing sodium
chloride (or sodium hydroxide or carbonate), and closed at both ends
by glass plates, be so powerfully heated that the salt volatilises,
then the sodium spectrum is not observable; but if the salt be replaced
by sodium, then either the bright line or the absorption spectra is
obtained, according to whether the light emitted by the incandescent
vapour be observed, or light passing through the tube. Thus the above
spectrum is not given by sodium chloride or other sodium compound, but
is proper to the metal sodium itself. This is also the case with other
analogous metals. The chlorides and other halogen _compounds_ of barium,
calcium, copper, &c., give independent spectra which differ from those
of the metals. If barium chloride be introduced into a flame, it gives
a mixed spectrum belonging to metallic barium and barium chloride. If
besides barium chloride, hydrochloric acid or sal-ammoniac be introduced
into the flame, then the spectrum of the metal disappears, and that of
the chloride remains, which differs distinctly from the spectrum of
barium fluoride, barium bromide, or barium iodide. A certain common
resemblance and certain common lines are observed in the spectra of
two different compounds of one and the same element obtained in the
above-described manner, and also in the spectrum of the metal, but they
all have their peculiarities. The independent spectra of the compounds
of copper are easily observed (fig. 75). Thus certain compounds which
exist in a state of vapour, and are luminous at a high temperature,
give their independent spectra. In the majority of cases the spectra of
compounds are composed of indistinct luminous lines and complete bright
bands, whilst metallic elements generally give a few clearly-defined
spectral lines.[36] There is no reason for supposing that the spectrum
of a compound is equal to the sum of the spectra of its elements--that
is, _every compound_ which is not decomposed by heat _has its own
proper spectrum_. This is best proved by absorption spectra, which are
essentially only reversed spectra observed at low temperatures. If every
salt of sodium, lithium, and potassium gives one and the same spectrum,
this must be ascribed to the presence in the flame of the free metals
liberated by the decomposition of their salts. Therefore _the phenomena
of the spectrum are determined by molecules, and not by atoms_--that
is, the molecules of the metal sodium, and not its atoms, produce those
particular vibrations which determine the spectrum of a sodium salt.
Where there is no free metallic sodium there is no sodium spectrum.

  [31] Heated metals begin to emit light (only visible in the dark) at
       about 420° (varying with the metal). On further heating, solids
       first emit red, then yellow, and lastly white light. Compressed or
       heavy gases (_see_ Chapter III., Note 44), when strongly heated,
       also emit white light. Heated liquids (for example, molten steel
       or platinum) also give a white compound light. This is readily
       understood. In a dense mass of matter the collisions of the
       molecules and atoms are so frequent that waves of only a few
       definite lengths cannot appear; the reverse is possible in
       rarefied gases or vapours.

  [32] Brewster, as is mentioned above, first distinguished the
       atmospheric, cosmical Fraunhofer lines from the solar lines.
       Janssen showed that the spectrum of the atmosphere contains lines
       which depend on the absorption produced by aqueous vapour.
       Egoreff, Olszewski, Janssen, and Liveing and Dewar showed by a
       series of experiments that the oxygen of the atmosphere gives rise
       to certain lines of the solar spectrum, especially the line A.
       Liveing and Dewar took a layer of 165 c.m. of oxygen compressed
       under a pressure of 85 atmospheres, and determined its absorption
       spectrum, and found that, besides the Fraunhofer lines A and B, it
       contained the following groups: 630-622, 581-568, 535, 480-475.
       The same lines were found for liquid oxygen.

  [32 bis] If the material of the whole heavenly space formed the
       absorbent medium, the spectra of the stars would be the same as
       the solar spectrum; but Huyghens, Lockyer, and others showed not
       only that this is the case for only a few stars, but that the
       majority of stars give spectra of a different character with dark
       and bright lines and bands.

  [33] Eruptions, like our volcanic eruptions, but on an incomparably
       larger scale, are of frequent occurrence on the sun. They are seen
       as protuberances visible during a total eclipse of the sun, in the
       form of vaporous masses on the edge of the solar disc and emitting
       a faint light. These protuberances of the sun are now observed at
       all times by means of the spectroscope (Lockyer's method), because
       they contain luminous vapours (giving bright lines) of hydrogen
       and other elements.

  [34] The great interest and vastness of astro-physical observations
       concerning the sun, comets, stars, nebulæ, &c., render this new
       province of natural science very important, and necessitate
       referring the reader to special works on the subject.

       The most important astro-physical data since the time of Kellner
       are those referring to the _displacement_ of the lines of the
       spectrum. Just as a musical note changes its pitch with the
       approach or withdrawal of the resonant object or the ear, so the
       pitch of the luminous note or wave-length of the light varies if
       the luminous (or absorbent) vapour and the earth from which we
       observe it approach or recede from each other; this expresses
       itself in a visible displacement of the spectral lines. The solar
       eruptions even give broken lines in the spectrum, because the
       rapidly moving eruptive masses of vapour and gases either travel
       in the direction of the eye or fall back towards the sun. As the
       earth travels with the solar system among the stars, so it is
       possible to determine the direction and velocity with which the
       sun travels in space by the displacement of the spectral lines and
       light of the stars. The changes proceeding on the sun in its mass,
       which must be pronounced as vaporous, and in its atmosphere, are
       now studied by means of the spectroscope. For this purpose, many
       special astro-physical observatories now exist where these
       investigations are carried on.

       We may remark that if the observer or luminous object moves with a
       velocity ±_v_, the ray, whose wave-length is [Greek: l], has an
       apparent wave-length [Greek: l](_n_±_v_)/_n_, where _n_ is the
       velocity of light. Thus Tolon, Huyghens, and others proved that
       the star Aldebaran approaches the solar system with a velocity of
       30 kilometres per second, while Arcturus is receding with a
       velocity of 45 kilometres. The majority of stars give a distinct
       hydrogen spectrum, besides which nebulæ also give the spectrum of
       nitrogen. Lockyer classes the stars from their spectra, according
       to their period of formation, showing that some stars are in a
       period of increasing temperature (of formation or aggregation),
       whilst others are in a period of cooling. Altogether, in the
       astro-physical investigation of the spectra of heavenly bodies we
       find one of the most interesting subjects of recent science.

  [35] Spectrum analysis has proved the indubitable existence in the sun
       and stars of a number of elements known in chemistry. Huyghens,
       Secchi, Lockyer, and others have furnished a large amount of
       material upon this subject. A compilation of existing information
       on it has been given by Prof. S. A. Kleiber, in the Journal of the
       Russian Physico-chemical Society for 1885 (vol. xviii. p. 146).
       Besides which, a peculiar element called helium has been
       discovered, which is characterised by a line (whose wave-length is
       587·5, situated near D), which is seen very brightly in the
       projections (protuberances) and spots of the sun, but which does
       not belong to any known element, and is not reproducible as a
       reversed, dark line. This may be a right conclusion--that is to
       say, it is possible that an element may be discovered to which the
       spectrum of helium corresponds--but it may be that the helium line
       belongs to one of the known elements, because spectra vary in the
       brilliancy and position of their lines with changes of temperature
       and pressure. Thus, for instance, Lockyer could only see the line
       423, at the very end of the calcium spectrum, at comparatively low
       temperatures, whilst the lines 397 and 393 appear at a higher
       temperature, and at a still higher temperature the line 423
       becomes quite invisible.

  [36] Spectroscopic observations are still further complicated by the
       fact that one and the same substance gives different spectra at
       different temperatures. This is especially the case with gases
       whose spectra are obtained by an electric discharge in tubes.
       Plücker, Wüllner, Schuster, and others showed that at low
       temperatures and pressures the spectra of iodine, sulphur,
       nitrogen, oxygen, &c. are quite different from the spectra of the
       same elements at high temperatures and pressures. This may either
       depend on the fact that the elements change their molecular
       structure with a change of temperature, just as ozone is converted
       into oxygen (for instance, from N_{2} molecules are obtained
       containing only one atom of nitrogen), or else it may be because
       at low temperature certain rays have a greater relative intensity
       than those which appear at higher temperatures. If we suppose that
       the molecules of a gas are in continual motion, with a velocity
       dependent on the temperature, then it must be admitted that they
       often strike against each other and rebound, and thus communicate
       peculiar motions to each other and the supposed ether, which
       express themselves in luminiferous phenomena. A rise of the
       temperature or an increase in the density of a gas must have an
       influence on the collision of its molecules and luminiferous
       motions thus produced, and this may be the cause of the difference
       of the spectra under these circumstances. It has been shown by
       direct experiment that gases compressed by pressure, when the
       collision of the molecules must be frequent and varied, exhibit a
       more complex spectrum on the passage of an electric spark than
       rarefied gases, and that even a continuous spectrum appears. In
       order to show the variability of the spectrum according to the
       circumstances under which it proceeds, it may be mentioned that
       potassium sulphate fused on a platinum wire gives, on the passage
       of a series of sparks, a distinct system of lines, 583-578, whilst
       when a series of sparks is passed through a solution of this salt
       this system of lines is faint, and when Roscoe and Schuster
       observed the absorption spectrum of the vapour of metallic
       potassium (which is green) they remarked a number of lines of the
       same intensity as the above system in the red, orange, and yellow
       portions.

       [Illustration: FIG. 76.--Method of showing the spectrum of
       substances in solution.]

       The spectra of solutions are best observed by means of Lecoq de
       Boisbaudran's arrangement, shown in fig. 76. A bent capillary
       tube, D F, inside which a platinum wire, A _a_ (from 0·3 to 0·5
       mm. in diameter) is fused, is immersed in a narrow cylinder, C (in
       which it is firmly held by a cork). The projecting end, _a_, of
       the wire is covered by a fine capillary tube, _d_, which extends
       1-2 mm. beyond the wire. Another straight capillary tube, E, with
       a platinum wire, B _b_, about 1 mm. in diameter (a finer wire soon
       becomes hot), is held (by a cork or in a stand) above the end of
       the tube, D. If the wire A be now connected with the positive, and
       the wire B with the negative terminal of a Ruhmkorff's coil (if
       the wires be connected in the opposite order, the spectrum of air
       is obtained), a series of sparks rapidly following each other
       appear between _a_ and _b_, and their light may be examined by
       placing the apparatus in front of the slit of a spectroscope. The
       variations to which a spectrum is liable may easily be observed by
       increasing the distance between the wires, altering the direction
       of the current or strength of the solution, &c.

_Spectrum analysis_ has not only endowed science with a knowledge of
the composition of distant heavenly bodies (of the sun, stars, nebulæ,
comets, &c.), but has also given a new _method_ for studying the matter
of the earth's surface. With its help Bunsen discovered two new elements
belonging to the group of the alkali metals, and thallium, indium, and
gallium were afterwards discovered by the same means. The spectroscope
is employed in the study of rare metals (which in solution often give
distinct absorption spectra), of dyes, and of many organic substances,
&c.[37] With respect to the metals which are analogous to sodium, they
all give similar very volatile salts and such very characteristic
spectra that the least traces of them[38] are discovered with great
ease by means of the spectroscope. For instance, _lithium_ gives a very
brilliant red coloration to a flame and a very bright red spectral line
(wave-length, 670 millionths mm.), which indicates the presence of this
metal in admixture with compounds of other alkali metals.

  [37] The importance of the spectroscope for the purpose of chemical
       research was already shown by Gladstone in 1856, but it did not
       become an accessory to the laboratory until after the discoveries
       of Kirchhoff and Bunsen. It may be hoped that in time
       spectroscopic researches will meet certain wants of the
       theoretical (philosophical) side of chemistry, but as yet all that
       has been done in this respect can only be regarded as attempts
       which have not yet led to any trustworthy conclusions. Thus many
       investigators, by collating the wave-lengths of all the light
       vibrations excited by a given element, endeavour to find the law
       governing their mutual relations; others (especially Hartley and
       Ciamician), by comparing the spectra of analogous elements (for
       instance, chlorine, bromine, and iodine), have succeeded in
       noticing definite features of resemblance in them, whilst others
       (Grünwald) search for relations between the spectra of compounds
       and their component elements, &c.; but--owing to the multiplicity
       of the spectral lines proper to many elements, and (especially in
       the ultra-red and ultra-violet ends of the spectrum) the existence
       of lines which are undistinguishable owing to their faintness, and
       also owing to the comparative novelty of spectroscopic
       research--this subject cannot be considered as in any way
       perfected. Nevertheless, in certain instances there is evidently
       some relationship between the wave-lengths of all the spectral
       lines formed by a given element. Thus, in the hydrogen spectrum
       the wave-length = 364·542 _m_^2/(_m_^{2}-4), if _m_ varies as a
       series of whole numbers from 3 to 15 (Walmer, Hagebach, and
       others). For example, when _m_ = 3, the wave-length of one of the
       brightest lines of the hydrogen spectrum is obtained (656·2), when
       _m_ = 7, one of the visible violet lines (396·8), and when _m_ is
       greater than 9, the ultra-violet lines of the hydrogen spectrum.

  [38] In order to show the degree of sensitiveness of spectroscopic
       reactions the following observation of Dr. Bence Jones may be
       cited: If a solution of 3 grains of a lithium salt be injected
       under the skin of a guinea-pig, after the lapse of four minutes,
       lithium can be discovered in the bile and liquids of the eye, and,
       after ten minutes, in all parts of the animal.

[Illustration: FIG. 77.--Preparation of lithium by the action of a
galvanic current on fused lithium chloride.]

_Lithium_, Li, is, like potassium and sodium, somewhat widely spread
in siliceous rocks, but only occurs in small quantities and as mere
traces in considerable masses of potassium and sodium salts. Only a
very few rather rare minerals contain more than traces of it,[39] for
example, spodumene and lithia mica. Many compounds of lithium are in all
respects closely analogous to the corresponding compounds of sodium
and potassium; but the _carbonate_ is sparingly soluble in cold water,
which fact is taken advantage of for separating lithium from potassium
and sodium. This salt, Li_{2}CO_{3}, is easily converted into the other
compounds of lithium. Thus, for instance, the lithium hydroxide, LiHO,
is obtained in exactly the same way as caustic soda, by the action of
lime on the carbonate, and it is soluble in water and crystallises
(from its solution in alcohol) as LiHO,H_{2}O. Metallic _lithium_ is
obtained by the action of a galvanic current on fused lithium chloride;
for this purpose a cast-iron crucible, furnished with a stout cover,
is filled with lithium chloride, heated until the latter fuses, and a
strong galvanic current is then passed through the molten mass. The
positive pole (fig. 77) consists of a dense carbon rod C (surrounded
by a porcelain tube P fixed in an iron tube BB), and the negative pole
of an iron wire, on which the metal is deposited after the current has
passed through the molten mass for a certain length of time. Chlorine
is evolved at the positive pole. When a somewhat considerable quantity
of the metal has accumulated on the wire it is withdrawn, the metal is
collected from it, and the experiment is then carried on as
before.[39 bis] Lithium is the lightest of all metals, its specific
gravity is 0·59, owing to which it floats even on naphtha; it melts at
180°, but does not volatilise at a red heat. Its appearance recalls that
of sodium, and, like it, it has a yellow tint. At 200° it burns in air
with a very bright flame, forming lithium oxide. In decomposing water it
does not ignite the hydrogen. The characteristic test for lithium
compounds is the _red coloration_ which they impart to a colourless
flame.[40]

  [39] Thus _spodumene_ contains up to 6 p.c. of lithium oxide, and
       _petolite_, and _lepidolite_ or lithia mica, about 3 p.c. of
       lithium oxide. This mica is met with in certain granites in a
       somewhat considerable quantity, and is therefore most frequently
       employed for the preparation of lithium compounds. The treatment
       of lepidolite is carried on on a large scale, because certain
       salts of lithium are employed in medicine as a remedy for certain
       diseases (stone, gouty affections), as they have the power of
       dissolving the insoluble uric acid which is then deposited.
       Lepidolite, which is unacted on by acids in its natural state,
       decomposes under the action of strong hydrochloric acid after it
       has been fused. After being subjected to the action of the
       hydrochloric acid for several hours all the silica is obtained in
       an insoluble form, whilst the metallic oxides pass into solution
       as chlorides. This solution is mixed with nitric acid to convert
       the ferrous salts into ferric, and sodium carbonate is then added
       until the liquid becomes neutral, by which means a precipitate is
       formed of the oxides of iron, alumina, magnesia, &c., as insoluble
       oxides and carbonates. The solution (with an excess of water) then
       contains the chlorides of the alkaline metals KCl, NaCl, LiCl,
       which do not give a precipitate with sodium carbonate in a dilute
       solution. It is then evaporated, and a strong solution of sodium
       carbonate added. This precipitates lithium carbonate, which,
       although soluble in water, is much less so than sodium carbonate,
       and therefore the latter precipitates lithium from strong
       solutions as carbonate, 2LiCl + Na_{2}CO_{3} = 2NaCl +
       Li_{2}CO_{3}. _Lithium carbonate_, which resembles sodium
       carbonate in many respects, is a substance which is very slightly
       soluble in cold water and is only moderately soluble in boiling
       water. In this respect lithium forms a transition between the
       metals of the alkalis and other metals, especially those of the
       alkaline earths (magnesium, barium), whose carbonates are only
       sparingly soluble. Oxide of lithium, Li_{2}O, may be obtained by
       heating lithium carbonate with charcoal. Lithium oxide in
       dissolving gives (per gram-molecule) 26,000 heat units; but the
       combination of Li_{2} with O evolves 140,000 calories--that is,
       more than Na_{2}O (100,000 calories) and K_{2}O (97,000 calories),
       as shown by Beketoff (1887). Oeuvrard (1892) heated lithium to
       redness in nitrogen, and observed the absorption of N and
       formation of Li_{3}N, like Na_{3}N (_see_ Chapter XII. Note 50).

       LiCl, LiBr, and LiI form crystallo-hydrates with H_{2}O, 2H_{2}O,
       and 3H_{2}O. As a rule, LiBr,2H_{2}O crystallises out, but
       Bogorodsky (1894) showed that a solution containing LiBr +
       3·7H_{2}O, cooled to -62°, separates out crystals LiBr,3H_{2}O,
       which decompose at +4° with the separation of H_{2}O. LiF is but
       slightly soluble (in 800 parts) in water (and still less so in a
       solution of NH_{4}F).

  [39 bis] Guntz (1893) recommends adding KCl to the LiCl in preparing Li
       by this method, and to act with a current of 10 ampères at 20
       volts, and not to heat above 450°, so as to avoid the formation of
       Li_{2}Cl.

  [40] In determining the presence of lithium in a given compound, it is
       best to treat the material under investigation with acid (in the
       case of mineral silicon compounds hydrofluoric acid must be
       taken), and to treat the residue with sulphuric acid, evaporate to
       dryness, and extract with alcohol, which dissolves a certain
       amount of the lithium sulphate. It is easy to discover lithium in
       such an alcoholic solution by means of the coloration imparted to
       the flame on burning it, and in case of doubt by investigating its
       light in a spectroscope, because lithium gives a red line, which
       is very characteristic and is found as a dark line in the solar
       spectrum. Lithium was first discovered in 1817 in petolite by
       Arfvedson.

Bunsen in 1860 tried to determine by means of the spectroscope whether
any other as yet unknown metals might not occur in different natural
products together with lithium, potassium, and sodium, and he soon
discovered two new alkali metals showing independent spectra. They are
named after the characteristic coloration which they impart to the
flame. One which gives a red and violet band is named _rubidium_, from
_rubidius_ (dark red), and the other is called _cæsium_, because it
colours a pale flame sky blue, which depends on its containing bright
blue rays, which appear in the spectrum of cæsium as two blue bands
(table on p. 565). Both metals accompany sodium, potassium, and lithium,
but in small quantities; rubidium occurs in larger quantity than cæsium.
The amount of the oxides of cæsium and rubidium in lepidolite does not
generally exceed one-half per cent. Rubidium has also been found in the
ashes of many plants, while the Stassfurt carnallite (the mother-liquor
obtained after having been treated for KCl) forms an abundant source
for rubidium and also partly for cæsium. Rubidium also occurs, although
in very small quantities, in the majority of mineral waters. In a very
few cases cæsium is not accompanied by rubidium; thus, in a certain
granite on the Isle of Elba, cæsium has been discovered, but not
rubidium. This granite contains a very rare mineral called _pollux_,
which contains as much as 34 per cent. of cæsium oxide. Guided by the
spectroscope, and aided by the fact that the double salts of platinic
chloride and rubidium and cæsium chlorides are still less soluble in
water than the corresponding potassium salt, K_{2}PtCl_{6},[41] Bunsen
succeeded in separating both metals from each other and from potassium,
and demonstrated the great resemblance they bear to each other. The
isolated metals,[42] rubidium and cæsium, have respectively the specific
gravities 1·52 and 2·366, and melting points 39° and 27° as N. N.
Beketoff showed (1894), he having obtained cæsium by heating CsAlO_{2}
with Mg([42 bis]).

  [41] The salts of the majority of metals are precipitated as carbonates
       on the addition of ammonium carbonate--for instance, the salts of
       calcium, iron, &c. The alkalis whose carbonates are soluble are
       not, however, precipitated in this case. On evaporating the
       resultant solution and igniting the residue (to remove the
       ammonium salts), we obtain salts of the alkali metals. They may he
       separated by adding hydrochloric acid together with a solution of
       platinic chloride. The chlorides of lithium and sodium give easily
       soluble double salts with platinic chloride, whilst the chlorides
       of potassium, rubidium, and cæsium form double salts which are
       sparingly soluble. A hundred parts of water at 0° dissolve 0·74
       part of the potassium platinochloride; the corresponding rubidium
       platinochloride is only dissolved to the amount of 0·134 part, and
       the cæsium salt, 0·024 part; at 100° 5·13 parts of potassium
       platinochloride, K_{2}PtCl_{6}, are dissolved, 0·634 part of
       rubidium platinochloride, and 0·177 part of cæsium
       platinochloride. From this it is clear how the salts of rubidium
       and cæsium may be isolated. The separation of cæsium from rubidium
       by this method is very tedious. It can be better effected by
       taking advantage of the difference of the solubility of their
       carbonates in alcohol; cæsium carbonate, Cs_{2}CO_{3}, is soluble
       in alcohol, whilst the corresponding salts of rubidium and
       potassium are almost insoluble. Setterberg separated these metals
       as alums, but the best method, that given by Scharples, is founded
       on the fact that from a mixture of the chlorides of potassium,
       sodium, cæsium, and rubidium in the presence of hydrochloric acid,
       stannic chloride precipitates a double salt of cæsium, which is
       very slightly soluble. The salts of Rb and Cs are closely
       analogous to those of potassium.

  [42] Bunsen obtained rubidium by distilling a mixture of the tartrate
       with soot, and Beketoff (1888) by heating the hydroxide with
       aluminium, 2RbHO + Al = RbAlO_{2} + H_{2} + Rb. By the action of
       85 grams of rubidium on water, 94,000 heat units are evolved.
       Setterberg obtained cæsium (1882) by the electrolysis of a fused
       mixture of cyanide of cæsium and of barium. Winkler (1890) showed
       that metallic magnesium reduces the hydrates and carbonates of Rb
       and Cs like the other alkaline metals. N. N. Beketoff obtained
       them with aluminium (see following note).

  [42 bis] Beketoff (1888) showed that metallic aluminium reduces the
       hydrates of the alkaline metals at a red heat (they should be
       perfectly dry) with the formation of aluminates (Chapter XVII.),
       RAlO_{2}--for example, 2KHO + Al = KAlO_{2} + K + H_{2}. It is
       evident that in this case only half of the alkaline metal is
       obtained free. On the other hand, K. Winkler (1889) showed that
       magnesium powder is also able to reduce the alkaline metals from
       their hydrates and carbonates. N. N. Beketoff and Tscherbacheff
       (1894) prepared cæsium upon this principle by heating its
       aluminate CsAlO_{2} with magnesium powder. In this case aluminate
       of magnesium is formed, and the whole of the cæsium is obtained as
       metal: 2CsAlO_{2} + Mg = MgOAl_{2}O_{5} + 2Cs. A certain excess of
       alumina was taken (in order to obtain a less hygroscopic mass of
       aluminate), and magnesium powder (in order to decompose the last
       traces of water); the CsAlO_{2} was prepared by the precipitation
       of cæsium alums by caustic baryta, and evaporating the resultant
       solution. We may add that N. N. Beketoff (1887) prepared oxide of
       potassium, K_{2}O, by heating the peroxide, KO, in the vapour of
       potassium (disengaged from its alloy with silver), and showed that
       in dissolving in an excess of water it evolves (for the
       above-given molecular weight) 67,400 calories (while 2KHO in
       dissolving in water evolves 24,920 cal.; so that K_{2}O + H_{2}O
       gives 42,480 cal.), whence (knowing that K_{2} + O+H_{2}O in an
       excess of water evolves 164,500) it follows that K_{2} + O evolves
       97,100 cal. This quantity is somewhat less than that (100,260
       cal.) which corresponds to sodium, and the energy of the action of
       potassium upon water is explained by the fact that K_{2}O evolves
       more heat than Na_{2}O in combining with water (_see_ Chapter II.
       Note 9). Just as hydrogen displaces half the Na from Na_{2}O
       forming NaHO, so also N. N. Beketoff found from experiment and
       thermochemical reasonings that hydrogen displaces half the
       potassium from K_{2}O forming KHO and evolving 7,190 calories.
       Oxide of lithium, Li_{2}O, which is easily formed by igniting
       Li_{2}CO_{3} with carbon (when Li_{2}O + 2CO is formed),
       disengages 26,000 cals. with an excess of water, while the
       reaction Li_{2} + O gives 114,000 cals. and the reaction Li_{2} +
       H_{2}O gives only 13,000 cals., and metallic lithium cannot be
       liberated from oxide of lithium with hydrogen (nor with carbon).
       Thus in the series Li, Na, K, the formation of R_{2}O gives most
       heat with Li and least with K, while the formation of RCl evolves
       most heat with K (105,000 cals.) and least of all with Li (93,500
       cals.). Rubidium, in forming Rb_{2}O, gives 94,000 cals.
       (Beketoff). Cæsium, in acting upon an excess of water, evolves
       51,500 cals., and the reaction Cs_{2} + O evolves about 100,000
       cals.--_i.e._ more than K and Rb, and almost as much as Na--and
       oxide of cæsium reacts with hydrogen (according to the equation
       Cs_{2}O + H = CsHO + Cs) more easily than any of the oxides of the
       alkali metals, and this reaction takes place at the ordinary
       temperature (the hydrogen is absorbed), as Beketoff showed (1893).
       He also obtained a mixed oxide, AgCsO, which was easily formed in
       the presence of silver, and absorbed hydrogen with the formation
       of CsHO.

Judging by the properties of the free metals, and of their corresponding
and even very complex compounds, lithium, sodium, potassium, rubidium,
and cæsium present an indubitable chemical resemblance. The fact that
the metals easily decompose water, and that their hydroxides RHO and
carbonates R_{2}CO_{3} are soluble in water, whilst the hydroxides
and carbonates of nearly all other metals are insoluble, shows that
these metals form a natural group of _alkali metals_. The halogens and
the alkali metals form, by their character, the two extremes of the
elements. Many of the other elements are metals approaching the alkali
metals, both in their capacity of forming salts and in not forming acid
compounds, but are not so energetic as the alkali metals, that is, they
form less energetic bases. Such are the common metals, silver, iron,
copper, &c. Some other elements, in the character of their compounds,
approach the halogens, and, like them, combine with hydrogen, but these
compounds do not show the energetic property of the halogen acids; in
a free state they easily combine with metals, but they do not then
form such saline compounds as the halogens do--in a word, the halogen
properties are less sharply defined in them than in the halogens
themselves. Sulphur, phosphorus, arsenic, &c. belong to this order of
elements. The clearest distinction of the properties of the halogens
and alkali metals is expressed in the fact that the former give acids
and do not form bases, whilst the latter, on the contrary, only give
bases. The first are true _acid elements_, the latter clearly-defined
_basic or metallic elements_. On combining together, the halogens form,
in a chemical sense, unstable compounds, and the alkali metals alloys
in which the character of the metals remains unaltered, just as in the
compound ICl the character of the halogens remains undisguised; thus
both classes of elements on combining with members of their own class
form non-characteristic compounds, which have the properties of their
components. On the other hand, the halogens on combining with the alkali
metals form compounds which are, in all respects, stable, and in which
the original characters of the halogens and alkali metals have entirely
disappeared. The formation of such compounds is accompanied by evolution
of a large amount of heat, and by an entire change of both the physical
and chemical properties of the substances originally taken. The alloy of
sodium and potassium, although liquid at the ordinary temperature, is
perfectly metallic, like both its components. The compound of sodium and
chlorine has neither the appearance nor the properties of the original
elements; sodium chloride melts at a higher temperature, and is more
difficultly volatile, than either sodium or chlorine.

With all these qualitative differences there is, however, an important
quantitative _resemblance between the halogens and the alkali metals_.
This resemblance is clearly expressed by stating that both orders of
elements belong to those which are univalent with respect to hydrogen.
It is thus correct to say that both the above-named orders of elements
replace hydrogen atom for atom. Chlorine is able to take the place of
hydrogen by metalepsis, and the alkali metals take the place of hydrogen
in water and acids. As it is possible to consecutively replace every
equivalent of hydrogen in a hydrocarbon by chlorine, so it is possible
in an acid containing several equivalents of hydrogen to replace the
hydrogen consecutively equivalent after equivalent by an alkali metal;
hence an atom of these elements is analogous to an atom of hydrogen,
which is taken, in all cases, as the unit for the comparison of the
other elements. In ammonia, and in water, chlorine and sodium are
able to bring about a direct replacement. According to the law of
substitution, the formation of sodium chloride, NaCl, at once shows the
equivalence of the atoms of the alkali metals and the halogens. The
halogens and hydrogen and the alkali metals combine with such elements
as oxygen, and it is easily proved that in such compounds one atom
of oxygen is able to retain two atoms of the halogens, of hydrogen,
and of the alkali metals. For this purpose it is enough to compare
the compounds KHO, K_{2}O, HClO, and Cl_{2}O, with water. It must not
be forgotten, however, that the halogens give, with oxygen, besides
compounds of the type R_{2}O, higher acid grades of oxidation, which the
alkali metals and hydrogen are not capable of forming. We shall soon
see that these relations are also subject to a special law, showing a
gradual transition of the properties of the elements from the alkali
metals to the halogens.[43]

  [43] We may here observe that the halogens, and especially iodine,
       may play the part of metals (hence iodine is more easily replaced
       by metals than the other halogens, and it approaches nearer to the
       metals in its physical properties than the other halogens).
       Schützenberger obtained a compound C_{2}H_{3}O(OCl), which he
       called chlorine acetate, by acting on acetic anhydride,
       (C_{2}H_{3}O)_{2}O, with chlorine monoxide, Cl_{2}O. With iodine
       this compound gives off chlorine and forms iodine acetate,
       C_{2}H_{3}O(OI), which also is formed by the action of iodine
       chloride on sodium acetate, C_{2}H_{3}O(ONa). These compounds are
       evidently nothing else than mixed anhydrides of hypochlorous and
       hypoiodous acids, or the products of the substitution of hydrogen
       in RHO by a halogen (_see_ Chapter XI., Notes 29 and 78 bis). Such
       compounds are very unstable, decompose with an explosion when
       heated, and are changed by the action of water and of many other
       reagents, which is in accordance with the fact that they contain
       very closely allied elements, as does Cl_{2}O itself, or ICl or
       KNa. By the action of chlorine monoxide on a mixture of iodine and
       acetic anhydride, Schützenberger also obtained the compound
       I(C_{2}H_{3}O_{2})_{3}, which is analogous to ICl_{3}, because the
       group C_{2}H_{3}O_{2} is, like Cl, a halogen, forming salts with
       the metals. Similar properties are found in iodosobenzene (Chapter
       XI., Note 79).

The atomic weights of the alkali metals, lithium 7, sodium 23, potassium
39, rubidium 85, and cæsium 133, show that here, as in the class of
halogens, the elements may be arranged according to their atomic weights
in order to compare the properties of the analogous compounds of the
members of this group. Thus, for example, the platinochlorides of lithium
and sodium are soluble in water; those of potassium, rubidium, and
cæsium sparingly soluble, and the greater the atomic weight of the metal
the less soluble is the salt. In other cases the reverse is observed--the
greater the atomic weight the more soluble are the corresponding salts.
The variation of properties with the variation in atomic weights even
shows itself in the metals themselves; thus lithium volatilises with
difficulty, whilst sodium is obtained by distillation, potassium
volatilises more easily than sodium, and rubidium and cæsium as we have
seen, are still more volatile.




                               CHAPTER XIV

    THE VALENCY AND SPECIFIC HEAT OF THE METALS. MAGNESIUM. CALCIUM,
                    STRONTIUM, BARIUM, AND BERYLLIUM


It is easy by investigating the composition of corresponding compounds,
to establish the _equivalent weights_ of the metals compared with
hydrogen--that is, the quantity which replaces one part by weight of
hydrogen. If a metal decomposes acids directly, with the evolution of
hydrogen, the equivalent weight of the metal may be determined by taking
a definite weight of it and measuring the volume of hydrogen evolved by
its action on an excess of acid; it is then easy to calculate the weight
of the hydrogen from its volume.[1] The same result may be arrived at
by determining the composition of the normal salts of the metal; for
instance, by finding the weight of metal which combines with 35·5 parts
of chlorine or 80 parts of bromine.[2] The equivalent of a metal may be
also ascertained by simultaneously (_i.e._ in one circuit) decomposing
an acid and a fused salt of a given metal by an electric current
and determining the relation between the amounts of hydrogen and
metal separated, because, according to Faraday's law, electrolytes
(conductors of the second order) are always decomposed in equivalent
quantities.[2 bis] The equivalent of a metal may even be found by simply
determining the relation between its weight and that of its salt-giving
oxide, as by this we know the quantity of the metal which combines with 8
parts by weight of oxygen, and this will be the equivalent, because 8
parts of oxygen combine with 1 part by weight of hydrogen. One method is
verified by another, and all the processes for the accurate determination
of equivalents require the greatest care to avoid the absorption of
moisture, further oxidation, volatility, and other accidental influences
which affect exact weighings. The description of the methods necessary
for the attainment of exact results belongs to the province of analytical
chemistry.

  [1] Under favourable circumstances (by taking all the requisite
      precautions), the weight of the equivalent may be accurately
      determined by this method. Thus Reynolds and Ramsay (1887)
      determined the equivalent of zinc to be 32·7 by this method (from
      the average of 29 experiments), whilst by other methods it has been
      fixed (by different observers) between 32·55 and 33·95.

      The differences in their equivalents may be demonstrated by taking
      equal weights of different metals, and collecting the hydrogen
      evolved by them (under the action of an acid or alkali).

  [2] The most accurate determinations of this kind were carried on by
      Stas, and will be described in Chapter XXIV.

  [2 bis] The amount of electricity in one coulomb according to the
      present nomenclature of electrical units (_see_ Works on Physics
      and Electro-technology) disengages 0·00001036 gram of hydrogen,
      0·00112 gram of silver, 0·0003263 gram of copper from the salts of
      the oxide, and 0·0006526 gram from the salts of the suboxide, &c.
      These amounts stand in the same ratio as the equivalents, _i.e._ as
      the quantities replaced by one part by weight of hydrogen. The
      intimate bond which is becoming more and more marked existing
      between the electrolytic and purely chemical relations of
      substances (especially in solutions) and the application of
      electrolysis to the preparation of numerous substances on a large
      scale, together with the employment of electricity for obtaining
      high temperatures, &c., makes me regret that the plan and
      dimensions of this book, and the impossibility of giving a concise
      and objective exposition of the necessary electrical facts, prevent
      my entering upon this province of knowledge, although I consider it
      my duty to recommend its study to all those who desire to take part
      in the further development of our science.

      There is only one side of the subject respecting the direct
      correlation between thermochemical data and electro-motive force,
      which I think right to mention here, as it justifies the general
      conception, enunciated by Faraday, that the galvanic current is an
      aspect of the transference of chemical motion or reaction along the
      conductors.

      From experiments conducted by Favre, Thomsen, Garni, Berthelot,
      Cheltzoff, and others, upon the amount of heat evolved in a closed
      circuit, it follows that the electro-motive force of the current or
      its capacity to do a certain work, E, is proportional to the whole
      amount of heat, Q, disengaged by the reaction forming the source of
      the current. If E be expressed in volts, and Q in thousands of
      units of heat referred to equivalent weights, then E = 0·0436Q. For
      example in a Daniells battery E = 1·09 both by experiment and
      theory, because in it there takes place the decomposition of
      CuSO_{4} into Cu + O together with the formation of Zn + O and ZnO
      + SO_{3}Aq, and these reactions correspond to Q = 25·06 thousand
      units of heat. So also in all other primary batteries (_e.g._
      Bunsen's, Poggendorff's, &c.) and secondary ones (for instance,
      those acting according to the reaction Pb + H_{2}SO_{4} + PbO_{2},
      as Cheltzoff showed) E = 0·0436Q.

For univalent metals, like those of the alkalis, the weight of the
equivalent is equal to the weight of the atom. For bivalent metals the
atomic weight is equal to the weight of two equivalents, for _n_-valent
metals it is equal to the weight of _n_ equivalents. Thus aluminium,
Al = 27, is trivalent, that is, its equivalent = 9; magnesium, Mg =
24, is bivalent, and its equivalent = 12. Therefore, if potassium or
sodium, or in general a univalent metal, M, give compounds M_{2}O, MHO,
MCl, MNO_{3}, M_{2}SO_{4}, &c., and in general MX, then for bivalent
metals like magnesium or calcium the corresponding compounds will be
MgO, Mg(HO)_{2}, MgCl_{2}, Mg(NO_{3})_{2}, MgSO_{4}, &c., or in general
MX_{2}.

By what are we to be guided in ascribing to some metals univalency and
to others bi-, ter-, quadri-, ... _n_-valency? What obliges us to make
this difference? Why are not all metals given the same valency--for
instance, why is not magnesium considered as univalent? If this be
done, taking Mg = 12 (and not 24 as now), not only is a simplicity
of expression of the composition of all the compounds of magnesium
attained, but we also gain the advantage that their composition will
be the same as those of the corresponding compounds of sodium and
potassium. These combinations were so expressed formerly--why has this
since been changed?

These questions could only be answered after the establishment of the
idea of multiples of the atomic weights as the minimum quantities of
certain elements combining with others to form compounds--in a word,
since the time of the establishment of Avogadro-Gerhardt's law (Chapter
VII.). By taking such an element as arsenic, which has many volatile
compounds, it is easy to determine the density of these compounds, and
therefore to establish their molecular weights, and hence to find the
indubitable atomic weight, exactly as for oxygen, nitrogen, chlorine,
carbon, &c. It appears that As = 75, and its compounds correspond, like
the compounds of nitrogen, with the forms AsX_{3}, and AsX_{5}; for
example, AsH_{3}, AsCl_{3}, AsF_{5}, As_{2}O_{5}, &c. It is evident that
we are here dealing with a metal (or rather element) of two valencies,
which moreover is never univalent, but tri- or quinqui-valent. This
example alone is sufficient for the recognition of the existence of
polyvalent atoms among the metals. And as antimony and bismuth are
closely analogous to arsenic in all their compounds, (just as potassium
is analogous to rubidium and cæsium); so, although very few volatile
compounds of bismuth are known, it was necessary to ascribe to them
formulæ corresponding with those ascribed to arsenic.

As we shall see in describing them, there are also many analogous metals
among the bivalent elements, some of which also give volatile compounds.
For example, zinc, which is itself volatile, gives several volatile
compounds (for instance, zinc ethyl, ZnC_{4}H_{10}, which boils at 118°,
vapour density = 61·3), and in the molecules of all these compounds
there is never less than 65 parts of zinc, which is equivalent to H_{2},
because 65 parts of zinc displace 2 parts by weight of hydrogen; so
that zinc is just such an example of the bivalent metals as oxygen,
whose equivalent = 8 (because H_{2} is replaced by O = 16), is a
representative of the bivalent elements, or as arsenic is of the tri-
and quinqui-valent elements. And, as we shall afterwards see, magnesium
is in many respects closely analogous to zinc, which fact obliges us to
regard magnesium as a bivalent metal.

Such metals as mercury and copper, which are able to give not one but
two bases, are of particular importance for distinguishing univalent and
bivalent metals. Thus copper gives the suboxide Cu_{2}O and the oxide
CuO--that is, the compounds CuX corresponding with the suboxide are
analogous (in the quantitative relations, by their composition) to NaX
or AgX, and the compounds of the oxide CuX_{2}, to MgX_{2}, ZnX_{2}, and
in general to the bivalent metals. It is clear that in such examples we
must make a distinction between atomic weights and equivalents.

In this manner the valency, that is, the number of equivalents entering
into the atom of the metals may in many cases be established by means
of comparatively few volatile metallic compounds, with the aid of a
search into their analogies (concerning which see Chapter XV.). _The law
of specific heats_ discovered by Dulong and Petit has frequently been
applied to the same purpose[3] in the history of chemistry, especially
since the development given to this law by the researches of Regnault,
and since Cannizzaro (1860) showed the agreement between the deductions
of this law and the consequences arising from Avogadro-Gerhardt's law.

  [3] The chief means by which we determine the valency of the elements,
      or what multiple of the equivalent should be ascribed to the atom,
      are: (1) The law of Avogadro-Gerhardt. This method is the most
      general and trustworthy, and has already been applied to a great
      number of elements. (2) The different grades of oxidation and their
      isomorphism or analogy in general; for example, Fe = 56 because the
      suboxide (ferrous oxide) is isomorphous with magnesium oxide, &c.,
      and the oxide (ferric oxide) contains half as much oxygen again as
      the suboxide. Berzelius, Marignac, and others took advantage of
      this method for determining the composition of the compounds of
      many elements. (3) The specific heat, according to Dulong and
      Petit's law. Regnault, and more especially Cannizzaro, used this
      method to distinguish univalent from bivalent metals. (4) The
      periodic law (_see_ Chapter XV.) has served as a means for the
      determination of the atomic weights of cerium, uranium, yttrium,
      &c., and more especially of gallium, scandium, and germanium. The
      correction of the results of one method by those of others is
      generally had recourse to, and is quite necessary, because,
      phenomena of dissociation, polymerisation, &c., may complicate the
      individual determinations by each method.

      It will be well to observe that a number of other methods,
      especially from the province of those physical properties which are
      clearly dependent on the magnitude of the atom (or equivalent) or
      of the molecule, may lead to the same result. I may point out, for
      instance, that even the specific gravity of solutions of the
      metallic chlorides may serve for this purpose. Thus, if beryllium
      he taken as trivalent--that is, if the composition of its chloride
      be taken as BeCl_{3} (or a polymeride of it), then the specific
      gravity of solutions of beryllium chloride will not fit into the
      series of the other metallic chlorides. But by ascribing to it an
      atomic weight Be = 7, or taking Be as bivalent, and the composition
      of its chloride as BeCl_{2}, we arrive at the general rule given in
      Chapter VII., Note 28. Thus W. G. Burdakoff determined in my
      laboratory that the specific gravity at 15°/4° of the solution
      BeCl_{2} + 200H_{2}O = 1·0138--that is, greater than the
      corresponding solution KCl + 200H_{2}O (= 1·0121), and less than
      the solution MgCl_{2} + 200H_{2}O (= 1·0203), as would follow from
      the magnitude of the molecular weight BeCl_{2} = 80, since KCl =
      74·5 and MgCl_{2} = 95.


Dulong and Petit, having determined the specific heat of a number of
solid elementary substances, observed that as the atomic weights of the
elements increase, their specific heats decrease, and that _the product
of the specific heat Q into the atomic weight A is an almost constant
quantity_. This means that to bring different elements into a known
thermal state an equal amount of work is required if atomic quantities
of the elements are taken; that is, the amounts of heat expended in
heating equal quantities by weight of the elements are far from equal,
but are in inverse proportion to the atomic weights. For thermal
changes the atom is a unit; all atoms, notwithstanding the difference
of weight and nature, are equal. This is the simplest expression of the
fact discovered by Dulong and Petit. The specific heat measures that
quantity of heat which is required to raise the temperature of _one
unit of weight_ of a substance by one degree. If the magnitude of the
specific heat of elements be multiplied by the atomic weight, then we
obtain the atomic heat--that is, the amount of heat required to raise
the temperature of the atomic weight of an element by one degree. It
is these products which for the majority of the elements prove to be
approximately, if not quite, identical. A complete identity cannot be
expected, because the specific heat of one and the same substance varies
with the temperature, with its passage from one state into another, and
frequently with even a simple mechanical change of density (for instance
by hammering), not to speak of allotropic changes, &c. We will cite
several figures[4] proving the truth of the conclusions arrived at by
Dulong and Petit with respect to solid elementary bodies.

                          Li         Na        Mg        P
                A =   7         23         24        31
                Q =   0·9408     0·2934     0·245     0·202
               AQ =   6·59       6·75       5·88      6·26

                         Fe         Cu         Zn         Br
                A =  56         63         65        80
                Q =   0·112      0·093      0·093     0·0843
               AQ =   6·27       5·86       6·04      6·74

                          Pd        Ag         Sn         I
                A = 106        108        118       127
                Q =   0·0592     0·056      0·055     0·541
               AQ =   6·28       6·05       6·49      6·87

                          Pt         Au         Hg        Pb
                A = 196        198        200       206
                Q =   0·0325     0·0324     0·0333    0·0315
               AQ =   6·37       6·41       6·66      6·49

  [4] The specific heats here given refer to different limits of
      temperature, but in the majority of cases between 0° and 100°; only
      in the case of bromine the specific heat is taken (for the solid
      state) at a temperature below -7°, according to Regnault's
      determination. _The variation of the specific heat with a change of
      temperature_ is a very complex phenomenon, the consideration of
      which I think would here be out of place. I will only cite a few
      figures as an example. According to Bystrom, the specific heat of
      iron at 0° = 0·1116, at 100° = 0·1114, at 200° = 0·1188, at 300° =
      0·1267, and at 1,400° = 0·4031. Between these last limits of
      temperature a change takes place in iron (a spontaneous heating,
      _recalescence_), as we shall see in Chapter XXII. For quartz
      SiO_{2} Pionchon gives Q = 0·1737 + 394_t_10^{-6}-27_t_^{2}10^{-9}
      up to 400°, for metallic aluminium (Richards, 1892) at 0° 0·222, at
      20° 0·224, at 100° 0·232; consequently, as a rule, the specific
      heat varies slightly with the temperature. Still more remarkable
      are H. E. Weber's observations on the great variation of the
      specific heat of charcoal, the diamond and boron:

                               0°   100°  200°  600°  900°
               Wood charcoal  0·15  0·23  0·29  0·44  0·46
               Diamond        0·10  0·19  0·22  0·44  0·45
               Boron          0·22  0·29  0·35    --    --

      These determinations, which have been verified by Dewar, Le
      Chatelier (Chapter VIII., Note 13), Moissan, and Gauthier, the
      latter finding for boron AQ = 6 at 400°, are of especial importance
      as confirming the universality of Dulong and Petit's law, because
      the elements mentioned above form exceptions to the general rule
      when the mean specific heat is taken for temperatures between 0°
      and 100°. Thus in the case of the diamond the product of A × Q at
      0° = 1·2, and for boron = 2·4. But if we take the specific heat
      towards which there is evidently a tendency with a rise of
      temperature, we obtain a product approaching to 6 as with other
      elements. Thus with the diamond and charcoal, it is evident that
      the specific heat tends towards 0·47, which multiplied by 12 gives
      5·6, the same as for magnesium and aluminium. I may here direct the
      reader's attention to the fact that for solid elements having a
      small atomic weight, the specific heat varies considerably if we
      take the average figures for temperatures 0° to 100°:

                      Li = 7    Be = 9    B = 11    C = 12
                  Q =  0·94      0·42      0·24      0·20
                 AQ =  6·6       3·8       2·6       2·4

      It is therefore clear that the specific heat of beryllium
      determined at a low temperature cannot serve for establishing its
      atomicity. On the other hand, the low atomic heat of charcoal,
      graphite, and the diamond, boron, &c., may perhaps depend on the
      complexity of the molecules of these elements. The necessity for
      acknowledging a great complexity of the molecules of carbon was
      explained in Chapter VIII. In the case of sulphur the molecule
      contains at least S_{6} and its atomic heat = 32 × 0·163 = 5·22,
      which is distinctly below the normal. If a large number of atoms of
      carbon are contained in the molecule of charcoal, this would to a
      certain extent account for its comparatively small atomic heat.
      With respect to the specific heat of compounds, it will not be out
      of place to mention here the conclusion arrived at by Kopp, that
      the molecular heat (that is, the product of MQ) may be looked on as
      the sum of the atomic heats of its component elements; but as this
      rule is not a general one, and can only be applied to give an
      approximate estimate of the specific heats of substances, I do not
      think it necessary to go into the details of the conclusions
      described in Liebig's 'Annalen Supplement-Band,' 1864, which
      includes a number of determinations made by Kopp.

It is seen from this that the product of the specific heat of the element
into the atomic weight is an almost constant quantity, which is nearly 6.
Hence it is possible to determine the valency by the specific heats of
the metals. Thus, for instance, the specific heats of lithium, sodium,
and potassium convince us of the fact that their atomic weights are
indeed those which we chose, because by multiplying the specific heats
found by experiment by the corresponding atomic weights we obtain the
following figures: Li, 6·59, Na, 6·75 and K, 6·47. Of the alkaline earth
metals the specific heats have been determined: of magnesium = 0·245
(Regnault and Kopp), of calcium = 0·170 (Bunsen), and of barium = 0·05
(Mendeléeff). If the same composition be ascribed to the compounds of
magnesium as to the corresponding compounds of potassium, then the
equivalent of magnesium will be equal to 12. On multiplying this atomic
weight by the specific heat of magnesium, we obtain a figure 2·94, which
is half that which is given by the other solid elements and therefore the
atomic weight of magnesium must be taken as equal to 24 and not to 12.
Then the atomic heat of magnesium = 24 × 0·245 = 5·9; for calcium, giving
its compounds a composition CaX_{2}--for example CaCl_{2}, CaSO_{4}, CaO
(Ca = 40)--we obtain an atomic heat = 40 × 0·17 = 6·8, and for barium it
is equal to 137 × 0·05 = 6·8; that is, they must be counted as bivalent,
or that their atom replaces H_{2}, Na_{2}, or K_{2}. This conclusion may
be confirmed by a method of analogy, as we shall afterwards see. The
application of the principle of specific heats to the determination of
the magnitudes of the atomic weights of those metals, the magnitude of
whose atomic weights could not be determined by Avogadro-Gerhardt's law,
was made about 1860 by the Italian professor Cannizzaro.

Exactly the same conclusions respecting the bivalence of magnesium and
its analogues are obtained by comparing the specific heats of their
compounds, especially of the halogen compounds as the most simple,
with the specific heats of the corresponding alkali compounds. Thus,
for instance, the specific heats of magnesium and calcium chlorides,
MgCl_{2} and CaCl_{2}, are 0·194 and 0·164, and of sodium and potassium
chlorides, NaCl and KCl, 0·214 and 0·172, and therefore their molecular
heats (or the products QM, where M is the weight of the molecule)
are 18·4 and 18·2, 12·5 and 12·8, and hence the atomic heats (or the
quotient of QM by the number of atoms) are all nearly 6, as with the
elements. Whilst if, instead of the actual atomic weights Mg = 24 and
Ca = 40, their equivalents 12 and 20 be taken, then the atomic heats of
the chlorides of magnesium and calcium would be about 4·6, whilst those
of potassium and sodium chlorides are about 6·3.[5] We must remark,
however, that as the specific heat or the amount of heat required to
raise the temperature of a unit of weight one degree[6] is a complex
quantity--including not only the increase of the energy of a substance
with its rise in temperature, but also the external work of expansion[7]
and the internal work accomplished in the molecules causing them to
decompose according to the rise of temperature[8]--therefore it is
impossible to expect in the magnitude of the specific heat the great
simplicity of relation to composition which we see, for instance, in
the density of gaseous substances. Hence, although the specific heat
is one of the important means for determining the atomicity of the
elements, still the mainstay for a true judgment of atomicity is only
given by Avogadro-Gerhardt's law, _i.e._ this other method can only be
accessory or preliminary, and when possible recourse should be had to
the determination of the vapour density.

  [5] It must be remarked that in the case of oxygen (and also hydrogen
      and carbon) compounds the quotient of MQ/_n_, where _n_ is the
      number of atoms in the molecule, is always less than 6 for solids;
      for example, for MgO = 5·0, CuO = 5·1, MnO_{2} = 4·6, ice (Q =
      0·504) = 3, SiO_{2} = 3·5, &c. At present it is impossible to say
      whether this depends on the smaller specific heat of the atom of
      oxygen in its solid compounds (Kopp, Note 4) or on some other
      cause; but, nevertheless, taking into account this decrease
      depending on the presence of oxygen, a reflection of the atomicity
      of the elements may to a certain extent be seen in the specific
      heat of the oxides. Thus for alumina, Al_{2}O_{3} (Q = 0·217), MQ =
      22·3, and therefore the quotient MQ/_n_ = 4·5, which is nearly that
      given by magnesium oxide, MgO. But if we ascribe the same
      composition to alumina, as to magnesia--that is, if aluminium were
      counted as divalent--we should obtain the figure 3·7, which is much
      less. In general, in compounds of identical atomic composition and
      of analogous chemical properties the molecular heats MQ are nearly
      equal, as many investigators have long remarked. For example, ZnS =
      11·7 and HgS = 11·8; MgSO_{4} = 27·0 and ZnSO_{4} = 28·0, &c.

  [6] If W be the amount of heat contained in a mass _m_ of a substance
      at a temperature _t_, and _d_W the amount expended in heating it
      from _t_ to _t_ + _dt_, then the specific heat Q = _d_W(_m_ ×
      _dt_). The specific heat not only varies with the composition and
      complexity of the molecules of a substance, but also with the
      temperature, pressure, and physical state of a substance. Even for
      gases the variation of Q with _t_ is to be observed. Thus it is
      seen from the experiments of Regnault and Wiedemann that the
      specific heat of carbonic anhydride at 0° = 0·19, at 100° = 0·22,
      and at 200° = 0·24. But the variation of the specific heat of
      permanent gases with the temperature is, as far as we know, very
      inconsiderable. According to Mallard and Le Chatelier it is =
      0·0006/M per 1°, where M is the molecular weight (for instance, for
      O_{2}, M = 32). Therefore the specific heat of those permanent
      gases which contain two atoms in the molecule (H_{2}, O_{2}, N_{2},
      CO, and NO) may be, as is shown by experiment, taken as not varying
      with the temperature. The constancy of the specific heat of perfect
      gases forms one of the fundamental propositions of the whole theory
      of heat and on it depends the determination of temperatures by
      means of gas-thermometers containing hydrogen, nitrogen, or air. Le
      Chatelier (1887), on the basis of existing determinations,
      concludes that the molecular heat--that is, the product MQ--of all
      gases varies in proportion to the temperature, and tends to become
      equal (= 6·8) at the temperature of absolute zero (that is,
      at -273°); and therefore MQ = 6·8 + _a_(273 + _t_), where _a_ is a
      constant quantity which increases with the complexity of the
      gaseous molecule and Q is the specific heat of the gas under a
      constant pressure. For permanent gases _a_ almost = 0, and
      therefore MQ = 6·8--that is, the atomic heat (if the molecule
      contains two atoms) = 3·4, as it is in fact (Chapter IX., Note
      17 bis). As regards liquids (as well as the vapours formed by
      them), the specific heat always rises with the temperature. Thus
      for benzene it equals 0·38 + 0·0014_t_. R. Schiff (1887) showed
      that the variation of the specific heat of many organic liquids is
      proportional to the change of temperature (as in the case of gases,
      according to Le Chatelier), and reduced these variations into
      dependence with their composition and absolute boiling point. It is
      very probable that the theory of liquids will make use of these
      simple relations which recall the simplicity of the variation of
      the specific gravity (Chapter II., Note 34), cohesion, and other
      properties of liquids with the temperature. They are all expressed
      by the linear function of the temperature, _a_ + _bt_, with the
      same degree of proximity as the property of gases is expressed by
      the equation _pv_ = _Rt_.

      As regards the relation between the specific heats of liquids (or
      of solids) and of their vapours, the specific heat of the vapour
      (and also of the solid) is always less than that of the liquid. For
      example, benzene vapour 0·22, liquid 0·38; chloroform vapour 0·13,
      liquid 0·23; steam 0·475, liquid water 1·0. But the complexity of
      the relations existing in specific heat is seen from the fact that
      the specific heat of ice = 0·502 is less than that of liquid water.
      According to Regnault, in the case of bromine the specific heat of
      the vapour = 0·055 at (150°), of the liquid = 0·107 (at 30°), and
      of solid bromine = 0·084 (at -15°). The specific heat of solid
      benzoic acid (according to experiment and calculation, Hess, 1888)
      between 0° and 100° is 0·31, and of liquid benzoic acid 0·50. One
      of the problems of the present day is the explanation of those
      complex relations which exist between the composition and such
      properties as specific heat, latent heat, expansion by heat,
      compression, internal friction, cohesion, and so forth. They can
      only be connected by a complete theory of liquids, which may now
      soon be expected, more especially as many sides of the subject have
      already been partially explained.

  [7] According to the above reasons the quantity of heat, Q, required to
      raise the temperature of one part by weight of a substance by one
      degree may be expressed by the sum Q = K + B + D, where K is the
      heat actually expended in heating the substance, or what is termed
      the absolute specific heat, B the amount of heat expended in the
      internal work accomplished with the rise of temperature, and D the
      amount of heat expended in external work. In the case of gases the
      last quantity may be easily determined, knowing their coefficient
      of expansion, which is approximately = 0·00368. By applying to this
      case the same argument given at the end of Note 11, Chapter I., we
      find that one cubic metre of a gas heated 1° produces an external
      work of 10333 × 0·00368, or 38·02 kilogrammetres, on which
      38·02/424 or 0·0897 heat units are expended. This is the heat
      expended for the external work produced by one cubic metre of a
      gas, but the specific heat refers to units of weight, and therefore
      it is necessary in order to know D to reduce the above quantity to
      a unit of weight. One cubic metre of hydrogen at 0° and 760 mm.
      pressure weighs 0·0896 kilo, a gas of molecular weight M has a
      density M/2, consequently a cubic metre weighs (at 0° and 760 mm.)
      0·0448M kilo, and therefore 1 kilogram of the gas occupies a volume
      1/0·0448M cubic metres, and hence the external work D in the
      heating of 1 kilo of the given gas through 1° = 0·0896/0·0448M, or
      D = 2/M.

      Taking the magnitude of the internal work B for gases as negligible
      if permanent gases are taken, and therefore supposing B = 0, we
      find the specific heat of gases at a constant pressure Q = K + 2 M,
      where K is the specific heat at a constant volume, or the true
      specific heat, and M the molecular weight. Hence K = Q-2/M. The
      magnitude of the specific heat Q is given by direct experiment.
      According to Regnault's experiments, for oxygen it = 0·2175, for
      hydrogen 3·405, for nitrogen 0·2438; the molecular weights of these
      gases are 32, 2, and 28, and therefore for oxygen K = 0·2175-0·0625
      = 0·1550, for hydrogen K = 3·4050-1·000 = 2·4050, and for nitrogen
      K = 0·2438-0·0714 = 0·1724. These true specific heats of elements
      are in inverse proportion to their atomic weights--that is, their
      product by the atomic weight is a constant quantity. In fact, for
      oxygen this product = 0·155 × 16 = 2·48, for hydrogen 2·40, for
      nitrogen 0·7724 × 14 = 2·414, and therefore if A stand for the
      atomic weight we obtain the expression K × A = a constant, which
      may be taken as 2·45. This is the true expression of Dulong and
      Petit's law, because K is the true specific heat and A the weight
      of the atom. It should be remarked, moreover, that the product of
      the observed specific heat Q into A is also a constant quantity
      (for oxygen = 3·48, for hydrogen = 3·40), because the external work
      D is also inversely proportional to the atomic weight.

      In the case of gases we distinguish the specific heat at a constant
      pressure _c´_ (we designated this quantity above by Q), and at a
      constant volume _c_. It is evident that _the relation between the
      two specific heats, k_, judging from the above, is the ratio of Q
      to K, or equal to the ratio of 2·45_n_ + 2 to 2·45_n_. When _n_ = 1
      this ratio _k_ = 1·8; when _n_ = 2, _k_ = 1·4, when _n_ = 3, _k_ =
      1·3, and with an exceedingly large number _n_, of atoms in the
      molecule, _k_ = 1. That is, the ratio between the specific heats
      decreases from 1·8 to 1·0 as the number of atoms, _n_, contained in
      the molecule increases. This deduction is verified to a certain
      extent by direct experiment. For such gases as hydrogen, oxygen,
      nitrogen, carbonic oxide, air, and others in which _n_ = 2, the
      magnitude of _k_ is determined by methods described in works on
      physics (for example, by the change of temperature with an
      alteration of pressure, by the velocity of sound, &c.) and is found
      in reality to be nearly 1·4, and for such gases as carbonic
      anhydride, nitric dioxide, and others it is nearly 1·3. Kundt and
      Warburg (1875), by means of the approximate method mentioned in
      Note 29, Chapter VII., determined _k_ for mercury vapour when _n_ =
      1, and found it to be = 1·67--that is, a larger quantity than for
      air, as would be expected from the above.

      It may be admitted that the true atomic heat of gases = 2·43, only
      under the condition that they are distant from a liquid state, and
      do not undergo a chemical change when heated--that is, when no
      internal work is produced in them (B = 0). Therefore this work may
      to a certain extent be judged by the observed specific heat. Thus,
      for instance, for chlorine (Q = 0·12, Regnault; _k_ = 1·33,
      according to Straker and Martin, and therefore K = 0·09, MK = 6·4),
      the atomic heat (3·2) is much greater than for other gases
      containing two atoms in a molecule, and it must be assumed,
      therefore, that when it is heated some great internal work is
      accomplished.

      In order to generalise the facts concerning the specific heat of
      gases and solids, it appears to me possible to accept the following
      general proposition: _the atomic heat_ (that is, AQ or QM/_n_,
      where M is the molecular weight and _n_ the number of molecules) is
      _smaller_ (in solids it attains its highest value 6·8 and in gases
      3·4), _the more complex the molecule_ (i.e. _the greater the number
      (n) of atoms forming it_) _and so much smaller, up to a certain
      point_ (in similar physical states) _the smaller the mean atomic
      weight M/n_.

  [8] As an example, it will be sufficient to refer to the specific heat
      of nitrogen tetroxide, N_{2}O_{4}, which, when heated, gradually
      passes into NO_{2}--that is, chemical work of decomposition
      proceeds, which consumes heat. Speaking generally, specific heat is
      a complex quantity, in which it is clear that thermal data (for
      instance, the heat of reaction) alone cannot give an idea either of
      chemical or of physical changes individually, but always depend on
      an association of the one and the other. If a substance be heated
      from _t__{0} to _t__{1} it cannot but suffer a chemical change
      (that is, the state of the atoms in the molecules changes more or
      less in one way or another) if dissociation sets in at a
      temperature _t__{1}. Even in the case of the elements whose
      molecules contain only one atom, a true chemical change is possible
      with a rise of temperature, because more heat is evolved in
      chemical reactions than that quantity which participates in purely
      physical changes. One gram of hydrogen (specific heat = 3·4 at a
      constant pressure) cooled to the temperature of absolute zero will
      evolve altogether about one thousand units of heat, 8 grams of
      oxygen half this amount, whilst in combining together they evolve
      in the formation of 9 grams of water more than thirty times as much
      heat. Hence the store of chemical energy (that is, of the motion of
      the atoms, vortex, or other) is much greater than the physical
      store proper to the molecules, but it is the change accomplished by
      the former that is the cause of chemical transformations. Here we
      evidently touch on those limits of existing knowledge beyond which
      the teaching of science does not yet allow us to pass. Many new
      scientific discoveries have still to be made before this is
      possible.

Among the bivalent metals the first place, with respect to their
distribution in nature, is occupied by _magnesium_ and _calcium_, just as
sodium and potassium stand first amongst the univalent metals. The
relation which exists between the atomic weights of these four metals
confirms the above comparison. In fact, the combining weight of magnesium
is equal to 24, and of calcium 40; whilst the combining weights of sodium
and potassium are 23 and 39--that is, the latter are one unit less than
the former.[9] They all belong to the number of _light metals_, as they
have but a small specific gravity, in which respect they differ from the
ordinary, generally known heavy, or ore, metals (for instance, iron,
copper, silver, and lead), which are distinguished by a much greater
specific gravity. There is no doubt that their low specific gravity has a
significance, not only as a simple point of distinction, but also as a
property which determines the fundamental properties of these metals.
Indeed, all the light metals have a series of points of resemblance with
the metals of the alkalis; thus both magnesium and calcium, like the
metals of the alkalis, decompose water (without the addition of acids),
although not so easily as the latter metals. The process of the
decomposition is essentially one and the same; for example, Ca + 2H_{2}O
= CaH_{2}O_{2} + H_{2}--that is, hydrogen is liberated and a hydroxide of
the metal formed. These hydroxides are bases which neutralise nearly all
acids. However, the hydroxides RH_{2}O_{2} of calcium and magnesium are
in no respect so energetic as the hydroxides of the true metals of the
alkalis; thus when heated they lose water, are not so soluble, develop
less heat with acids, and form various salts, which are less stable and
more easily decomposed by heat than the corresponding salts of sodium and
potassium. Thus calcium and magnesium carbonates easily part with
carbonic anhydride when ignited; the nitrates are also very easily
decomposed by heat, calcium and magnesium oxides, CaO and MgO, being left
behind. The chlorides of magnesium and calcium, when heated with water,
evolve hydrogen chloride, forming the corresponding hydroxides, and when
ignited the oxides themselves. All these points are evidence of a
weakening of the alkaline properties.

  [9] As if NaH = Mg and KH = Ca, which is in accordance with their
      valency. KH includes two monovalent elements, and is a bivalent
      group like Ca.

These metals have been termed _the metals of the alkaline earths_,
because they, like the alkali metals, form energetic bases. They are
called alkaline _earths_ because they are met with in nature in a state
of combination, forming the insoluble mass of the earth, and because as
oxides, RO, they themselves have an earthy appearance. Not a few salts of
these metals are known which are insoluble in water, whilst the
corresponding salts of the alkali metals are generally soluble--for
example, the carbonates, phosphates, borates, and other salts of the
alkaline earth metals are nearly insoluble. This enables us to separate
the metals of the alkaline earths from the metals of the alkalis. For
this purpose a solution of ammonium carbonate is added to a mixed
solution of salts of both kinds of metals, when by a double decomposition
the insoluble carbonates of the metals of the alkaline earths are formed
and fall as a precipitate, whilst the metals of the alkalis remain in
solution: RX_{2} + Na_{2}CO_{3} = RCO_{3} + 2NaX.

We may here remark that the oxides of the metals of the alkaline earths
are frequently called by special names: MgO is called magnesia or bitter
earth; CaO, lime; SrO, strontia; and BaO, baryta.

In the primary rocks the oxides of calcium and magnesium are combined
with silica, sometimes in variable quantities, so that in some cases
the lime predominates and in other cases the magnesium. The two oxides,
being analogous to each other, replace each other in equivalent
quantities. The various forms of _augite_, _hornblende_, or _amphibole_,
and of similar minerals, which enter into the composition of nearly
all rocks, contain lime and magnesia and silica. The majority of the
primary rocks also contain alumina, potash, and soda. These rocks,
under the action of water (containing carbonic acid) and air, give
up lime and magnesia to the water, and therefore they are contained
in all kinds of water, and especially in sea-water. The _carbonates_
CaCO_{3} and MgCO_{3}, frequently met with in nature, _are soluble
in an excess of water saturated with carbonic anhydride_,[10] and
therefore many natural waters contain these salts, and are able to
yield them when evaporated. However, one kilogram of water saturated
with carbonic anhydride does not dissolve more than three grams of
calcium carbonate. By gradually expelling the carbonic anhydride from
such water, an insoluble precipitate of calcium carbonate separates
out. It may confidently be stated that the formation of the very widely
distributed strata of calcium and magnesium carbonates was of this
nature, because these strata are of a sedimentary character--that is,
such as would be exhibited by a gradually accumulating deposit on the
bottom of the sea, and, moreover, frequently containing the remains of
marine plants, and animals, shells, &c. It is very probable that the
presence of these organisms in the sea has played the chief part in the
precipitation of the carbonates from the sea water, because the plants
absorb CO_{2}, and many of the organisms CaCO_{3}, and after death give
deposits of carbonate of lime; for instance, chalk, which is almost
entirely composed of the minute remains of the calcareous shields of
such organisms. These deposits of calcium and magnesium carbonates are
the most important sources of these metals. Lime generally predominates,
because it is present in rocks and running water in greater quantity
than magnesia, and in this case these sedimentary rocks are termed
_limestone_. Some common flagstones used for paving, &c., and chalk may
be taken as examples of this kind of formation. Those limestones in
which a considerable portion of the calcium is replaced by magnesium are
termed _dolomites_. The dolomites are distinguished by their hardness,
and by their not parting with the whole of their carbonic anhydride
so easily as the limestones under the action of acids. Dolomites[11]
sometimes contain an equal number of molecules of calcium carbonate and
magnesium carbonate, and they also sometimes appear in a crystalline
form, which is easily intelligible, because calcium carbonate itself
is exceedingly common in this form in nature, and is then known as
_calc spar_, whilst natural crystalline magnesium carbonate is termed
_magnesite_. The formation of the crystalline varieties of the insoluble
carbonates is explained by the possibility of a slow deposition from
solutions containing carbonic acid. Besides which (Chapter X.) calcium
and magnesium sulphates are obtained from sea water, and therefore they
are met with both as deposits and in springs. It must be observed that
magnesium is held in considerable quantities in sea water, because the
sulphate and chloride of magnesium are very soluble in water, whilst
calcium sulphate is but little soluble, and is used in the formation
of shells; and therefore if the occurrence of considerable deposits of
magnesium sulphate cannot be expected in nature, still, on the other
hand, one would expect (and they do actually occur) large masses of
calcium sulphate or _gypsum_, CaSO_{4},2H_{2}O. Gypsum sometimes forms
strata of immense size, which extend over many hectometres--for example,
in Russia on the Volga, and in the Donetz and Baltic provinces.

  [10] Sodium carbonate and other carbonates of the alkalis give acid
       salts which are less soluble than the normal; here, on the
       contrary, with an excess of carbonic anhydride, a salt is formed
       which is more soluble than the normal, but this acid salt is more
       unstable than sodium hydrogen carbonate, NaHCO_{3}.

  [11] The formation of dolomite may be explained, if only we imagine
       that a solution of a magnesium salt acts on calcium carbonate.
       Magnesium carbonate may be formed by double decomposition, and it
       must be supposed that this process ceases at a certain limit
       (Chapter XII.), when we shall obtain a mixture of the carbonates
       of calcium and magnesium. Haitinger heated a mixture of calcium
       carbonate, CaCO_{3}, with a solution of an equivalent quantity of
       magnesium sulphate, MgSO_{4}, in a closed tube at 200°, and then a
       portion of the magnesia actually passed into the state of
       magnesium carbonate, MgCO_{3}, and a portion of the lime was
       converted into gypsum, CaSO_{4}. Lubavin (1892) showed that
       MgCO_{3} is more soluble than CaCO_{3} in salt water, which is of
       some significance in explaining the composition of sea water.

Lime and magnesia also, but in much smaller quantities (only to the
amount of several fractions of a per cent. and rarely more), enter into
the composition of every fertile soil, and without these bases the soil
is unable to support vegetation. Lime is particularly important in this
respect, and its presence in a larger quantity generally improves the
harvest, although purely calcareous soils are as a rule infertile. For
this reason the soil is fertilised both with lime[12] itself and with
marl--that is, with clay mixed with a certain quantity of calcium
carbonate, strata of which are found nearly everywhere.

  [12] The undoubted action of lime in increasing the fertility of
       soils--if not in every case, at all events, with ordinary soils
       which have long been under corn--is based not so much on the need
       of plants for the lime itself as on those chemical and physical
       changes which it produces in the soil, as a particularly powerful
       base which aids the alteration of the mineral and organic elements
       of the soil.

From the soil the lime and magnesia (in a smaller quantity) pass into the
substance of _plants_, where they occur as salts. Certain of these salts
separate in the interior of plants in a crystalline form--for example,
calcium oxalate. The lime occurring in plants serves as the source for
the formation of the various calcareous secretions which are so common in
_animals_ of all classes. The bones of the highest animal orders, the
shells of mollusca, the covering of the sea-urchin, and similar solid
secretions of sea animals, contain calcium salts; namely, the shells
mainly calcium carbonate, and the bones mainly calcium phosphate. Certain
limestones are almost entirely formed of such deposits. Odessa is
situated on a limestone of this kind, composed of shells. Thus magnesium
and calcium occur throughout the entire realm of nature, but calcium
predominates.

As lime and magnesia form bases which are in many respects analogous,
they were not distinguished from each other for a long time. Magnesia
was obtained for the first time in the seventeenth century from Italy,
and used as a medicine; and it was only in the last century that Black,
Bergmann, and others distinguished magnesia from lime.

_Metallic magnesium_ (and calcium also) is not obtained by heating
magnesium oxide or the carbonate with charcoal, as the alkali metals
are obtained,[13] but is liberated by the action of a galvanic current
on fused magnesium chloride (best mixed with potassium chloride); Davy
and Bussy obtained metallic magnesium by acting on magnesium chloride
with the vapours of potassium. At the present time (Deville's process)
magnesium is prepared in rather considerable quantities by a similar
process, only the potassium is replaced by sodium. Anhydrous magnesium
chloride, together with sodium chloride and calcium fluoride, is fused
in a close crucible. The latter substances only serve to facilitate the
formation of a fusible mass before and after the reaction, which is
indispensable in order to prevent the access and action of air. One part
of finely divided sodium to five parts of magnesium chloride is thrown
into the strongly heated molten mass, and after stirring the reaction
proceeds very quickly, and magnesium separates, MgCl_{2} + Na_{2} = Mg
+ 2NaCl. In working on a large scale, the powdery metallic magnesium
is then subjected to distillation at a white heat. The distillation
of the magnesium is necessary, because the undistilled metal is not
homogeneous[14] and burns unevenly: the metal is prepared for the
purpose of illumination. Magnesium is a white metal, like silver; it
is not soft like the alkali metals, but is, on the contrary, hard like
the majority of the ordinary metals. This follows from the fact that it
melts at a somewhat high temperature--namely, about 500°--and boils at
about 1000°. It is malleable and ductile, like the generality of metals,
so that it can be drawn into wires and rolled into ribbon; it is most
frequently used for lighting purposes in the latter form. Unlike the
alkali metals, magnesium does not decompose the atmospheric moisture at
the ordinary temperature, so that it is almost unacted on by air; it is
not even acted on by water at the ordinary temperature, so that it may
be washed to free it from sodium chloride. Magnesium only decomposes
water with the evolution of hydrogen at the boiling point of water,[15]
and more rapidly at still higher temperatures. This is explained by the
fact that in decomposing water magnesium forms an insoluble hydroxide,
MgH_{2}O_{2}, which covers the metal and hinders the further action
of the water. Magnesium easily displaces hydrogen from acids, forming
magnesium salts. When ignited it _burns_, not only in oxygen but in air
(and even in carbonic anhydride), forming a white powder of magnesium
oxide, or magnesia; in burning it emits a white and exceedingly
_brilliant light_. The strength of this light naturally depends on the
fact that magnesium (24 parts by weight) in burning evolves about
140 thousand heat units, and that the product of combustion, MgO,
is infusible by heat; so that the vapour of the burning magnesium
contains an ignited powder of non-volatile and infusible magnesia,
and consequently presents all the conditions for the production of
a brilliant light. The light emitted by burning magnesium contains
many rays which act chemically, and are situated in the violet and
ultra-violet parts of the spectrum. For this reason burning magnesium
may be employed for producing photographic images.[16]

  [13] Sodium and potassium only decompose magnesium oxide at a white
       heat and very feebly, probably for two reasons. In the first
       place, because the reaction Mg + O develops more heat (about 140
       thousand calories) than K_{2} + O or Na_{2} + O (about 100
       thousand calories); and, in the second place, because magnesia is
       not fusible at the heat of a furnace and cannot act on the
       charcoal, sodium, or potassium--that is, it does not pass into
       that mobile state which is necessary for reaction. The first
       reason alone is not sufficient to explain the absence of the
       reaction between charcoal and magnesia, because iron and charcoal
       in combining with oxygen evolve less heat than sodium or
       potassium, yet, nevertheless, they can displace them. With respect
       to magnesium chloride, it acts on sodium and potassium, not only
       because their combination with chlorine evolves more heat than the
       combination of chlorine and magnesium (Mg + Cl_{2} gives 150 and
       Na_{2} + Cl_{2} about 195 thousand calories), but also because a
       fusion, both of the magnesium chloride and of the double salt,
       takes place under the action of heat. It is probable, however,
       that a reverse reaction will take place. A reverse reaction might
       probably be expected, and Winkler (1890) showed that Mg reduces
       the oxides of the alkali metals (Chapter XIII., Note 42).

  [14] Commercial magnesium generally contains a certain amount of
       magnesium nitride (Deville and Caron), Mg_{3}N_{2}--that is, a
       product of substitution of ammonia which is directly formed (as is
       easily shown by experiment) when magnesium is heated in nitrogen.
       It is a yellowish green powder, which gives ammonia and magnesia
       with water, and cyanogen when heated with carbonic anhydride.
       Pashkoffsky (1893) showed that Mg_{3}N_{2} is easily formed and is
       the sole product when Mg is heated to redness in a current of
       NH_{3}. Perfectly pure magnesium may be obtained by the action of
       a galvanic current.

  [15] Hydrogen peroxide (Weltzien) dissolves magnesium. The reaction has
       not been investigated.

  [16] A special form of apparatus is used for burning magnesium. It is a
       clockwork arrangement in which a cylinder rotates, round which a
       ribbon or wire of magnesium is wound. The wire is subjected to a
       uniform unwinding and burning as the cylinder rotates, and in this
       manner the combustion may continue uniform for a certain time. The
       same is attained in special lamps, by causing a mixture of sand
       and finely divided magnesium to fall from a funnel-shaped
       reservoir on to the flame. In photography it is best to blow
       finely divided magnesium into a colourless (spirit or gas) flame,
       and for instantaneous photography to light a cartridge of a
       mixture of magnesium and chlorate of potassium by means of a spark
       from a Ruhmkorff's coil (D. Mendeléeff, 1889).

Owing to its great affinity for oxygen, magnesium _reduces_ many metals
(zinc, iron, bismuth, antimony, cadmium, tin, lead, copper, silver, and
others) from solutions of their salts at the ordinary temperature,[17]
and at a red heat finely divided magnesium takes up the oxygen from
silica, alumina, boric anhydride, &c.; so that silicon and similar
elements may be obtained by directly heating a mixture of powdered silica
and magnesium in an infusible glass tube.[18]

  [17] According to the observations of Maack, Comaille, Böttger, and
       others. The reduction by heat mentioned further on was pointed out
       by Geuther, Phipson, Parkinson and Gattermann.

  [18] This action of metallic magnesium in all probability depends,
       although only partially (_see_ Note 13), on its volatility, and on
       the fact that, in combining with a given quantity of oxygen, it
       evolves more heat than aluminium, silicon, potassium, and other
       elements.

The affinity of magnesium for the halogens is much more feeble than for
oxygen,[19] as is at once evident from the fact that a solution of iodine
acts feebly on magnesium; still magnesium burns in the vapours of iodine,
bromine, and chlorine. The character of magnesium is also seen in the
fact that all its salts, especially in the presence of water, are
decomposable at a comparatively moderate temperature, the elements of the
acid being evolved, and the magnesium oxide, which is non-volatile and
unchangeable by heat, being left. This naturally refers to those acids
which are themselves volatilised by heat. Even magnesium sulphate is
completely decomposed at the temperature at which iron melts, oxide of
magnesium remaining behind. This decomposition of magnesium salts by heat
proceeds much more easily than that of calcium salts. For example,
magnesium carbonate is totally decomposed at 170°, magnesium oxide being
left behind. This _magnesia_, or _magnesium oxide_, is met with both in
an anhydrous and hydrated state in nature (the anhydrous magnesia as the
mineral _periclase_, MgO, and the hydrated magnesia as _brucite_,
MgH_{2}O_{2}). Magnesia is a well-known medicine (calcined
magnesia--_magnesia usta_). It is a white, extremely fine, and very
voluminous powder, of specific gravity 3·4; it is infusible by heat, and
only shrinks or shrivels in an oxyhydrogen flame. After long contact the
anhydrous magnesia combines with water, although very slowly, forming the
hydroxide Mg(HO)_{2}, which, however, parts with its water with great
ease when heated even below a red heat, and again yields anhydrous
magnesia. This hydroxide is obtained directly as a gelatinous amorphous
substance when a soluble alkali is mixed with a solution of any magnesium
salt, MgCl_{2} + 2KHO = Mg(HO)_{2} + 2KCl. This decomposition is
complete, and nearly all the magnesium passes into the precipitate; and
this clearly shows the almost perfect insolubility of magnesia in water.
Water dissolves a scarcely perceptible quantity of magnesium
hydroxide--namely, one part is dissolved by 55,000 parts of water. Such a
solution, however, has an alkaline reaction, and gives, with a salt of
phosphoric acid, a precipitate of magnesium phosphate, which is still
more insoluble. Magnesia is not only dissolved by acids, forming salts,
but it also displaces certain other bases--for example, ammonia from
ammonium salts when boiled; and the hydroxide also absorbs carbonic
anhydride from the air. The magnesium salts, like those of calcium,
potassium, and sodium, are colourless if they are formed from colourless
acids. Those which are soluble have a bitter taste, whence magnesia has
been termed _bitter-earth_. In comparison with the alkalis magnesia is a
feeble base, inasmuch as it forms somewhat unstable salts, easily gives
basic salts, forms acid salts with difficulty, and is able to give double
salts with the salts of the alkalis, which facts are characteristic of
feeble bases, as we shall see in becoming acquainted with the different
metals.

  [19] Davy, on heating magnesia in chlorine, concluded that there was a
       complete substitution, because the volume of the oxygen was half
       the volume of the chlorine; it is probable, however, that owing to
       the formation of chlorine oxide (Chapter XI., Note 30) the
       decomposition is not complete and is limited by a reverse
       reaction.

The power of magnesium salts to form double and basic salts is very
frequently shown in reactions, and is specially marked as
regards ammonium salts. If saturated solutions of magnesium and
ammonium sulphates are mixed together, a crystalline double salt
Mg(NH_{4})_{2}(SO_{4})_{2},6H_{2}O,[20] is immediately precipitated.
A strong solution of ordinary ammonium carbonate dissolves magnesium
oxide or carbonate, and precipitates crystals of a double salt,
Mg(NH_{4})_{2}(CO_{3})_{2},4H_{2}O, from which water extracts the
ammonium carbonate. With an excess of an ammonium salt the double
salt passes into solution,[21] and therefore if a solution contain
a magnesium salt and an excess of an ammonium salt--for instance,
sal-ammoniac--then sodium carbonate will no longer precipitate
magnesium carbonate. A mixture of solutions of magnesium and ammonium
chlorides, on evaporation or refrigeration, gives a double salt,
Mg(NH_{4})Cl_{3},6H_{2}O.[22] The salts of potassium, like those
of ammonium, are able to enter into combination with the magnesium
salts.[23] For instance, the double salt, MgKCl_{3},6H_{2}O, which is
known as _carnallite_,[24] and occurs in the salt mines of Stassfurt,
may be formed by freezing a saturated solution of potassium chloride
with an excess of magnesium chloride. A saturated solution of magnesium
sulphate dissolves potassium sulphate, and solid magnesium sulphate
is soluble in a saturated solution of potassium sulphate. A double
salt, K_{2}Mg(SO_{4})_{2},6H_{2}_O, which closely resembles the
above-mentioned ammonium salt, crystallises from these solutions.[25]
The nearest analogues of magnesium are able to give exactly similar
double salts, both in crystalline form (monoclinic system) and
composition; they, like this salt (_see_ Chapter XV.), are easily
able (at 140°) to part with all their water of crystallisation,
and correspond with the salts of sulphuric acid, whose type may
be taken as _magnesium sulphate_, MgSO_{4}.[26] It occurs at
Stassfurt as _kieserite_, MgSO_{4},H_{2}O, and generally separates
from solutions as a heptahydrated salt, MgSO_{4},7H_{2}O, and from
supersaturated solutions as a hexahydrated salt, MgSO_{4},6H_{2}O; at
temperatures below 0° it crystallises out as a dodecahydrated salt,
MgSO_{4},12H_{2}O, and a solution of the composition MgSO_{4},2H_{2}O
solidifies completely at -5°.[27] Thus between water and magnesium
sulphate there may exist several definite and more or less stable
degrees of equilibrium; the double salt MgSO_{4}K_{2}SO_{4},6H_{2}O
may be regarded as one of these equilibrated systems, the more so
since it contains 6H_{2}O, whilst MgSO_{4} forms its most stable
system with 7H_{2}O, and the double salt may be considered as this
crystallo-hydrate in which one molecule of water is replaced by the
molecule K_{2}SO_{4}.[28]

  [20] Even a solution of ammonium chloride gives this salt with
       magnesium sulphate. Its sp. gr. is 1·72; 100 parts of water at 0°
       dissolve 9, at 20° 17·9 parts of the anhydrous salt. At about 130°
       it loses all its water.

  [21] This is an example of equilibrium and of the influence of mass;
       the double salt is decomposed by water, but if instead of water we
       take a solution of that soluble part which is formed in the
       decomposition of the double salt, then the latter dissolves as a
       whole.

  [22] If an excess of ammonia be added to a solution of magnesium
       chloride, only half the magnesium is thrown down in the
       precipitate, 2MgCl_{2} + 2NH_{4}.OH = Mg(OH)_{2} + Mg.NH_{4}Cl_{3}
       + NH_{4}Cl. A solution of ammonium chloride reacts with magnesia,
       evolving ammonia and forming a solution of the same salt, MgO +
       3NH_{4}Cl = MgNH_{4}Cl_{3} + H_{2}O + 2NH_{3}.

       Among the double salts of ammonium and magnesium, the phosphate,
       MgNH_{4}PO_{4},6H_{2}O, is almost insoluble in water (0·07 gram is
       soluble in a litre), even in the presence of ammonia. Magnesia is
       very frequently precipitated as this salt from solutions in which
       it is held by ammonium salts. As lime is not retained in solution
       by the presence of ammonium salts, but is precipitated
       nevertheless by sodium carbonate, &c., it is very easy to separate
       calcium from magnesium by taking advantage of these properties.

  [23] In order to see the nature and cause of formation of double salts,
       it is sufficient (although this does not embrace the whole essence
       of the matter) to consider that one of the metals of such salts
       (for instance, potassium) easily gives acid salts, and the other
       (in this instance, magnesium) basic salts; the properties of
       distinctly basic elements predominate in the former, whilst in the
       latter these properties are enfeebled, and the salts formed by
       them bear the character of acids--for example, the salts of
       aluminium or magnesium act in many cases like acids. By their
       mutual combination these two opposite properties of the salts are
       both satisfied.

  [24] Carnallite has been mentioned in Chapter X. (Note 4) and in
       Chapter XIII. These deposits also contain much _kainite_,
       KMgCl(SO_{4}),3H_{2}O (sp. gr. 2·13; 100 parts of water dissolve
       79·6 parts at 18°). This double salt contains two metals and two
       haloids. Feit (1889) also obtained a bromide corresponding to
       carnallite.

  [25] The component parts of certain double salts diffuse at different
       rates, and as the diffused solution contains a different
       proportion of the component salts than the solution taken of the
       double salt, it shows that such salts are decomposed by water.
       According to Rüdorff, the double salts, like carnallite,
       MgK_{2}(SO_{4})_{2},6H_{2}O, and the alums, all belong to this
       order (1888). But such salts as tartar emetic, the double
       oxalates, and double cyanides are not separated by diffusion,
       which in all probability depends both on the relative rate of the
       diffusion of the component salts and on the degree of affinity
       acting between them. Those complex states of equilibrium which
       exist between water, the individual salts MX and NY, and the
       double salt MNXY, have been already partially analysed (as will be
       shown hereafter) in that case when the system is heterogeneous
       (that is, when something separates out in a solid state from the
       liquid solution), but in the case of equilibria in a homogeneous
       liquid medium (in a solution) the phenomenon is not so clear,
       because it concerns that very theory of solution which cannot yet
       be considered as established (Chapter I., Note 9, and others). As
       regards the heterogeneous decomposition of double salts, it has
       long been known that such salts as carnallite and
       K_{2}Mg(SO_{4})_{2} give up the more soluble salt if an
       insufficient quantity of water for their complete solution be
       taken. The complete saturation of 100 parts of water requires at
       0° 14·1, at 20° 25, and at 60° 50·2 parts of the latter double
       salt (anhydrous), while 100 parts of water dissolve 27 parts of
       magnesium sulphate at 0°, 36 parts at 20°, and 55 parts at 60°, of
       the anhydrous salt taken. Of all the states of equilibrium
       exhibited by double salts the most fully investigated as yet is
       the system containing water, sodium sulphate, magnesium sulphate,
       and their double salt, Na_{2}Mg(SO_{4})_{2}, which crystallises
       with 4 and 6 mol. OH_{2}. The first crystallo-hydrate,
       MgNa_{2}(SO_{4})_{2},4H_{2}O, occurs at Stassfurt, and as a
       sedimentary deposit in many of the salt lakes near Astrakhan, and
       is therefore called _astrakhanite_. The specific gravity of the
       monoclinic prisms of this salt is 2·22. If this salt, in a finely
       divided state, be mixed with the necessary quantity of water
       (according to the equation MgNa_{2}(SO_{4})_{2},4H_{2}O + 13H_{2}O
       = Na_{2}SO_{4},10H_{2}O + MgSO_{4},7H_{2}O), the mixture
       solidifies like plaster of Paris into a homogeneous mass if the
       temperature be _below_ 22° (Van't Hoff und Van Deventer, 1886;
       Bakhuis Roozeboom, 1887); but if the temperature be above this
       _transition-point_ the water and double salt do not react on each
       other: that is, they do not solidify or give a mixture of sodium
       and magnesium sulphates. If a mixture (in equivalent quantities)
       of solutions of these salts be evaporated, and crystals of
       astrakhanite and of the individual salts capable of proceeding
       from it be added to the concentrated solution to avoid the
       possibility of a supersaturated solution, then at temperatures
       above 22° astrakhanite is exclusively formed (this is the method
       of its production), but at lower temperatures the individual salts
       are alone produced. If equivalent amounts of Glauber's salt and
       magnesium sulphate be mixed together in a solid state, there is no
       change at temperatures below 22°, but at higher temperatures
       astrakhanite and water are formed. The volume occupied by
       Na_{2}SO_{4},10H_{2}O in grams = 322/1·46 = 220·5 cubic
       centimetres, and by MgSO_{4},7H_{2}O = 246/1·68 = 146·4 c.c.;
       hence their mixture in equivalent quantities occupies a volume of
       366·9 c.c. The volume of astrakhanite = 334/2·22 = 150·5 c.c., and
       the volume of 13H_{2}O = 234 c.c., hence their sum = 380·5 c.c.,
       and therefore it is easy to follow the formation of the
       astrakhanite in a suitable apparatus (a kind of thermometer
       containing oil and a powdered mixture of sodium and magnesium
       sulphates), and to see by the variation in volume that below 22°
       it remains unchanged, and at higher temperatures proceeds the more
       quickly the higher the temperature. At the transition temperature
       the solubility of astrakhanite and of the mixture of the component
       salts is one and the same, whilst at higher temperatures a
       solution which is saturated for a mixture of the individual salts
       would be supersaturated for astrakhanite, and at lower
       temperatures the solution of astrakhanite will be supersaturated
       for the component salts, as has been shown with especial detail by
       Karsten, Deacon, and others. Roozeboom showed that there are two
       limits to the composition of the solutions which can exist for a
       double salt; these limits are respectively obtained by dissolving
       a mixture of the double salt with each of its component simple
       salts. Van't Hoff demonstrated, besides this, that the tendency
       towards the formation of double salts has a distinct influence on
       the progress of double decomposition, for at temperatures above
       31° the mixture 2MgSO_{4},7H_{2}O + 2NaCl passes into
       MgNa_{2}(SO_{4})_{2},4H_{2}O + MgCl_{2},6H_{2}O + 4H_{2}O, whilst
       below 31° there is not this double decomposition, but it proceeds
       in the opposite direction, as may be demonstrated by the
       above-described methods. Van der Heyd obtained a potassium
       astrakhanite, K_{2}SO_{4}MgSO_{4},4H_{2}O, from solutions of the
       component salts at 100°.

       From these experiments on double salts we see that there is as
       close a dependence between the temperature and the formation of
       substances as there is between the temperature and a change of
       state. It is a case of Deville's principles of dissociation,
       extended in the direction of the passage of a solid into a liquid.
       On the other hand, we see here how essential a _rôle_ water plays
       in the formation of compounds, and how the affinity for water of
       crystallisation is essentially analogous to the affinity between
       salts, and hence also to the affinity of acids for bases, because
       the formation of double salts does not differ in any essential
       point (except the degree of affinity--that is, from a quantitative
       aspect) from the formation of salts themselves. When sodium
       hydroxide with nitric acid gives sodium nitrate and water the
       phenomenon is essentially the same as in the formation of
       astrakhanite from the salts Na_{2}SO_{4},10H_{2}O and
       MgSO_{4},7H_{2}O. Water is disengaged in both cases, and hence the
       volumes are altered.

  [26] This salt, and especially its crystallo-hydrate with 7H_{2}O, is
       generally known as Epsom salts. It has long been used as a
       purgative. It is easily obtained from magnesia and sulphuric acid,
       and it separates on the evaporation of sea water and of many
       saline springs. When carbonic anhydride is obtained by the action
       of sulphuric acid on magnesite, magnesium sulphate remains in
       solution. When dolomite--that is, a mixture of magnesium and
       calcium carbonates--is subjected to the action of a solution of
       hydrochloric acid until about half of the salt remains, the
       calcium carbonate is mostly dissolved and magnesium carbonate is
       left, which by treatment with sulphuric acid gives a solution of
       magnesium sulphate.

  [27] The anhydrous salt, MgSO_{4} (sp. gr. 2·61), attracts moisture
       (7 mol. H_{2}O) from moist air; when heated in steam or hydrogen
       chloride it gives sulphuric acid, and when heated with carbon it
       is decomposed according to the equation 2MgSO_{4} + C = 2SO_{2} +
       CO_{2} + 2MgO. The monohydrated salt (kieserite), MgSO_{4},H_{2}O
       (sp. gr. 2·56), dissolves in water with difficulty; it is formed
       by heating the other crystallo-hydrates to 135°. The hexahydrated
       salt is dimorphous. If a solution, saturated at the boiling-point,
       be prepared, and cooled without access of crystals of the
       heptahydrated salt, then MgSO_{4},6H_{2}O crystallises out in
       _monoclinic_ prisms (Loewel, Marignac), which are quite as
       unstable as the salt, Na_{2}SO_{4},7H_{2}O; but if prismatic
       crystals of the cubic system of the copper-nickel salts of the
       composition MSO_{4},6H_{2}O be added, then crystals of
       MgSO_{4},6H_{2}O are deposited on them as prisms of the _cubic_
       system (Lecoq de Boisbaudran). The common crystallo-hydrate,
       MgSO_{4},7H_{2}O, Epsom salts, belongs to the _rhombic_ system,
       and is obtained by crystallisation below 30°. Its specific gravity
       is 1·69. In a vacuum, or at 100°, it loses 5H_{2}O, at 132°
       6H_{2}O, and at 210° all the 7H_{2}O (Graham). If crystals of
       ferrous or cobaltic sulphate be placed in a saturated solution,
       _hexagonal_ crystals of the heptahydrated salt are formed (Lecoq
       de Boisbaudran); they present an unstable state of equilibrium,
       and soon become cloudy, probably owing to their transformation
       into the more stable common form. Fritzsche, by cooling saturated
       solutions below 0°, obtained a mixture of crystals of ice and of a
       dodecahydrated salt, which easily split up at temperatures above
       0°. Guthrie showed that dilute solutions of magnesium sulphate,
       when refrigerated, separate ice until the solution attains a
       composition MgSO_{4},24H_{2}O, which will completely freeze into a
       crystallo-hydrate at -5·3°. According to Coppet and Rüdorff, the
       temperature of the formation of ice falls by 0·073° for every part
       by weight of the heptahydrated salt per 100 of water. This figure
       gives (Chapter I., Note 49) _i_ = 1 for both the heptahydrated and
       the anhydrous salt, from which it is evident that it is impossible
       to judge the state of combination in which a dissolved substance
       occurs by the temperature of the formation of ice.

       The solubility of the different crystallo-hydrates of magnesium
       sulphate, according to Loewel, also varies, like those of sodium
       sulphate or carbonate (_see_ Chapter XII., Notes 7 and 18). At 0°
       100 parts of water dissolves 40·75 MgSO_{4} in the presence of the
       hexahydrated salt, 34·67 MgSO_{4} in the presence of the hexagonal
       heptahydrated salt, and only 26 parts of MgSO_{4} in the presence
       of the ordinary heptahydrated salt--that is, solutions giving the
       remaining crystallo-hydrates will be supersaturated for the
       ordinary heptahydrated salt.

       All this shows how many diverse aspects of more or less stable
       equilibria may exist between water and a substance dissolved in
       it; this has already been enlarged on in Chapter I.

       Carefully purified magnesium sulphate in its aqueous solution
       gives, according to Stcherbakoff, an alkaline reaction with
       litmus, and an acid reaction with phenolphthalein.

       The specific gravity of solutions of certain salts of magnesium
       and calcium reduced to 15°/4° (see my work cited, Chapter I., Note
       19), are, if water at 4° = 10,000,

              MgSO_{4}: _s_ = 9,992 + 99·89_p_ + 0·553_p_^2
              MgCl_{2}: _s_ = 9,992 + 81·31_p_ + 0·372_p_^2
              CaCl_{2}: _s_ = 9,992 + 80·24_p_ + 0·476_p_^2

  [28] Graham even distinguished the last equivalent of the water of
       crystallisation of the heptahydrated salt as that which is
       replaced by other salts, pointing out that double salts like
       MgK_{2}(SO_{4})_{2},6H_{2}O lose all their water at 135°, whilst
       MgSO_{4},7H_{2}O only parts with 6H_{2}O.

_The power of forming basic salts_ is a very remarkable peculiarity of
magnesia and other feeble bases, and especially of those corresponding
with polyvalent metals. The very powerful bases corresponding with
univalent metals--like potassium and sodium--do not form basic salts,
and, indeed, are more prone to give acid salts, whilst magnesium easily
and frequently forms basic salts, especially with feeble acids, although
there are some oxides--as, for example, copper and lead oxides--which
still more frequently give basic salts. If a cold solution of magnesium
sulphate be mixed with a solution of sodium carbonate there is formed a
gelatinous precipitate of a basic salt, Mg(HO)_{2},4MgCO_{3},9H_{2}O;
but all the magnesia is not precipitated in this case, as a portion of it
remains in solution as an acid double salt. If sodium carbonate be added
to a boiling solution of magnesium sulphate a precipitate of a still more
basic salt is formed, 4MgSO_{4} + 4Na_{2}CO_{3} + 4H_{2}O = 4Na_{2}SO_{4}
+ CO_{2} + Mg(OH)_{2},3MgCO_{3},3H_{2}O. This basic salt forms the
ordinary drug _magnesia_ (_magnesia alba_), in the form of light porous
lumps. Other basic salts are formed under certain modifications of
temperature and conditions of decomposition. But _the normal salt_,
MgCO_{3}, which occurs in nature as magnesite in the form of rhombohedra
of specific gravity 3·056, cannot be obtained by such a method of
precipitation. In fact, the formation of the different basic salts shows
the power of water to decompose the normal salt. It is possible, however,
to obtain this salt both in an anhydrous and hydrated state. A solution
of magnesium carbonate in water containing carbonic acid is taken for
this purpose. The reason for this is easily understood--carbonic
anhydride is one of the products of the decomposition of magnesium
carbonate in the presence of water. If this solution be left to evaporate
spontaneously the normal salt separates in a hydrated form, but in the
evaporation of a heated solution, through which a stream of carbonic
anhydride is passed, the anhydrous salt is formed as a crystalline mass,
which remains unaltered in the air, like the natural mineral.[29] The
decomposing influence of water on the salts of magnesium, which is
directly dependent on the feeble basic properties of magnesia,[30] is
most clearly seen in _magnesium chloride_, MgCl_{2}. This salt is
contained[31] in the last mother-liquors of the evaporation of sea-water.
On cooling a sufficiently concentrated solution, the crystallo-hydrate,
MgCl_{2},6H_{2}O, separates;[32] but if it be further heated (above
106°) to remove the water, then hydrochloric acid passes off together
with the latter, so that there ultimately remains magnesia with a small
quantity of magnesium chloride.[33] From what has been said it is evident
that anhydrous magnesium chloride cannot be obtained by simple
evaporation. But if sal-ammoniac or sodium chloride be added to a
solution of magnesium chloride, then the evolution of hydrochloric acid
does not take place, and after complete evaporation the residue is
perfectly soluble in water. This renders it possible to obtain anhydrous
magnesium chloride from its aqueous solution. Indeed the mixture with
sal-ammoniac (in excess) may be dried (the residue consists of an
anhydrous double salt, MgCl_{2},2NH_{4}Cl) and then ignited (460°), when
the sal-ammoniac is converted into vapour and a fused mass of anhydrous
magnesium chloride remains behind. The anhydrous chloride evolves a very
considerable amount of heat on the addition of water, which shows the
great affinity the salt has for water.[34] Anhydrous magnesium chloride
is not only obtained by the above method, but is also formed by the
direct combination of chlorine and magnesium, and by the action of
chlorine on magnesium oxide, oxygen being evolved; this proceeds still
more easily _by heating magnesia with charcoal in a stream of chlorine_,
when the charcoal serves to take up the oxygen. This latter method is
also employed for the preparation of chlorides which are formed in an
anhydrous condition with still greater difficulty than magnesium
chloride. Anhydrous magnesium chloride forms a colourless, transparent
mass, composed of flexible crystalline plates of a pearly lustre. It
fuses at a low red heat (708°) into a colourless liquid, remains
unchanged in a dry state, but under the action of moisture is partially
decomposed even at the ordinary temperature, with formation of
hydrochloric acid. When heated in the presence of oxygen (air) it gives
chlorine and the basic salt, which is formed with even greater facility
under the action of heat in the presence of steam, when HCl is formed,
according to the equation 2MgCl_{2} + H_{2}O = MgOMgCl_{2} + 2HCl.[34 bis]

  [29] The crystalline form of the anhydrous salt obtained in
       this manner is not the same as that of the natural salt. The
       former gives rhombohedra, like those in which calcium carbonate
       appears as calc spar, whilst the natural salt appears as rhombic
       prisms, like those sometimes presented by the same carbonate as
       aragonite, which will shortly be described.

  [30] Magnesium sulphate enters into certain reactions which are proper
       to sulphuric acid itself. Thus, for instance, if a carefully
       prepared mixture of equivalent quantities of hydrated magnesium
       sulphate and sodium chloride be heated to redness, the evolution
       of hydrochloric acid is observed just as in the action of
       sulphuric acid on common salt, MgSO_{4} + 2NaCl + H_{2}O =
       Na_{2}SO_{4} + MgO + 2HCl. Magnesium sulphate acts in a similar
       manner on nitrates, with the evolution of nitric acid. A mixture
       of it with common salt and manganese peroxide gives chlorine.
       Sulphuric acid is sometimes replaced by magnesium sulphate in
       galvanic batteries--for example, in the well-known Meidinger
       battery. In the above-mentioned reactions we see a striking
       example of the similarity of the reactions of acids and salts,
       especially of salts which contain such feeble bases as magnesia.

  [31] As sea-water contains many salts, MCl and MgX_{2}, it follows,
       according to Berthollet's teaching, that MgCl_{2} is also present.

  [32] As the crystallo-hydrates of the salts of sodium often contain
       10H_{2}O, so many of the salts of magnesium contain 6H_{2}O.

  [33] This decomposition is most simply defined as the result of the
       two reverse reactions, MgCl_{2} ÷ H_{2}O = MgO + 2HCl and MgO +
       2HCl = MgCl_{2} + H_{2}O, or as a distribution between O and
       Cl_{2} on the one hand and H_{2} and Mg on the other. (With O,
       MgCl_{2} gives chlorine, _see_ Chapter X., Note 33, and Chapter
       II., Note 3 bis and others, where the reactions and
       applications of MgCl_{2} are given.) It is then clear that,
       according to Berthollet's doctrine, the mass of the hydrochloric
       acid converts the magnesium oxide into chloride, and the mass of
       the water converts the magnesium chloride into oxide. The
       crystallo-hydrate, MgCl_{2},6H_{2}O, forms the limit of the
       reversibility. But an intermediate state of equilibrium may exist
       in the form of basic salts. On mixing ignited magnesia with a
       solution of magnesium chloride of specific gravity about 1·2, a
       solid mass is obtained which is scarcely decomposed by water at
       the ordinary temperature (_see_ Chapter XVI., Note 4). A similar
       means is employed for cementing sawdust into a solid mass, called
       cylolite, used for flooring, &c.

       We may remark that MgBr_{2} crystallises not only with 6H_{2}O
       (temperature of fusion 152°), but also with 10H_{2}O (temperature
       of fusion +12°, formed at -18°). (Panfiloff, 1894).

  [34] According to Thomsen, the combination of MgCl_{2} with 6H_{2}O
       evolves 33,000 calories, and its solution in an excess of water
       36,000.

  [34 bis] Hence MgCl_{2} may be employed for the preparation of chlorine
       and hydrochloric acid (Chapters X. and XI.). In general magnesium
       chloride, which is obtained in large quantities from sea water and
       Stassfurt carnallite, may find numerous practical uses.

_Calcium_ (or the metal of lime) and its compounds in many respects
present a great resemblance to magnesium compounds, but are also clearly
distinguished from them by many properties.[35] In general, calcium
stands to magnesium in the same relation as potassium occupies in respect
to sodium. Davy obtained metallic calcium, like potassium, as an amalgam
by the action of a galvanic current; but neither charcoal nor iron
decomposes calcium oxide, and even sodium decomposes calcium chloride[36]
with difficulty. But a galvanic current easily decomposes calcium
chloride, and metallic sodium somewhat easily decomposes calcium iodide
when heated. As in the case of hydrogen, potassium, and magnesium, the
affinity of iodine for calcium is feebler than that of chlorine (and
oxygen), and therefore it is not surprising that calcium iodide may be
subjected to that decomposition, which the chloride and oxide undergo
with difficulty.[37] _Metallic calcium_ is of a yellow colour, and has a
considerable lustre, which it preserves in dry air. Its specific gravity
is 1·58. Calcium is distinguished by its great ductility; it melts at a
red heat and then burns in the air with a very brilliant flame; the
brilliancy is due to the formation of finely divided infusible calcium
oxide. Judging from the fact that calcium in burning gives a very large
flame, it is probable that this metal is volatile. Calcium decomposes
water at the ordinary temperature, and is oxidised in moist air, but not
so rapidly as sodium. In burning, it gives its oxide or _lime_, CaO, a
substance which is familiar to every one, and of which we have already
frequently had occasion to speak. This oxide is not met with in nature in
a free state, because it is an energetic base which everywhere encounters
acid substances forming salts with them. It is generally combined with
silica, or occurs as calcium carbonate or sulphate. The carbonate and
nitrate are decomposed, at a red heat, with the formation of lime. As a
rule, the carbonate, which is so frequently met with in nature, serves as
the source of the calcium oxide, both commercial and pure. When heated,
calcium carbonate dissociates: CaCO_{3} = CaO + CO_{2}. In practice the
decomposition is conducted at a bright red heat, in the presence of
steam, or a current of a foreign gas, in heaps or in special kilns.[38]

  [35] There are many other methods of separating calcium from magnesium
       besides that mentioned above (Note 22). Among them it will be
       sufficient to mention the behaviour of these bases towards a
       solution of sugar; hydrated _lime_ is exceedingly _soluble in an
       aqueous solution of sugar_, whilst magnesia is but little soluble.
       All the lime may be extracted from dolomite by burning it, slaking
       the mixture of oxides thus obtained, and adding a 10 p.c. solution
       of sugar. Carbonic anhydride precipitates calcium carbonate from
       this solution. The addition of sugar (molasses) to the lime used
       for building purposes powerfully increases the binding power of
       the mortar, as I have myself found. I have been told that in the
       East (India, Japan) the addition of sugar to cement has long been
       practised.

  [36] Moreover Caron obtained an alloy of calcium and zinc by fusing
       calcium chloride with zinc and sodium. The zinc distilled from
       this alloy at a white heat, leaving calcium behind (Note 50).

  [37] Calcium iodide may be prepared by saturating lime with hydriodic
       acid. It is a very soluble salt (at 20° one part of the salt
       requires 0·49 part and at 43° 0·35 part of water for solution), is
       deliquescent in the air, and resembles calcium chloride in many
       respects. It changes but little when evaporated, and like calcium
       chloride fuses when heated, and therefore all the water may be
       driven off by heat. If anhydrous calcium iodide be heated with an
       equivalent quantity of sodium in a closely covered iron crucible,
       sodium iodide and metallic calcium are formed (Liés-Bodart). Dumas
       advises carrying on this reaction in a closed space under
       pressure.

  [38] Kilns which act either intermittently or continuously are built
       for this purpose. Those of the first kind are filled with
       alternate layers of fuel and limestone; the fuel is lighted, and
       the heat developed by its combustion serves for decomposing the
       limestone. When the process is completed the kiln is allowed to
       cool somewhat, the lime raked out, and the same process repeated.
       In the continuously acting furnaces, constructed like that shown
       in fig. 78, the kiln itself only contains limestone, and there are
       lateral hearths for burning the fuel, whose flame passes through
       the limestone and serves for its decomposition. Such furnaces are
       able to work continuously, because the unburnt limestone may be
       charged from above and the burnt lime raked out from below. It is
       not every limestone that is suitable for the preparation of lime,
       because many contain impurities, principally clay, dolomite, and
       sand. Such limestones when burnt either fuse partially or give an
       impure lime, called _poor_ lime in distinction from that obtained
       from purer limestone, which is called _rich_ lime. The latter kind
       is characterised by its disintegrating into a fine powder when
       treated with water, and is suitable for the majority of uses to
       which lime is applied, and for which the poor lime is sometimes
       quite unfit. However, certain kinds of poor lime (as we shall see
       in Chapter XVIII., Note 25) are used in the preparation of
       hydraulic cements, which solidify into a hard mass under water.

       In order to obtain perfectly pure lime it is necessary to take the
       purest possible materials. In the laboratory, marble or shells are
       used for this purpose as a pure form of calcium carbonate. They
       are first burnt in a furnace, then put in a crucible and moistened
       with a small quantity of water, and finally strongly ignited, by
       which means a pure lime is obtained. Pure lime may be more rapidly
       prepared by taking calcium nitrate, CaN_{2}O_{6}, which is easily
       obtained by dissolving limestone in nitric acid. The solution
       obtained is boiled with a small quantity of lime in order to
       precipitate the foreign oxides which are insoluble in water. The
       oxides of iron, aluminium, &c., are precipitated by this means.
       The salt is then crystallised and ignited: CaN_{2}O_{6} = CaO +
       2NO_{2} + O.

       In the decomposition of calcium carbonate the lime preserves the
       form of the lumps subjected to ignition; this is one of the signs
       distinguishing quicklime when it is freshly burnt and unaltered by
       air. It attracts moisture from the air and then disintegrates to a
       powder; if left long exposed in the air, it also attracts carbonic
       anhydride and increases in volume; it does not entirely pass into
       carbonate, but forms a compound of the latter with caustic lime.

[Illustration: FIG. 78.--Continually-acting kiln for burning lime. The
lime is charged from above and calcined by four lateral grates, R, M. D,
fire-bars. B, space for withdrawing the burnt lime. K, stoke-house. M.
fire grate. Q, R, under-grate.]

Calcium oxide--that is, quicklime--is a substance (sp. gr. 3·15) which
is unaffected by heat,[39] and may therefore serve as a fire-resisting
material, and was employed by Deville for the construction of furnaces in
which platinum was melted, and silver volatilised by the action of the
heat evolved by the combustion of detonating gas. The hydrated lime,
slaked lime, or calcium hydroxide, CaH_{2}O_{2} (specific gravity 2·07)
is a most common alkaline substance, employed largely in building for
making mortars or cements, in which case its binding property is mainly
due to the absorption of carbonic anhydride.[40] Lime, like other
alkalis, acts on many animal and vegetable substances, and for this
reason has many practical uses--for example, for removing fats, and in
agriculture for accelerating the decomposition of organic substances in
the so-called _composts_ or accumulations of vegetable and animal remains
used for fertilising land. Calcium hydroxide easily loses its water at a
moderate heat (530°), but it does not part with water at 100°. When mixed
with water, lime forms a pasty mass known as _slaked lime_ and in a more
dilute form as _milk of lime_, because when shaken up in water it remains
suspended in it for a long time and presents the appearance of a milky
liquid. But, besides this, lime is directly soluble in water, not to any
considerable extent, but still in such a quantity that _lime water_ is
precipitated by carbonic anhydride, and has clearly distinguishable
alkaline properties. One part of lime requires at the ordinary
temperature about 800 parts of water for solution. At 100° it requires
about 1500 parts of water, and therefore lime-water becomes cloudy when
boiled. If lime-water be evaporated in a vacuum, calcium hydroxide
separates in six-sided crystals.[41] If lime-water be mixed with hydrogen
peroxide minute crystals of _calcium peroxide_, CaO_{2},8H_{2}O,
separate; this compound is very unstable and, like barium peroxide, is
decomposed by heat. Lime, as a powerful base, combines with all acids,
and in this respect presents a transition from the true alkalis to
magnesia. Many of the salts of calcium (the carbonate, phosphate,
borate, and oxalate) are insoluble in water; besides which the sulphate
is only sparingly soluble. As a more energetic base than magnesia, lime
forms salts, CaX_{2}, which are distinguished by their stability in
comparison with the salts MgX_{2}; neither does lime so easily form basic
and double salts as magnesia.

  [39] Lime, when raised to a white heat in the vapour of potassium,
       gives calcium, and in chlorine it gives off oxygen. Sulphur,
       phosphorus, &c., when heated with lime, are absorbed by it.

  [40] The greater quantity of lime is used in making mortar for binding
       bricks or stones together, in the form of _lime_ or _cement_, or
       the so-called _slaked lime_. For this purpose the lime is mixed
       with water and sand, which serves to separate the particles of
       lime from each other. If only lime paste were put between two
       bricks they would not hold firmly together, because after the
       water had evaporated the lime would occupy a smaller space than
       before, and therefore cracks and powder would form in its mass, so
       that it would not at all produce that complete cementation of the
       bricks which it is desired to attain. Pieces of stone--that is,
       sand--mixed with the lime hinder this process of disintegration,
       because the lime binds together the individual grains of sand
       mixed with it, and forms one concrete mass, in consequence of a
       process which proceeds after the desiccation or removal of the
       water. The process of the solidification of lime, taken as slaked
       lime, consists first in the direct evaporation of the water and
       crystallisation of the hydrate, so that the lime binds the stones
       and sand mixed with it, just as glue binds two pieces of wood. But
       this preliminary binding action of lime is feeble (as is seen by
       direct experiment) unless there be further alteration of the lime
       leading to the formation of carbonates, silicates, and other salts
       of calcium which are distinguished by their great cohesiveness.
       With the progress of time the cement is partially subjected to the
       action of the carbonic anhydride in the air, owing to which
       calcium carbonate is formed, but not more than half the lime is
       thus converted into carbonate. Besides which, the lime partially
       acts on the silica of the bricks, and it is owing to these new
       combinations simultaneously forming in the cement that it
       gradually becomes stronger and stronger. Hence the binding action
       of the lime becomes stronger with the lapse of time. This is the
       reason (and not, as is sometimes said, because the ancients knew
       how to build stronger than we do) why buildings which have stood
       for centuries possess a very strongly binding cement. Hydraulic
       cements will be described later (Chapter XVIII., Note 25).

  [41] Professor Glinka measured the transparent bright crystals of
       calcium hydroxide which are formed in common hydraulic (Portland)
       cement.

Anhydrous lime does not absorb dry carbonic anhydride at the ordinary
temperature. This was already known by Scheele, and Prof. Schuliachenko
showed that there is no absorption even at 360°. It only proceeds at a
red heat,[42] and then only leads to the formation of a mixture of
calcium oxide and carbonate (Rose). But if the lime be slaked or
dissolved, the absorption of carbonic anhydride proceeds rapidly and
completely. These phenomena are connected with the _dissociation of
calcium carbonate_, studied by Debray (1867) under the influence of the
conceptions of dissociation introduced into science by Henri Saint-Claire
Deville. Just as there is no vapour tension for non-volatile substances,
so there is no dissociation tension of carbonic anhydride for calcium
carbonate at the ordinary temperature. Just as every volatile substance
has a maximum possible vapour tension for every temperature, so also
calcium carbonate has its corresponding _dissociation tension_; this at
770° (the boiling point of cadmium) is about 85 mm. (of the mercury
column), and at 930° (the boiling point of Zn) it is about 520 mm. As, if
the tension be greater, there will be no evaporation, so also there will
he no decomposition. Debray took crystals of calc spar, and could not
observe the least change in them at the boiling point of zinc (930°) in
an atmosphere of carbonic anhydride taken at the atmospheric pressure
(760 mm.), whilst on the other hand calcium carbonate may be completely
decomposed at a much lower temperature if the tension of the carbonic
anhydride be kept below the dissociation tension, which may be done
either by directly pumping away the gas with an air-pump, or by mixing it
with some other gas--that is, by diminishing the partial pressure of the
carbonic anhydride,[43] just as an object may be dried at the ordinary
temperature by removing the aqueous vapour or by carrying it off in a
stream of another gas. Thus it is possible to obtain calcium carbonate
from lime and carbonic anhydride at a certain temperature above that at
which dissociation begins, and conversely to decompose calcium carbonate
at the same temperature into lime and carbonic anhydride.[44] At the
ordinary temperature the reaction of the first order (combination) cannot
proceed because the second (decomposition, dissociation) cannot take
place, and thus all the most important phenomena with respect to the
behaviour of lime towards carbonic anhydride are explained by starting
from one common basis.[45]

  [42] The act of heating brings the substance into that state of
       internal motion which is required for reaction. It should be
       considered that by the act of heating not only is the bond between
       the parts, or cohesion of the molecules, altered (generally
       diminished), not only is the motion or store of energy of the
       whole molecule increased, but also that in all probability the
       motion of the atoms themselves in molecules undergoes a change.
       The same kind of change is accomplished by the act of solution, or
       of combination in general, judging from the fact that a dissolved
       or combined substance--for instance, lime with water--reacts on
       carbonic anhydride as it does under the action of heat. For the
       comprehension of chemical phenomena it is exceedingly useful to
       recognise clearly this parallelism. Rose's observation on the
       formation (by the slow diffusion of solutions of calcium chloride
       and sodium carbonate) of aragonite from dilute, and of calc spar
       from strong, solutions is easily understood from this point of
       view. As aragonite is always formed from hot solutions, it appears
       that dilution with water acts like heat. The following experiment
       of Kühlmann is particularly instructive in this sense. Anhydrous
       (perfectly dry) barium oxide does not react with monohydrated
       sulphuric acid, H_{2}SO_{4} (containing neither free water nor
       anhydride, SO_{3}). But if either an incandescent object or a
       moist substance is brought into contact with the mixture a violent
       reaction immediately begins (it is essentially the same as
       combustion), and the whole mass reacts.

       The influence of solution on the process of reaction is
       instructively illustrated by the following experiment. Lime, or
       barium oxide, is placed in a flask or retort having an upper
       orifice and connected with a tube immersed in mercury. A funnel
       furnished with a stopcock and filled with water is fixed into the
       upper orifice of the retort, which is then filled with dry
       carbonic anhydride. There is no absorption. When a constant
       temperature is arrived at, the unslaked oxide is made to absorb
       all the carbonic anhydride by carefully admitting water. A vacuum
       is formed, as is seen by the mercury rising in the neck of the
       retort. With water the absorption goes on to the end, whilst under
       the action of heat there remains the dissociating tension of the
       carbonic anhydride. Furthermore, we here see that, with a certain
       resemblance, there is also a distinction, depending on the fact
       that at low temperatures calcium carbonate does not dissociate;
       this determines the complete absorption of the carbonic anhydride
       in the aqueous solution.

  [43] Experience has shown that by moistening partially-burnt lime with
       water and reheating it, it is easy to drive off the last traces of
       carbonic anhydride from it, and that, in general, by blowing air
       or steam through the lime, and even by using moist fuel, it is
       possible to accelerate the decomposition of the calcium carbonate.
       The partial pressure is decreased by these means.

  [44] Before the introduction of Deville's theory of dissociation, the
       _modus operandi_ of decompositions like that under consideration
       was understood in the sense that decomposition starts at a certain
       temperature, and that it is accelerated by a rise of temperature,
       but it was not considered possible that combination could proceed
       at the same temperature as that at which decomposition goes on.
       Berthollet and Deville introduced the conception of equilibrium
       into chemical science, and elucidated the question of reversible
       reactions. Naturally the subject is still far from being
       clear--the questions of the rate and completeness of reaction, of
       contact, &c., still intrude themselves--but an important step has
       been made in chemical mechanics, and we have started on a new path
       which promises further progress, towards which much has been done
       not only by Deville himself, but more especially by the French
       chemists Debray, Troost, Lemoine, Hautefeuille, Le Chatelier, and
       others. Among other things those investigators have shown the
       close resemblance between the phenomena of evaporation and
       dissociation, and pointed out that the amount of heat absorbed by
       a dissociating substance may be calculated according to the law of
       the variation of dissociation-pressure, in exactly the same manner
       as it is possible to calculate the latent heat of the evaporation
       of water, knowing the variation of the tension with the
       temperature, on the basis of the second law of the mechanical
       theory of heat. Details of this subject must be looked for in
       special works on physical chemistry. _One and the same conception_
       of the mechanical theory of heat _is applicable to dissociation_
       and _evaporation_.

  [45] But the question as to the formation of a basic calcium carbonate
       with a rise of temperature still remains undecided. The presence
       of water complicates all the relations between lime and carbonic
       anhydride, all the more as the existence of an attraction between
       calcium carbonate and water is seen from its being able to give a
       _crystallo-hydrate_, CaCO_{3},5H_{2}O (Pelouze), which
       crystallises in rhombic prisms of sp. gr. about 1·77 and loses its
       water at 20°. These crystals are obtained when a solution of lime
       in sugar and water is left long exposed to the air and slowly
       attracts carbonic anhydride from it, and also by the evaporation
       of such a solution at a temperature of about 3°. On the other
       band, it is probable that an _acid salt_, CaH_{2}(CO_{3})_{2}, is
       formed in an aqueous solution, not only because water containing
       carbonic acid dissolves calcium carbonate, but more especially in
       view of the researches of Schloesing (1872), which showed that at
       16° a litre of water in an atmosphere of carbonic anhydride
       (pressure 0·984 atmosphere) dissolves 1·086 gram of calcium
       carbonate and 1·778 gram of carbonic anhydride, which corresponds
       with the formation of calcium hydrogen carbonate, and the solution
       of carbonic anhydride in the remaining water. Caro showed that a
       litre of water is able to dissolve as much as 3 grams of calcium
       carbonate if the pressure be increased to 4 and more atmospheres.
       The calcium carbonate is precipitated when the carbonic anhydride
       passes off in the air or in a current of another gas; this also
       takes place in many natural springs. Tufa, stalactites, and other
       like formations from waters containing calcium carbonate and
       carbonic acid in solution are formed in this manner. The
       solubility of calcium carbonate itself at the ordinary temperature
       does not exceed 13 milligrams per litre of water.

_Calcium carbonate_, CaCO_{3}, is sometimes met with in nature in a
crystalline form, and it forms an example of the phenomenon termed
_dimorphism_--that is, it appears in two crystalline forms. When it
exhibits combinations of forms belonging to the hexagonal system
(six-sided prisms, rhombohedra, &c.) it is called _calc spar_. Calc spar
has a specific gravity of 2·7, and is further characterised by a distinct
cleavage along the planes of the fundamental rhombohedron having an angle
of 105°. Perfectly transparent Iceland spar presents a clear example of
double refraction (for which reason it is frequently employed in physical
apparatus). The other form of calcium carbonate occurs in crystals
belonging to the rhombic system, and it is then called _aragonite_; its
specific gravity is 3·0. If calcium carbonate be artificially produced by
slow crystallisation at the ordinary temperature, it appears in the
rhombohedral form, but if the crystallisation be aided by heat it then
appears as aragonite. It may therefore be supposed that calc spar
presents the form corresponding with a low temperature, and aragonite
with a higher temperature during crystallisation.[46]

  [46] Dimorphous bodies differ from true isomers and polymers in that
       they do not differ in their chemical reactions, which are
       determined by a difference in the distribution (motion) of the
       atoms in the molecules, and therefore dimorphism is usually
       ascribed to a difference in the distribution of similar molecules,
       building up a crystal. Although such a hypothesis is quite
       admissible in the spirit of the atomic and molecular theory, yet,
       as in such a redistribution of the molecules a perfect
       conservation of the distribution of the atoms in them cannot be
       imagined, and in every effort of chemical reaction there must take
       place a certain motion among the atoms; so in my opinion there is
       no firm basis for distinguishing dimorphism from the general
       conception of isomerism, under which the cases of those organic
       bodies which are dextro and lævo rotatory (with respect to
       polarised light) have recently been brought with such brilliant
       success. When calcium carbonate separates out from solutions, it
       has at first a gelatinous appearance, which leads to the
       supposition that this salt appears in a colloidal state. It only
       crystallises with the progress of time. The colloidal state of
       calcium carbonate is particularly clear from the following
       observations made by Prof. Famintzin, who showed that when it
       separates from solutions it is obtained under certain conditions
       in the form of grains having the peculiar paste-like structure
       proper to starch, which fact has not only an independent interest,
       but presents an example of a mineral substance being obtained in a
       form until then only known in the organic substances elaborated in
       plants. This shows that the forms (cells, vessels, &c.) in which
       vegetable and animal substances occur in organisms do not present
       in themselves anything peculiar to organisms, but are only the
       result of those particular conditions in which these substances
       are formed. Traube and afterwards Monnier and Vogt (1882) obtained
       formations which, under the microscope, were in every respect
       identical in appearance with vegetable cells, by means of a
       similar slow formation of precipitates (by reacting on sulphates
       of different metals with sodium silicate or carbonate).

_Calcium sulphate_ in combination with two equivalents of water,
CaSO_{4},2H_{2}O, is very widely distributed in nature, and is known as
_gypsum_. Gypsum loses one and a half and two equivalents of water at a
moderate temperature,[47] and anhydrous or burnt gypsum is then obtained,
which is also known as plaster of Paris, and is employed in large
quantities for modelling.[48] This use depends on the fact that burnt and
finely-divided and sifted gypsum forms a paste when mixed with water;
after a certain time this paste becomes slightly heated and solidifies,
owing to the fact that the anhydrous calcium sulphate, CaSO_{4}, again
combines with water. When the plaster of Paris and water are first made
into a paste they form a mechanical mixture, but when the mass
solidifies, then a compound of the calcium sulphate with two molecules of
water is produced; and this may be regarded as derived from S(OH)_{6} by
the substitution of two atoms of hydrogen by one atom of bivalent
calcium. Natural gypsum sometimes appears as perfectly colourless, or
variegated, marble-like, masses, and sometimes in perfectly colourless
crystals, _selenite_, of specific gravity 2·33. The semi-transparent
gypsum, or _alabaster_, is often carved into small statues. Besides
which an anhydrous calcium sulphate, CaSO_{4}, called _anhydrite_
(specific gravity 2·97), occurs in nature. It sometimes occurs along with
gypsum. It is no longer capable of combining directly with water, and
differs in this respect from the anhydrous salt obtained by gently
igniting gypsum. If gypsum be very strongly heated it shrinks and loses
its power of combining with water.[48 bis] One part of calcium sulphate
requires at 0° 525 parts of water for solution, at 38° 466 parts, and at
100° 571 parts of water. The maximum solubility of gypsum is at about
36°, which is nearly the same temperature as that at which sodium
sulphate is most soluble.[49]

  [47] According to Le Chatelier (1888), 1-1/2H_{2}O is lost at
       120°--that is, H_{2}O,2CaSO_{4} is formed, but at 194° all the
       water is expelled. According to Shenstone and Cundall (1888)
       gypsum begins to lose water at 70° in dry air. The semi-hydrated
       compound H_{2}O,2CaSO_{4} is also formed when gypsum is heated
       with water in a closed vessel at 150° (Hoppe-Seyler).

  [48] For stucco-work it is usual to add lime and sand, as the mass is
       then harder and does not solidify so quickly. For imitating
       marble, glue is added to the plaster, and the mass is polished
       when thoroughly dry. Re-burnt gypsum cannot be used over again, as
       that which has once solidified is, like the natural anhydride, not
       able to recombine with water. It is evident that the structure of
       the molecules in the crystallised mass, or in general in any dense
       mass, exerts an influence on the chemical action, which is more
       particularly evident in metals in their different forms (powder,
       crystalline, rolled, &c.)

  [48 bis] According to MacColeb, gypsum dehydrated at 200° has a
       specific gravity 2·577, and heated to its point of fusion, 2·654.
       Potilitzin (1894) also admits the two above-named modifications of
       anhydrous gypsum, which, moreover, always contain the
       semi-hydrated hydrate (Note 47), and he explains by their relation
       to water the phenomena observed in the solidification of a mixture
       of burnt gypsum and water.

  [49] As Marignac showed, gypsum, especially when desiccated at 120°,
       easily gives supersaturated solutions with respect to
       CaSO_{4},2H_{2}O, which contain as much as 1 part of CaSO_{4} to
       110 parts of water. Boiling dilute hydrochloric acid dissolves
       gypsum, forming calcium chloride. The behaviour of gypsum towards
       the alkaline carbonates has been described in Chapter X. Alcohol
       precipitates gypsum from its aqueous solutions, because, like the
       sulphates in general, it is sparingly soluble in alcohol. Gypsum,
       like all the sulphates, when heated with charcoal, gives up its
       oxygen, forming the sulphide, CaS.

       Calcium ulphate, like magnesium sulphate, is capable of forming
       double salts, but with difficulty, and they are chemically less
       stable. They contain, as is always the case with double salts,
       less water of crystallisation than the component salts. Rose,
       Struvé, and others obtained the salt CaK_{2}(SO_{4})_{2},H_{2}O; a
       mixture of gypsum with an equivalent amount of potassium sulphate
       and water solidifies into a homogeneous mass. Fritzsche obtained
       the corresponding sodium salt in a hydrated and anhydrous state,
       by heating a mixture of gypsum with a saturated solution of sodium
       sulphate. The anhydrous salt occurs in nature as _glauberite_.
       Fritzsche also obtained _gaylussite_,
       Na_{2}Ca(CO_{3})_{2},5H_{2}O, by pouring a saturated solution of
       sodium carbonate on to freshly-precipitated calcium carbonate.
       Calcium also forms basic salts, but only a few. Veeren (1892)
       obtained Ca(NO_{3})_{2}Ca(OH)_{2},2-1/2H_{2}O by leaving powdered
       caustic lime in a saturated solution of Ca(NO_{3})_{2} until it
       solidified. This salt is decomposed by water.

As lime is a more energetic base than magnesia, so _calcium chloride_,
CaCl_{2}, is not so easily decomposed by water, and its solutions only
disengage a small quantity of hydrochloric acid when evaporated, and when
the evaporation is conducted in a stream of hydrochloric acid it easily
gives an anhydrous salt which fuses at 719°; otherwise an aqueous
solution yields a crystallo-hydrate, CaCl_{2},6H_{2}O, which melts at
30°.[50]

  [50] Calcium chloride has a specific gravity 2·20, or, when fused,
       2·12, and the sp. gr. of the crystallised salt CaCl_{2},6H_{2}O is
       1·69. If the volume of the crystals at 0° = 1, then at 29° it is
       1·020, and the volume of the fused mass at the same temperature is
       1·118 (Kopp) (specific gravity of solutions, _see_ Note 27). The
       solution containing 50 p.c. CaCl_{2} boils at 130°, 70 p.c. at
       158°. Superheated steam decomposes calcium chloride with more
       difficulty than magnesium chloride and with greater ease than
       barium chloride (Kuhnheim). Sodium does not decompose fused
       calcium chloride even on prolonged heating (Liés-Bodart), but an
       alloy of sodium with zinc, lead, and bismuth decomposes it,
       forming an alloy of calcium with one of the above-named metals
       (Caron). The zinc alloy may be obtained with as much as 15 p.c. of
       calcium. Calcium chloride is soluble in alcohol and absorbs
       ammonia.

       A gram molecular weight of calcium chloride in dissolving in an
       excess of water evolves 18,723 calories, and in dissolving in
       alcohol 17,555 units of heat, according to Pickering.

       Roozeboom made detailed researches on the crystallo-hydrates of
       calcium chloride (1889), and found that CaCl_{2},6H_{2}O melts at
       30·2°, and is formed at low temperatures from solutions containing
       not more than 103 parts of calcium chloride per 100 parts of
       water; if the amount of salt (always to 100 parts of water)
       reaches 120 parts, then tabular crystals of
       CaCl_{2},4H_{2}O[Greek: b] are formed, which at temperatures above
       38·4° are converted into the crystallo-hydrates CaCl_{2},2H_{2}O,
       whilst at temperatures below 18° the [Greek: b] variety passes
       into the more stable CaCl_{2},4H_{2}O[Greek: a], which process is
       aided by mechanical friction. Hence, as is the case with magnesium
       sulphate (Note 27), one and the same crystallo-hydrate appears in
       two forms--the [Greek: b], which is easily produced but is
       unstable, and the [Greek: a], which is stable. The solubility of
       the above-mentioned hydrates of chloride of calcium, or amount of
       calcium chloride per 100 parts of water, is as follows:--

                                      0°  20°   30°    40°  60°

          CaCl_{2},6H_{2}O            60   75   100     (102·8)
          CaCl_{2},4H_{2}O[Greek: a]  --   90   101   117} (154·2)
          CaCl_{2},4H_{2}O[Greek: b]  --  104   114    --}
          CaCl_{2},2H_{2}O            --   -- (308·3) 128  137

       The amount of calcium chloride to 100 parts of water in the
       crystallo-hydrate is given in brackets. The point of intersection
       of the curves of solubility lies at about 30° for the first two
       salts and about 45° for the salts with 4H_{2}O and 2H_{2}O. The
       crystals CaCl_{2},2H_{2}O may, however, be obtained (Ditte) at the
       ordinary temperature from solutions containing hydrochloric acid.
       The vapour tension of this crystallo-hydrate equals the
       atmospheric at 165°, and therefore the crystals may be dried in an
       atmosphere of steam and obtained without a mother liquor, whose
       vapour tension is greater. This crystallo-hydrate decomposes at
       about 175° into CaCl_{2},H_{2}O and a solution; this is easily
       brought about in a closed vessel when the pressure is greater than
       the atmosphere. This crystallo-hydrate is destroyed at
       temperatures above 260°, anhydrous calcium chloride being formed.

       Neglecting the unstable modification CaCl_{2},4H_{2}O[Greek: b],
       we will give the temperatures _t_ at which the passage of one
       hydrate into another takes place and at which the solution
       CaCl_{2} + _n_H_{2}O, the two solids A and B and aqueous vapour,
       whose tension is given as _p_ in millimetres, are able to exist
       together in stable equilibrium, according to Roozeboom's
       determinations:

        _t_       _n_            A                 B           _p_
       -55°      14·5           ice        CaCl_{2},6H_{2}O     0
       +29·8°     6·1    CaCl_{2},6H_{2}O  CaCl_{2},4H_{2}O     6·8
        45·3°     4·7    CaCl_{2},4H_{2}O  CaCl_{2},2H_{2}O    11·8
       175·5°     2·1    CaCl_{2},2H_{2}O  CaCl_{2},H_{2}O    842
       260°       1·8    CaCl_{2},H_{2}O   CaCl_{2}           Several
                                                            atmospheres

       Solutions of calcium chloride may serve as a convenient example
       for the study of the supersaturated state, which in this case
       easily occurs, because different hydrates are formed. Thus at 25°
       solutions containing more than 83 parts of anhydrous calcium
       chloride per 100 of water will be supersaturated for the hydrate
       CaCl_{2},6H_{2}O.

       On the other hand, Hammerl showed that solutions of calcium
       chloride, when frozen, deposit ice if they contain less than 43
       parts of salt per 100 of water, and if more the crystallo-hydrate
       CaCl_{2},6H_{2}O separates, and that a solution of the above
       composition (CaCl_{2},14H_{2}O requires 44·0 parts calcium
       chloride per 100 of water) solidifies as a cryohydrate at
       about -55°.

Just as for potassium, K = 39 (and sodium, Na = 23), there are the near
analogues, Rb = 85 and Cs = 133, and also another, Li = 7, so in exactly
the same manner for calcium, Ca = 40 (and magnesium, Mg = 24), there is
another analogue of lighter atomic weight, beryllium, Be = 9, besides the
near analogues strontium, Sr = 87, and barium, Ba = 137. As rubidium and
cæsium are more rarely met with in nature than potassium, so also
strontium and barium are rarer than calcium (in the same way that bromine
and iodine are rarer than chlorine). Since they exhibit many points of
resemblance with calcium, strontium and barium may be characterised after
a very short acquaintance with their chief compounds; this shows the
important advantages gained by distributing the elements according to
their natural groups, to which matter we shall turn our attention in the
next chapter.

Among the compounds of barium met with in nature the commonest is the
_sulphate_, BaSO_{4}, which forms anhydrous crystals of the rhombic
system, which are identical in their crystalline form with anhydrite,
and generally occur as transparent and semi-transparent masses of
tabular crystals having a high specific gravity, namely 4·45, for which
reason this salt bears the name of _heavy spar_ or _barytes_. Analogous
to it is _celestine_, SrSO_{4}, which is, however, more rarely met with.
Heavy spar frequently forms the gangue separated on dressing metallic
ores from the vein stuff; this mineral is the source of all other barium
compounds; for the carbonate, although more easily transformed into the
other compounds (because acids act directly on it, evolving carbonic
anhydride), is a comparatively rare mineral (BaCO_{3} forms the mineral
_witherite_; SrCO_{3}, _strontianite_; both are rare, the latter is
found at Etna). The treatment of barium sulphate is rendered difficult
from the fact that it is insoluble both in water and acids, and has
therefore to be treated by a method of reduction.[51] Like sodium
sulphate and calcium sulphate, heavy spar when heated with charcoal
parts with its oxygen and forms barium sulphide, BaS. For this purpose a
pasty mixture of powdered heavy spar, charcoal, and tar is subjected to
the action of a strong heat, when BaSO_{4} + 4C = BaS + 4CO. The residue
is then treated with water, in which the barium sulphide is soluble.[52]
When boiled with hydrochloric acid, barium chloride, BaCl_{2},
is obtained in solution, and the sulphur is disengaged as gaseous
sulphuretted hydrogen, BaS + 2HCl = BaCl_{2} + H_{2}S. In this manner
barium sulphate is converted into barium chloride,[53] and the latter
by double decomposition with strong nitric acid or nitre gives the less
soluble barium nitrate, Ba(NO_{3})_{2},[54] or with sodium carbonate a
precipitate of barium carbonate, BaCO_{3}. Both these salts are able to
give _barium oxide_, or _baryta_, BaO, and the hydroxide, Ba(HO)_{2},
which differs from lime by its great solubility in water,[55] and by the
ease with which it forms a crystallo-hydrate, BaH_{2}O_{2},8H_{2}O, from
its solutions. Owing to its solubility, baryta is frequently employed
in manufactures and in practical chemistry as an alkali which has the
very important property that it may be always entirely removed from
solution by the addition of sulphuric acid, which entirely separates
it as the insoluble barium sulphate, BaSO_{4}. It may also be removed
whilst it remains in an alkaline state (for example, the excess which
may remain when it is used for saturating acids) by means of carbonic
anhydride, which also completely precipitates baryta as a sparingly
soluble, colourless, and powdery carbonate. Both these reactions show
that baryta has such properties as would very greatly extend its use
were its compounds as widely distributed as those of sodium and calcium,
and were its soluble compounds not poisonous. Barium nitrate is directly
decomposed by the action of heat, barium oxide being left behind. The
same takes place with barium carbonate, especially that form of it
precipitated from solutions, and when mixed with charcoal or ignited
in an atmosphere of steam. Barium oxide combines with water with the
development of a large amount of heat, and the resultant hydroxide is
very stable in its retention of the water, although it parts with it
when strongly ignited.[55 bis] With oxygen the anhydrous oxide gives, as
already mentioned in Chapters III. and IV., a _peroxide_, BaO_{2}.[56]
Neither calcium nor strontium oxides are able to give such a peroxide
directly, but they form peroxides under the action of hydrogen peroxide.

  [51] The action of barium sulphate on sodium and potassium carbonates
       is given on p. 437.

  [52] Barium sulphide is decomposed by water, BaS + 2H_{2}O = H_{2}S
       + Ba(OH)_{2} (the reaction is reversible), but both substances are
       soluble in water, and their separation is complicated by the fact
       that barium sulphide absorbs oxygen and gives insoluble barium
       sulphate. The hydrogen sulphide is sometimes removed from the
       solution by boiling with the oxides of copper or zinc. If sugar be
       added to a solution of barium sulphide, barium saccharate is
       precipitated on heating; it is decomposed by carbonic anhydride,
       so that barium carbonate is formed. An equivalent mixture of
       sodium sulphate with barium or strontium sulphates when ignited
       with charcoal gives a mixture of sodium sulphide and barium or
       strontium sulphide, and if this mixture be dissolved in water and
       the solution evaporated, barium or strontium hydroxide
       crystallises out on cooling, and sodium hydrosulphide, NaHS, is
       obtained in solution. The hydroxides BaH_{2}O_{2} and SrH_{2}O_{2}
       are prepared on a large scale, being applied to many reactions;
       for example, strontium hydroxide is prepared for sugar works for
       extracting crystallisable sugar from molasses.

       We may remark that Boussingault, by igniting barium sulphate in
       hydrochloric acid gas, obtained a complete decomposition, with the
       formation of barium chloride. Attention should also be turned to
       the fact that Grouven, by beating a mixture of charcoal and
       strontium sulphate with magnesium and potassium sulphates, showed
       the easy decomposability depending on the formation of double
       salts, such as SrS,K_{2}S, which are easily soluble in water, and
       give a precipitate of strontium carbonate with carbonic anhydride.
       In such examples as these we see that the force which binds double
       salts may play a part in directing the course of reactions, and
       the number of double salts of silica on the earth's surface shows
       that nature takes advantage of these forces in her chemical
       processes. It is worthy of remark that Buchner (1893), by mixing a
       40 per cent. solution of barium acetate with a 60 per cent.
       solution of sulphate of alumina, obtained a thick glutinous mass,
       which only gave a precipitate of BaSO_{4} after being diluted with
       water.

  [53] Barium sulphate is sometimes converted into barium chloride in the
       following manner: finely-ground barium sulphate is heated with
       coal and manganese chloride (the residue from the manufacture of
       chlorine). The mass becomes semi-liquid, and when it evolves
       carbonic oxide the heating is stopped. The following double
       decompositions proceed during this operation: first the carbon
       takes up the oxygen from the barium sulphate, and gives sulphide,
       BaS, which enters into double decomposition with the chloride of
       manganese, MnCl_{2}, forming manganese sulphide, MnS, which is
       insoluble in water, and soluble barium chloride. This solution is
       easily obtained pure because many foreign impurities, such as
       iron, remain in the insoluble portion with the manganese. The
       solution of barium chloride is chiefly used for the preparation of
       barium sulphate, which is precipitated by sulphuric acid, by which
       means _barium sulphate_ is re-formed as a powder. This salt is
       characterised by the fact that it is unacted on by the majority of
       chemical reagents, is insoluble in water, and is not dissolved by
       acids. Owing to this, artificial barium sulphate forms a permanent
       white paint which is used instead of (and mixed with) white lead,
       and has been termed 'blanc fixé' or 'permanent white.'

       The solution of one part of calcium chloride at 20° requires 1·36
       part of water, the solution of one part of strontium chloride
       requires 1·88 part of water at the same temperature, and the
       solution of barium chloride 2·88 parts of water. The solubility of
       the bromides and iodides varies in the same proportion. The
       chlorides of barium and strontium crystallise out from solution
       with great ease in combination with water; they form
       BaCl_{2},2H_{2}O and SrCl_{2},6H_{2}O. The latter (which separates
       out at 40°) resembles the salts of Ca and Mg in composition, and
       Étard (1892) obtained SrCl_{2},2H_{2}O from solutions at 90-130°.
       We may also observe that the crystallo-hydrates BaBr_{2},H_{2}O
       and BaI_{2},7H_{2}O are known.

  [54] The nitrates Sr(NO_{3})_{2} (in the cold its solutions give a
       crystallo-hydrate containing 4H_{2}O) and Ba(NO_{3})_{2} are so
       very sparingly soluble in water that they separate in considerable
       quantities when a solution of sodium nitrate is added to a strong
       solution of either barium or strontium chloride. They are obtained
       by the action of nitric acid on the carbonates or oxides. 100
       parts of water at 15° dissolve 6·5 parts of strontium nitrate and
       8·2 parts of barium nitrate, whilst more than 300 parts of calcium
       nitrate are soluble at the same temperature. Strontium nitrate
       communicates a crimson coloration to the flame of burning
       substances, and is therefore frequently used for Bengal fire,
       fireworks, and signal lights, for which purpose the salts of
       lithium are still better fitted. Calcium nitrate is exceedingly
       hygroscopic. Barium nitrate, on the contrary, does not show this
       property in the least degree, and in this respect it resembles
       potassium nitrate, and is therefore used instead of the latter for
       the preparation of a gunpowder which is called 'saxifragin powder'
       (76 parts of barium nitrate, 2 parts of nitre, and 22 parts of
       charcoal).

  [55] The dissociation of the crystallo-hydrate of baryta is given in
       Chapter I., Note 65. 100 parts of water dissolve

                          0°   20°  40°  60°   80°
                     BaO  1·5  3·5  7·4  18·8  90·8
                     SrO  0·3  0·7  1·4   3     9


       Supersaturated solutions are easily formed.

       The anhydrous oxide BaO fuses in the oxyhydrogen flame. When
       ignited in the vapour of potassium, the latter takes up the
       oxygen; whilst in chlorine, oxygen is separated and barium
       chloride formed.

  [55 bis] Brugellmann, by heating BaH_{2}O_{2} in a graphite or clay
       crucible, obtained BaO in needles, sp. gr. 5·32, and by heating in
       a platinum crucible--in crystals belonging to the cubical system,
       sp. gr. 5·74. SrO is obtained in the latter form from the nitrate.
       The following are the specific gravities of the oxides from
       different sources:--

                                 MgO       CaO       SrO
               from RN_{2}O_{6}  3·38      3·25      4·75
                 "  RCO_{3}      3·48      3·26      4·45
                 "  RH_{2}O_{2}  3·41      3·25      4·57

  [56] The property of barium oxide of absorbing oxygen when heated, and
       giving the peroxide, BaO_{2}, is very characteristic for this
       oxide (_see_ Chapter III., Note 7). It only belongs to the
       anhydrous oxide. The hydroxide does not absorb oxygen. Peroxides
       of calcium and strontium may be obtained by means of hydrogen
       peroxide. Barium peroxide is insoluble in water, but is able to
       form a hydrate with it, and also to combine with hydrogen
       peroxide, forming a very unstable compound having the composition
       BaH_{2}O_{4} (obtained by Professor Schöne), which in course of
       time evolves oxygen (Chapter IV., Note 21).

Barium oxide is decomposed when heated with potassium; fused barium
chloride is decomposed, as Davy showed, by the action of a galvanic
current, forming metallic _barium_; and Crookes (1862) obtained an
amalgam of barium from which the mercury could easily be driven off,
by heating sodium amalgam in a saturated solution of barium chloride.
Strontium is obtained by the same processes. Both metals are soluble in
mercury, and seem to be non-volatile or only very slightly volatile.
They are both heavier than water; the specific gravity of barium is
3·6, and of strontium 2·5. They both decompose water at the ordinary
temperature, like the metals of the alkalis.

Barium and strontium as saline elements are characterised by their
powerful basic properties, so that they form acid salts with difficulty,
and scarcely form basic salts. On comparing them together and with
calcium, it is evident that the alkaline properties in this group (as
in the group potassium, rubidium, cæsium) increase with the atomic
weight, and this succession clearly shows itself in many of their
corresponding compounds. Thus, for instance, the solubility of the
hydroxides RH_{2}O_{2} and the specific gravity[57] rise in passing from
calcium to strontium and barium, while the solubility of the sulphates
decreases,[58] and therefore in the case of magnesium and beryllium,
as metals whose atomic weights are still less, we should expect the
solubility of the sulphates to be greater, and this is in reality the
case.

  [57] Even in solutions a gradual progression in the increase of the
       specific gravity shows itself, not only for equivalent solutions
       (for instance, RCl_{2} + 200H_{2}O), but even with an equal
       percentage composition, as is seen from the curves giving the
       specific gravity (water 4° = 10,000) at 15° (for barium chloride,
       according to Bourdiakoff's determinations):

              BeCl_{2} : S = 9,992 + 67·21_p_ + 0·111_p_^2
              CaCl_{2} : S = 9,992 + 80·24_p_ + 0·476_p_^2
              SrCl_{2} : S = 9,992 + 85·57_p_ + 0·733_p_^2
              BaCl_{2} : S = 9,992 + 86·56_p_ + 0·813_p_^2

  [58] One part of calcium sulphate at the ordinary temperature requires
       about 500 parts of water for solution, strontium sulphate about
       7,000 parts, barium sulphate about 400,000 parts, whilst beryllium
       sulphate is easily soluble in water.

Just as in the series of the alkali metals we saw the metals potassium,
rubidium, and cæsium approaching near to each other in their properties,
and allied to them two metals having smaller combining weights--namely,
sodium, and the lightest of all, lithium, which all exhibited certain
peculiar characteristic properties--so also in the case of the metals
of the alkaline earths we find, besides calcium, barium, and strontium,
the metal magnesium and also _beryllium_ or _glucinum_. In respect
to the magnitude of its atomic weight, this last occupies the same
position in the series of the metals of the alkaline earths as lithium
does in the series of the alkali metals, for the combining weight of
beryllium, Be or Gl = 9. This combining weight is greater than that
of lithium (7), as the combining weight of magnesium (24) is greater
than that of sodium (23), and as that of calcium (40) is greater than
that of potassium (39), &c.[59] Beryllium was so named because it
occurs in the mineral _beryl_. The metal is also called glucinum (from
the Greek word [Greek: glykys], 'sweet'), because its salts have a
sweet taste. It occurs in beryl, aquamarine, the emerald, and other
minerals, which are generally of a green colour; they are sometimes
found in considerable masses, but as a rule are comparatively rare
and, as transparent crystals, form precious stones. The composition of
beryl and of the emerald is as follows: Al_{2}O_{3},3BeO,6SiO_{2}. The
Siberian and Brazilian beryls are the best known. The specific gravity
of beryl is about 2·7. Beryllium oxide, from the feebleness of its basic
properties, presents an analogy to aluminium oxide in the same way
that lithium oxide is analogous to magnesium oxide.[60] Owing to its
rare occurrence in nature, to the absence of any especially distinct
individual properties, and to the possibility of foretelling them to
a certain extent on the basis of the periodic system of the elements
given in the following chapter, and owing to the brevity of this
treatise, we will not discuss at any length the compounds of beryllium,
and will only observe that their individuality was pointed out in 1798
by Vauquelin, and that metallic beryllium was obtained by Wöhler and
Bussy. Wöhler obtained _metallic beryllium_ (like magnesium) by acting
on beryllium chloride, BeCl_{2}, with potassium (it is best prepared
by fusing K_{2}BeF_{4} with Na). Metallic beryllium has a specific
gravity 1·64 (Nilson and Pettersson). It is very infusible, melting at
nearly the same temperature as silver, which it resembles in its white
colour and lustre. It is characterised by the fact that it is very
difficultly oxidised, and even in the oxidising flame of the blowpipe is
only superficially covered by a coating of oxide; it does not burn in
pure oxygen, and does not decompose water at the ordinary temperature
or at a red heat, but gaseous hydrochloric acid is decomposed by it
when slightly heated, with evolution of hydrogen and development of a
considerable amount of heat. Even dilute hydrochloric acid acts in the
same manner at the ordinary temperature. Beryllium also acts easily on
sulphuric acid, but it is remarkable that neither dilute nor strong
nitric acid acts on beryllium, which seems especially able to resist
oxidising agents. Potassium hydroxide acts on beryllium as on aluminium,
hydrogen being disengaged and the metal dissolved, but ammonia has no
action on it. These properties of metallic beryllium seem to isolate it
from the series of the other metals described in this chapter, but if we
compare the properties of calcium, magnesium, and beryllium we shall see
that magnesium occupies a position intermediate between the other two.
Whilst calcium decomposes water with great ease, magnesium does so with
difficulty, and beryllium not at all. The peculiarities of beryllium
among the metals of the alkaline earths recall the fact that in the
series of the halogens we saw that fluorine differed from the other
halogens in many of its properties and had the smallest atomic weight.
The same is the case with regard to beryllium among the other metals of
the alkaline earths.

  [59] We refer beryllium to the class of the bivalent metals of the
       alkaline earths--that is, we ascribe to its oxide the formula BeO,
       and do not consider it as trivalent (Be = 13·5, Chapter VII., Note
       21), although that view has been upheld by many chemists. The true
       atomic composition of beryllium oxide was first given by the
       Russian chemist, Avdéeff (1819), in his researches on the
       compounds of this metal. He compared the compounds of beryllium to
       those of magnesium, and refuted the notion prevalent at the time,
       of the resemblance between the oxides of beryllium and aluminium,
       by proving that beryllium sulphate presents a greater resemblance
       to magnesium sulphate than to aluminium sulphate. It was
       especially noticed that the analogues of alumina give alums,
       whilst beryllium oxide, although it is a feeble base, easily
       giving, like magnesia, basic and double salts, does not form true
       alums. The establishment of the periodic system of the elements
       (1869), considered in the following chapter, immediately indicated
       that Avdéeff's view corresponded with the truth--that is, that
       beryllium is bivalent, which therefore necessitated the denial of
       its trivalency. This scientific controversy resulted in a long
       series of researches (1870-80) concerning this element, and ended
       in Nilson and Pettersson--two of the chief advocates of the
       trivalency of beryllium--determining the vapour density of
       BeCl_{2} = 40, (Chapter VII., Note 21), which gave an undoubted
       proof of its bivalency (_see_ also Note 3).

  [60] Beryllium oxide, like aluminium oxide, is precipitated from
       solutions of its salts by alkalis as a gelatinous hydroxide,
       BeH_{2}O_{2}, which, like alumina, is soluble in an excess of
       caustic potash or soda. This reaction may be taken advantage of
       for distinguishing and separating beryllium from aluminium,
       because when the alkaline solution is diluted with water and
       boiled, beryllium hydroxide is precipitated, whilst the alumina
       remains in solution. The solubility of the beryllium oxide at once
       clearly indicates its feeble basic properties, and, as it were,
       separates this oxide from the class of the alkaline earths. But on
       arranging the oxides of the above-described metals of the alkaline
       earths according to their decreasing atomic weights we have the
       series

                      BaO,  SrO,  CaO,  MgO,  BeO,

       in which the basic properties and solubility of the oxides
       consecutively and distinctly decrease until we reach a point when,
       had we not known of the existence of the beryllium oxide, we
       should expect to find in its place an oxide insoluble in water and
       of feeble basic properties. If an alcoholic solution of caustic
       potash be saturated with the hydrate of BeO, and evaporated under
       the receiver of an air pump, it forms silky crystals BeK_{2}O_{2}.

       Another characteristic of the salts of beryllium is that they give
       with aqueous ammonia a gelatinous precipitate which is soluble in
       an excess of ammonium carbonate like the precipitate of magnesia;
       in this beryllium oxide differs from the oxide of aluminium.
       Beryllium oxide easily forms a carbonate which is insoluble in
       water, and resembles magnesium carbonate in many respects.
       Beryllium sulphate is distinguished by its considerable solubility
       in water--thus, at the ordinary temperature it dissolves in an
       equal weight of water; it crystallises out from its solutions in
       well-formed crystals, which do not change in the air, and contain
       BeSO_{4},4H_{2}O. When ignited it leaves beryllium oxide, but this
       oxide, after prolonged ignition, is re-dissolved by sulphuric
       acid, whilst aluminium sulphate, after a similar treatment, leaves
       aluminium oxide, which is no longer soluble in acids. With a few
       exceptions, the salts of beryllium crystallise with great
       difficulty, and to a considerable extent resemble the salts of
       magnesium; thus, for instance, beryllium chloride is analogous to
       magnesium chloride. It is volatile in an anhydrous state, and in a
       hydrated state it decomposes, with the evolution of hydrochloric
       acid.

In addition to the above characteristics of the compounds of the metals
of the alkaline earths, we must add that they, like the alkali metals,
combine with nitrogen and hydrogen, and while sodium nitride (obtained
by igniting the amide of sodium, Chapter XII., Note 44 bis) and lithium
nitride (obtained by heating lithium in nitrogen, Chapter XIII., Note
39) have the composition R_{3}N, so the nitrides of magnesium (Note 14),
calcium, strontium, and barium have the composition R_{3}N_{2}, for
example, Ba_{3}N_{2}, as might be expected from the diatomicity of the
metals of the alkaline earths and from the relation of the nitrides to
ammonia, which is obtained from all of these compounds by the action of
water. The _nitrides_ of Ca, Sr, and Ba are formed directly (Maquenne,
1892) by heating the metals in nitrogen. They all have the appearance of
an amorphous powder of dark colour; as regards their reactions, it is
known that besides disengaging ammonia with water, they form cyanides
when heated with carbonic oxide; for instance, Ba_{3}N_{2} + 2CO =
Ba(CN)_{2} + 2BaO.[61]

  [61] Thus in the nitrides of the metals we have substances by means
       of which we can easily obtain from the nitrogen of the air, not
       only ammonia, but also with the aid of CO, by synthesis, a whole
       series of complex carbon and nitrogen compounds.

The metals of the alkaline earths, just like Na and K, absorb hydrogen
under certain conditions, and form pulverulent easily oxidisable
metallic hydrides, whose composition corresponds exactly to that of
Na_{2}H and K_{2}H, with the substitution of K_{2} and Na_{2} by the
atoms Be, Mg, Ca, Sr, and Ba. The _hydrides of the metals of the
alkaline earths_ were discovered by C. Winkler (1891) in investigating
the reducibility of these metals by magnesium. In reducing their oxides
by heating them with magnesium powder in a stream of hydrogen, Winkler
observed that the hydrogen was absorbed (but very slowly), _i.e._ at
the moment of their separation all the metals of the alkaline earths
combine with hydrogen. This absorptive power increases in passing from
Be to Mg, Ca, Sr, and Ba, and the resultant hydrides retain the combined
hydrogen[62] when heated, so that these hydrides are distinguished
for their considerable stability under heat, but they oxidise very
easily.[63]

  [62] As the hydrides of calcium, magnesium, &c. are very stable under
       the action of heat, and these metals and hydrogen occur in the
       sun, it is likely that the formation of their hydrides may take
       place there. (Private communication from Prof. Winkler, 1894.) It
       is probable that in the free metals of the alkaline earths
       hitherto obtained a portion was frequently in combination with
       nitrogen and hydrogen.

  [63] Thus, for instance, a mixture of 56 parts of CaO and 24 parts of
       magnesium powder is heated in an iron pipe (placed over a row of
       gas burners as in the combustion furnace used for organic
       analysis) in a stream of hydrogen. After being heated for 1/2 hour
       the mixture is found to absorb hydrogen (it no longer passes over
       the mixture, but is retained by it). The product, which is light
       grey, and slightly coherent, disengages a mass of hydrogen when
       water is poured over it, and burns when heated in air. The
       resultant mass contains 33 per cent. CaH, about 28 per cent. CaO,
       and about 38 per cent. MgO. Neither CaH nor any other MH has yet
       been obtained in a pure state.

       The acetylene derivatives of the metals of the alkaline earths
       C_{2}M (Chapter VIII., Note 12 bis), for instance, C_{2}Ba,
       obtained by Maquenne and Moissan, belong to the same class of
       analogous compounds. It must here be remarked that the oxides MO
       of the metals of the alkaline earths, although not reducible by
       carbon at a furnace heat, yet under the action of the heat
       attained in electrical furnaces, not only give up their oxygen to
       carbon (probably partly owing to the action of the current), but
       also combine with carbon. The resultant compounds, C_{2}M, evolve
       acetylene, C_{2}H_{2}, with HCl, just as N_{2}M_{3} give ammonia.
       We may remark moreover that the series of compounds of the metals
       of the alkaline earths with hydrogen, nitrogen and carbon is a
       discovery of recent years, and that probably further research will
       give rise to similar unexpected compounds, and by extending our
       knowledge of their reactions prove to be of great interest.

Thus the analogies and correlation of the metals of these two groups
are now clearly marked, not only in their behaviour towards oxygen,
chlorine, acids, &c., but also in their capability of combining with
nitrogen and hydrogen.

                    *       *       *       *       *

                         END OF THE FIRST VOLUME

                               PRINTED BY

                 SPOTTISWOODE AND CO., NEW-STREET SQUARE

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                                    A
                          CLASSIFIED CATALOGUE
                                   OF
                            SCIENTIFIC WORKS

                              PUBLISHED BY
                     MESSRS. LONGMANS, GREEN, & CO.

                    LONDON: 39 PATERNOSTER ROW, E.C.
                     NEW YORK: 91 & 93 FIFTH AVENUE.
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                                CONTENTS.
                                                                    PAGE
  _ADVANCED SCIENCE MANUALS_                                          30
  AGRICULTURE                                                         27
  ASTRONOMY                                                           14
  BACTERIOLOGY                                                        25
  BIOLOGY                                                             25
  BOTANY                                                              26
  BUILDING CONSTRUCTION                                               10
  CHEMISTRY                                                            2
  DYNAMICS                                                             6
  ELECTRICITY                                                         11
  _ELEMENTARY SCIENCE MANUALS_                                        30
  ENGINEERING                                                         12
  GEOLOGY                                                             17
  HEALTH AND HYGIENE                                                  17
  HEAT                                                                 8
  HYDROSTATICS                                                         6
  LIGHT                                                                8
  _LONDON SCIENCE CLASS-BOOKS_                                        32
  _LONGMANS' CIVIL ENGINEERING SERIES_                                13
  MACHINE DRAWING AND DESIGN                                          13
  MAGNETISM                                                           11
  MANUFACTURES                                                        17
  MECHANICS                                                            6
  MEDICINE AND SURGERY                                                19
  METALLURGY                                                          14
  MINERALOGY                                                          14
  NATURAL HISTORY                                                     18
  NAVIGATION                                                          14
  OPTICS                                                               8
  PHOTOGRAPHY                                                          8
  PHYSICS                                                              5
  PHYSIOGRAPHY                                                        17
  PHYSIOLOGY                                                          25
  _PRACTICAL ELEMENTARY SCIENCE SERIES_                               32
  _PROCTOR'S (R. A.) WORKS_                                           15
  SOUND                                                                8
  STATICS                                                              6
  STEAM, OIL, AND GAS ENGINES                                          9
  STRENGTH OF MATERIALS                                               12
  TECHNOLOGY                                                          17
  TELEGRAPHY                                                          12
  TELEPHONE                                                           12
  _TEXT-BOOKS OF SCIENCE_                                             29
  THERMODYNAMICS                                                       8
  _TYNDALL'S (JOHN) WORKS_                                            28
  VETERINARY MEDICINE, ETC.                                           24
  WORKSHOP APPLIANCES                                                 14
  ZOOLOGY                                                             25




                               CHEMISTRY.

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                              PHYSICS, ETC.


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            MECHANICS, DYNAMICS, STATICS, HYDROSTATICS, ETC.


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                         OPTICS AND PHOTOGRAPHY.


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                 SOUND, LIGHT, HEAT, AND THERMODYNAMICS.


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  _TYNDALL._--Works by JOHN TYNDALL, D.C.L., F.R.S. See p. 27.

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    ADVANCED HEAT. With 136 Diagrams and numerous Examples and
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                      STEAM, OIL, AND GAS ENGINES.


  _BALE._--A HANDBOOK FOR STEAM USERS; being Rules for Engine Drivers and
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  STEAM ENGINEERING. [In the press.

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  _STROMEYER._--MARINE BOILER MANAGEMENT AND CONSTRUCTION. Being a
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                         BUILDING CONSTRUCTION.


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      Examiner in Building Construction, Science and Art Department,
      South Kensington; with numerous Illustrations. Medium 8vo., 16_s._


              RIVINGTONS' COURSE OF BUILDING CONSTRUCTION.

  NOTES ON BUILDING CONSTRUCTION. Arranged to meet the requirements of
      the syllabus of the Science and Art Department of the Committee of
      Council on Education, South Kensington. Medium 8vo.

  Part I. First Stage, or Elementary Course. With 552 Woodcuts, 10_s._
  6_d._

  Part II. Commencement of Second Stage, or Advanced Course. With 479
  Woodcuts, 10_s._ 6_d._

  Part III. Materials. Advanced Course, and Course for Honours. With 188
  Woodcuts, 21_s._

  Part IV. Calculations for Building Structures. Course for Honours. With
  597 Woodcuts, 15_s._




                       ELECTRICITY AND MAGNETISM.


  _CUMMING._--ELECTRICITY TREATED EXPERIMENTALLY. For the Use of Schools
      and Students. By LINNÆUS CUMMING, M.A. With 242 Illustrations.
      Crown 8vo., 4_s._ 6_d._

  _DAY._--EXERCISES IN ELECTRICAL AND MAGNETIC MEASUREMENTS, with
      Answers. By R. E. DAY. 12mo., 3_s._ 6_d._

  _DU BOIS._--THE MAGNETIC CIRCUIT IN THEORY AND PRACTICE. By Dr. H.
      DU BOIS, Privatdocent in the University of Berlin. Translated by E.
      ATKINSON, Ph.D. With 94 Illustrations. 8vo., 12_s._ net.

  _EBERT._--MAGNETIC FIELDS OF FORCE: an Exposition of the Phenomena of
      Magnetism, Electro-Magnetism and Induction, based on the Conception
      of Lines of Force. By H. EBERT, Professor of Physics in the
      University of Kiel. Translated by C. V. BURTON, D.Sc. Part I. With
      93 Illustrations. 8vo., 10_s._ 6_d._ net.

  _GORE._--THE ART OF ELECTRO-METALLURGY, including all known Processes
      of Electro-Deposition. By G. GORE, LL.D., F.R.S. With 56 Woodcuts.
      Fcp. 8vo., 6_s._

  _JENKIN._--ELECTRICITY AND MAGNETISM. By FLEEMING JENKIN, F.R.S.S., L.
      and E., M.I.C.E. With 177 Illustrations. Fcp. 8vo., 3_s._ 6_d._

  _JOUBERT._--ELEMENTARY TREATISE ON ELECTRICITY AND MAGNETISM. Founded
      on JOUBERT'S 'Traité Élémentairé d'Electricité'. By G. C. FOSTER,
      F.R.S., and E. ATKINSON, Ph.D. With 381 Illustrations. Crown 8vo.,
      7_s._ 6_d._

  _JOYCE._--EXAMPLES IN ELECTRICAL ENGINEERING. By SAMUEL JOYCE, A.I.E.E.
      Crown 8vo., 5_s._

  _LARDEN._--ELECTRICITY FOR PUBLIC SCHOOLS AND COLLEGES. By W. LARDEN,
      M.A. With 215 Illustrations, and a Series of Examination Papers,
      with Answers. Crown 8vo., 6_s._

  _MERRIFIELD._--MAGNETISM AND DEVIATION OF THE COMPASS. For the Use of
      Students in Navigation and Science Schools. By JOHN MERRIFIELD,
      LL.D., F.R.A.S., 18mo., 2_s._ 6_d._

  _POYSER._--Works by A. W. POYSER, M.A., Grammar School, Wisbech.

    MAGNETISM AND ELECTRICITY. With 235 Illustrations. Crown 8vo., 2_s._
      6_d._

    ADVANCED ELECTRICITY AND MAGNETISM. With 317 Illustrations. Crown
      8vo., 4_s._ 6_d._

  _SLINGO AND BROOKER._--Works by W. SLINGO and A. BROOKER.

    ELECTRICAL ENGINEERING FOR ELECTRIC LIGHT ARTISANS AND STUDENTS. With
      346 Illustrations. Crown 8vo., 12_s._

    PROBLEMS AND SOLUTIONS IN ELEMENTARY ELECTRICITY AND MAGNETISM.
      Embracing a Complete Set of Answers to the South Kensington Papers
      for the years 1885-1894, and a Series of Original Questions. With
      67 Original Illustrations. Crown 8vo., 2_s._

  _TYNDALL._--Works by JOHN TYNDALL, D.C.L., F.R.S. See p. 27.




                      TELEGRAPHY AND THE TELEPHONE.


  _BENNETT._--THE TELEPHONE SYSTEMS OF CONTINENTAL EUROPE. By A. R.
      BENNETT, Member of the Institute of Electrical Engineers; late
      General Manager in Scotland of the National Telephone Company, and
      General Manager and Electrician of the Mutual and New Telephone
      Companies. With 169 Illustrations. Crown 8vo., 15_s._

  _CULLEY._--A HANDBOOK OF PRACTICAL TELEGRAPHY. By R. S. CULLEY,
      M.I.C.E., late Engineer-in-Chief of Telegraphs to the Post Office.
      With 135 Woodcuts and 17 Plates. 8vo., 16_s._

  _PREECE AND SIVEWRIGHT._--TELEGRAPHY. By W. H. PREECE, C.B., F.R.S.,
      V.P.Inst., C.E., etc., Engineer-in-Chief and Electrician Post
      Office Telegraphs; and Sir J. SIVEWRIGHT, K.C.M.G., General
      Manager, South African Telegraphs. With 258 Woodcuts. Fcp. 8vo.,
      6_s._




                ENGINEERING, STRENGTH OF MATERIALS, ETC.


  _ANDERSON._--THE STRENGTH OF MATERIALS AND STRUCTURES: the Strength of
      Materials as depending on their Quality and as ascertained by
      Testing Apparatus. By Sir J. ANDERSON, C.E., LL.D., F.R.S.E. With
      66 Woodcuts. Fcp. 8vo., 3_s._ 6_d._

  _BARRY._--RAILWAY APPLIANCES: a Description of Details of Railway
      Construction subsequent to the completion of the Earthworks and
      Structures. By Sir JOHN WOLFE BARRY, K.C.B., F.R.S., M.I.C.E. With
      218 Woodcuts. Fcp. 8vo., 4_s._ 6_d._

  _SMITH._--GRAPHICS, or the Art of Calculation by Drawing Lines, applied
      especially to Mechanical Engineering. By ROBERT H. SMITH, Professor
      of Engineering, Mason College, Birmingham. Part I. With separate
      Atlas of 29 Plates containing 97 Diagrams. 8vo., 15_s._

  _STONEY._--THE THEORY OF THE STRESSES ON GIRDERS AND SIMILAR STRUCTURES.
      With Practical Observations on the Strength and other Properties of
      Materials. By BINDON B. STONEY, LL.D., F.R.S., M.I.C.E. With 5
      Plates and 143 Illustrations in the Text. Royal 8vo., 36_s._

  _UNWIN._--Works by WILLIAM CAWTHORNE UNWIN, F.R.S., B.S.C.

    THE TESTING OF MATERIALS OF CONSTRUCTION. Embracing the description
      of Testing Machinery and Apparatus Auxiliary to Mechanical Testing,
      and an Account of the most Important Researches on the Strength of
      Materials. With 141 Woodcuts and 5 Folding-out Plates. 8vo., 21_s._

    ON THE DEVELOPMENT AND TRANSMISSION OF POWER FROM CENTRAL STATIONS:
      being the Howard Lectures delivered at the Society of Arts in 1893.
      With 81 Diagrams. 8vo., 10_s._ net.

  _WARREN._--ENGINEERING CONSTRUCTION IN IRON, STEEL, AND TIMBER. By
      WILLIAM HENRY WARREN, Challis Professor of Civil and Mechanical
      Engineering, University of Sydney. With 13 Folding Plates, and 375
      Diagrams. Royal 8vo., 16_s._ net.




                       MACHINE DRAWING AND DESIGN.


  _LOW AND BEVIS._--A MANUAL OF MACHINE DRAWING AND DESIGN. By DAVID
      ALLAN LOW (Whitworth Scholar), M.I.Mech.E., Professor of
      Engineering, East London Technical College, People's Palace,
      London; and ALFRED WILLIAM BEVIS (Whitworth Scholar), M.I.Mech.E.,
      Director of Manual Training to the Birmingham School Board. With
      700 Illustrations. 8vo., 7_s._ 6_d._

  _LOW._--Works by DAVID ALLAN LOW, Professor of Engineering, East London
      Technical College.

    IMPROVED DRAWING SCALES. 4_d._ in case.

    AN INTRODUCTION TO MACHINE DRAWING AND DESIGN. With 97 Illustrations
      and Diagrams. Crown 8vo., 2_s._

    MECHANICAL ENGINEER'S POCKET-BOOK. [_In the press._

  _UNWIN._--THE ELEMENTS OF MACHINE DESIGN. By W. CAWTHORNE UNWIN,
      F.R.S., Professor of Engineering at the Central Institute of the
      City and Guilds of London Institute.

  Part I. General Principles, Fastenings, and Transmissive Machinery.
  With 304 Diagrams, etc. Fcp. 8vo., 6_s._

  Part II. Chiefly on Engine Details. With 174 Woodcuts. Fcp. 8vo., 4_s._
  6_d._




                   LONGMANS' CIVIL ENGINEERING SERIES.

        Edited by the Author of 'Notes on Building Construction'.


  TIDAL RIVERS: their (1) Hydraulics, (2) Improvement, (3) Navigation.
      By W. H. WHEELER, M.Inst.C.E., author of 'The Drainage of Fens and
      Low Lands by Gravitation and Steam Power'. With 75 Illustrations.
      Medium 8vo., 16_s._ net.

  NOTES ON DOCKS AND DOCK CONSTRUCTION. By C. COLSON, M.Inst.C.E.,
      Assistant Director of Works, Admiralty. With 365 Illustrations.
      Medium 8vo., 21_s._ net.

  PRINCIPLES AND PRACTICE OF HARBOUR CONSTRUCTION. By WILLIAM SHIELD,
      F.R.S.E., M.Inst.C.E., and Executive Engineer, National Harbour of
      Refuge, Peterhead, N.B. With 97 Illustrations. Medium 8vo., 15_s._
      net.

  CALCULATIONS FOR ENGINEERING STRUCTURES. By T. CLAXTON FIDLER,
      M.I.C.E., Professor of Engineering in the University of Dundee;
      Author of 'A Practical Treatise on Bridge Construction'.
                                                       [_In preparation._

  PRINCIPLES AND PRACTICE OF CIVIL ENGINEERING. By L. F. VERNON-HARCOURT,
      M.Inst.C.E., Professor of Civil Engineering at University College,
      London.                                          [_In preparation._

  RAILWAY CONSTRUCTION. By W. H. MILLS, M.I.C.E., Engineer-in-Chief,
      Great Northern Railway, Ireland.                 [_In  preparation._




                        WORKSHOP APPLIANCES, ETC.


  _NORTHCOTT._--LATHES AND TURNING, Simple, Mechanical and Ornamental. By
      W. H. NORTHCOTT. With 338 Illustrations. 8vo., 18_s._

  _SHELLEY._--WORKSHOP APPLIANCES, including Descriptions of some of the
      Gauging and Measuring Instruments, Hand-cutting Tools, Lathes,
      Drilling, Plaining, and other Machine Tools used by Engineers. By
      C. P. B. SHELLEY, M.I.C.E. With an additional Chapter on Milling by
      R. R. LISTER. With 323 Woodcuts. Fcp. 8vo., 5_s._




                      MINERALOGY, METALLURGY, ETC.


  _BAUERMAN._--Works by HILARY BAUERMAN, F.G.S.

    SYSTEMATIC MINERALOGY. With 373 Woodcuts and Diagrams. Fcp. 8vo.,
      6_s._

    DESCRIPTIVE MINERALOGY. With 236 Woodcuts and Diagrams. Fcp. 8vo.,
      6_s._

  _GORE._--THE ART OF ELECTRO-METALLURGY, including all known Processes
      of Electro-Deposition. By G. GORE, LL.D., F.R.S. With 56 Woodcuts.
      Fcp. 8vo., 6_s._

  _HUNTINGTON AND M'MILLAN._--METALS: their Properties and Treatment. By
      A. K. HUNTINGTON, Professor of Metallurgy in King's College,
      London, and W. G. M'MILLAN, Lecturer on Metallurgy in Mason's
      College, Birmingham. With 122 Illustrations. Fcp. 8vo., 7_s._ 6_d._

  _RHEAD._--METALLURGY. An Elementary Text Book. By E. C. RHEAD, Lecturer
      on Metallurgy at the Municipal Technical School, Manchester. With
      94 Illustrations. Fcp. 8vo., 3_s._ 6_d._

  _RUTLEY._--THE STUDY OF ROCKS: an Elementary Text-book of Petrology. By
      F. RUTLEY, F.G.S. With 6 Plates and 88 Woodcuts. Fcp. 8vo., 4_s._
      6_d._




                       ASTRONOMY, NAVIGATION, ETC.


  _ABBOTT._--ELEMENTARY THEORY OF THE TIDES: the Fundamental Theorems
      Demonstrated without Mathematics and the Influence on the Length of
      the Day Discussed. By T. K. ABBOTT, B.D., Fellow and Tutor, Trinity
      College, Dublin. Crown 8vo., 2_s._

  _BALL._--Works by Sir ROBERT S. BALL, LL.D., F.R.S.

    ELEMENTS OF ASTRONOMY. With 130 Figures and Diagrams. Fcp. 8vo.,
      6_s._ 6_d._

    A CLASS-BOOK OF ASTRONOMY. With 41 Diagrams. Fcp. 8vo., 1_s._ 6_d._

  _CLERKE._--THE SYSTEM OF THE STARS. By AGNES M. CLERKE. With 6 Plates,
      and numerous Illustrations. 8vo., 21_s._

  _GOODWIN._--AZIMUTH TABLES FOR THE HIGHER DECLINATIONS. (Limits of
      Declination 24° to 30°, both inclusive.) Between the Parallels of
      Latitude 0° and 60°. With Examples of the Use of the Tables in
      English and French. By H. B. GOODWIN, Naval Instructor, Royal Navy.
      Royal 8vo., 7_s._ 6_d._

  _HERSCHEL._--OUTLINES OF ASTRONOMY.--By Sir JOHN F. W. HERSCHEL, Bart.,
      K. H., etc. With 9 Plates, and numerous Diagrams. 8vo., 12_s._

  _LOWELL._--MARS. By PERCIVAL LOWELL, Fellow American Academy, Member
      Royal Asiatic Society, Great Britain and Ireland, etc. With 24
      Plates. 8vo., 12_s._ 6_d._

  _MARTIN._--NAVIGATION AND NAUTICAL ASTRONOMY. Compiled by Staff
      Commander W. R. MARTIN, R.N. Royal 8vo., 18_s._

  _MERRIFIELD._--A TREATISE ON NAVIGATION. For the Use of Students. By
      J. MERRIFIELD, LL.D., F.R.A.S., F.M.S. With Charts and Diagrams.
      Crown 8vo., 5_s._

  _PARKER._--ELEMENTS OF ASTRONOMY. With Numerous Examples and
      Examination Papers. By GEORGE W. PARKER, M.A., of Trinity College,
      Dublin. With 84 Diagrams. 8vo., 5_s._ net.

  _WEBB._--CELESTIAL OBJECTS FOR COMMON TELESCOPES. By the Rev. _T. W.
      Webb_, M.A., F.R.A.S. Fifth Edition, Revised and greatly Enlarged
      by the Rev. T. E. ESPIN, M.A., F.R.A.S. (Two Volumes.) Vol. I.,
      with Portrait and a Reminiscence of the Author, 2 Plates, and
      numerous Illustrations. Crown 8vo., 6_s._ Vol. II., with numerous
      Illustrations. Crown 8vo., 6_s._ 6_d._




                      WORKS BY RICHARD A. PROCTOR.


  OLD AND NEW ASTRONOMY. With 21 Plates and 472 Illustrations in the
      Text. 4to., 21_s._

  THE MOON: Her Motions, Aspect, Scenery, and Physical Condition. With
      many Plates and Charts, Wood Engravings, and 2 Lunar Photographs.
      Crown 8vo., 3_s._ 6_d._

  THE UNIVERSE OF STARS: Researches into, and New Views respecting the
      Constitution of the Heavens. With 22 Charts (4 ), and 22
      Diagrams. 8vo., 10_s._ 6_d._

  OTHER WORLDS THAN OURS: the Plurality of Worlds Studied Under the Light
      of Recent Scientific Researches. With 14 Illustrations; Map,
      Charts, etc. Crown 8vo., 3_s._ 6_d._

  OUR PLACE AMONG INFINITIES: a Series of Essays contrasting our Little
      Abode in Space and Time with the Infinities around us. Crown 8vo.,
      3_s._ 6_d._

  MYTHS AND MARVELS OF ASTRONOMY. Crown 8vo., 3_s._ 6_d._

  LIGHT SCIENCE FOR LEISURE HOURS: Familiar Essays on Scientific
      Subjects. Natural Phenomena, etc. 3 vols., crown 8vo., 5_s._ each.

  THE ORBS AROUND US; Essays on the Moon and Planets, Meteors and Comets,
      the Sun and  Pairs of Suns. Crown 8vo., 3_s._ 6_d._

  THE EXPANSE OF HEAVEN: Essays on the Wonders of the Firmament. Crown
      8vo., 3_s._ 6_d._

  OTHER SUNS THAN OURS: a Series of Essays on Suns--Old, Young, and Dead.
      With other Science Gleanings. Two Essays on Whist, and
      Correspondence with Sir John Herschel. With 9 Star-Maps and
      Diagrams. Crown 8vo., 3_s._ 6_d._

  HALF-HOURS WITH THE TELESCOPE: a Popular Guide to the Use of the
      Telescope as a means of Amusement and Instruction. With 7 Plates.
      Fcp. 8vo., 2_s._ 6_d._

  NEW STAR ATLAS FOR THE LIBRARY, the School, and the Observatory, in
      Twelve Circular Maps (with Two Index-Plates). With an Introduction
      on the Study of the Stars. Illustrated by 9 Diagrams. Crown 8vo.,
      5_s._

  THE SOUTHERN SKIES: a Plain and Easy Guide to the Constellations of the
      Southern Hemisphere. Showing in 12 Maps the position of the
      principal Star-Groups night after night throughout the year. With
      an Introduction and a separate Explanation of each Map. True for
      every Year. 4to., 5_s._

  HALF-HOURS WITH THE STARS: a Plain and Easy Guide to the Knowledge of
      the Constellations. Showing in 12 Maps the position of the
      principal Star-Groups night after night throughout the year. With
      Introduction and a separate Explanation of each Map. True for every
      Year. 4to., 3_s._ 6_d._

  LARGER STAR ATLAS FOR OBSERVERS AND STUDENTS. In Twelve Circular Maps,
      showing 6000 Stars, 1500 Double Stars, Nebulæ, etc. With 2
      Index-Plates. Folio, 15_s._

  THE STARS IN THEIR SEASONS: an Easy Guide to a Knowledge of the
      Star-Groups. In 12 Large Maps. Imperial 8vo., 5_s._

  ROUGH WAYS MADE SMOOTH. Familiar Essays on Scientific Subjects. Crown
      8vo., 3_s._ 6_d._

  PLEASANT WAYS IN SCIENCE. Crown 8vo., 3_s._ 6_d._

  NATURE STUDIES. By R. A. PROCTOR, GRANT ALLEN, A. WILSON, T. FOSTER,
      and E. CLODD. Crown 8vo., 3_s._ 6_d._

  LEISURE READINGS. By R. A. PROCTOR, E. CLODD, A. WILSON, T. FOSTER, and
      A. C. RANYARD. Crown 8vo., 3_s._ 6_d._




                     MANUFACTURES, TECHNOLOGY, ETC.


  _BELL._--JACQUARD WEAVING AND DESIGNING. By F. T. BELL, Medallist in
      Honours and Certificated Teacher in 'Linen Manufacturing' and in
      'Weaving and Pattern Designing,' City and Guilds of London
      Institute. With 199 Diagrams. 8vo., 12_s._ net.

  _LUPTON._--MINING. An Elementary Treatise on the Getting of Minerals.
      By ARNOLD LUPTON, M.I.C.E., F.G.S., etc. With 596 Diagrams and
      Illustrations. Crown 8vo., 9_s._ net.

  _MORRIS AND WILKINSON._--THE ELEMENTS OF COTTON SPINNING. By JOHN
      MORRIS and F. WILKINSON. With a Preface by Sir B. A. DOBSON, C.E.,
      M.I.M.E. With 169 Diagrams and Illustrations. Crown 8vo, 7_s._
      6_d._ net.

  _SHARP._--BICYCLES AND TRICYCLES: an Elementary Treatise on their
      Design and Construction. With Examples and Tables. By ARCHIBALD
      SHARP, B.Sc., Whitworth Scholar; Associate Member of the
      Institution of Civil Engineers. With 565 Illustrations and
      Diagrams. Crown 8vo., 15_s._

  _TAYLOR._--COTTON WEAVING AND DESIGNING. By JOHN T. TAYLOR. With 373
      Diagrams. Crown 8vo., 7_s._ 6_d._ net.

  _WATTS._--AN INTRODUCTORY MANUAL FOR SUGAR GROWERS. By FRANCIS WATTS,
      F.C.S., F.I.C. With 20 Illustrations. Crown 8vo., 6_s._




                        PHYSIOGRAPHY AND GEOLOGY.


  _BIRD._--Works by CHARLES BIRD, B.A.

    ELEMENTARY GEOLOGY. With Geological Map of the British Isles, and 247
      Illustrations. Crown 8vo., 2_s._ 6_d._

  ADVANCED GEOLOGY. A Manual for Students in Advanced Classes and for
      General Readers. With over 300 Illustrations, a Geological Map of
      the British Isles (), and a set of Questions for
      Examination. Crown 8vo., 7_s._ 6_d._

  _THORNTON._--Works by J. THORNTON, M.A.

    ELEMENTARY PRACTICAL PHYSIOGRAPHY (for Section I. of the New Syllabus
      of the Science and Art Department). With 215 Illustrations. Crown
      8vo, 2_s._ 6_d._

    ELEMENTARY PHYSIOGRAPHY: an Introduction to the Study of Nature. With
      12 Maps and 247 Illustrations. With Appendix on Astronomical
      Instruments and Measurements. Crown 8vo., 2_s._ 6_d._

    ADVANCED PHYSIOGRAPHY. With 6 Maps and 203 Illustrations. Crown 8vo.,
      4_s._ 6_d._




                           HEALTH AND HYGIENE.


  _BRODRIBB._--MANUAL OF HEALTH AND TEMPERANCE. By T. BRODRIBB, M.A. With
      Extracts from Gough's 'Temperance Orations'. Revised and Edited by
      the Rev. W. RUTHVEN PYM, M.A., Member of the Sheffield School
      Board. Crown 8vo., 1_s._ 6_d._

  _BUCKTON._--HEALTH IN THE HOUSE; Twenty-five Lectures on Elementary
      Physiology. By Mrs. C. M. BUCKTON. With 41 Woodcuts and Diagrams.
      Crown 8vo., 2_s._

  _CORFIELD._--THE LAWS OF HEALTH. By W. H. CORFIELD, M.A.. M.D. Fcp.
      8vo., 1_s._ 6_d._

  _NOTTER AND FIRTH._--Works by J. L. NOTTER, M.A., M.D., and R. H.
      FIRTH, F.R.C.S.

    HYGIENE. With 95 Illustrations. Crown 8vo., 3_s._ 6_d._

    PRACTICAL DOMESTIC HYGIENE. With 83 Illustrations. Crown 8vo., 2_s._
      6_d._

  _POORE._--Works by GEORGE VIVIAN POORE, M.D.

    ESSAYS ON RURAL HYGIENE. Crown 8vo., 6_s._ 6_d._

    THE DWELLING-HOUSE. With 36 Illustrations. Crown 8vo., 3_s._ 6_d._

  _WILSON._--A MANUAL OF HEALTH-SCIENCE: adapted for use in Schools and
      Colleges. By ANDREW WILSON, F.R.S.E., F.L.S., etc. With 74
      Illustrations. Crown 8vo., 2_s._ 6_d._




                            NATURAL HISTORY.


  _FURNEAUX._--Works by WILLIAM S. FURNEAUX, F.R.G.S.

    THE OUTDOOR WORLD; or, The Young Collector's Handbook. With 18
      Plates, 16 of which are , and 549 Illustrations in the
      Text. Crown 8vo., 7_s._ 6_d._

    LIFE IN PONDS AND STREAMS. With 8  Plates and 331
      Illustrations in the Text. Crown 8vo., 7_s._ 6_d._

    BUTTERFLIES AND MOTHS (British). With 12  Plates and 241
      Illustrations in the Text. 7_s._ 6_d._

  _HUDSON._--BRITISH BIRDS. By W. H. HUDSON, C.M.Z.S. With 8 
      Plates from Original Drawings by A. THORBURN, and 8 Plates and 100
      Figures by C. E. LODGE, and 3 Illustrations from Photographs. Crown
      8vo., 7_s._ 6_d._

  _STANLEY._--A FAMILIAR HISTORY OF BIRDS. By E. STANLEY, D.D., formerly
      Bishop of Norwich. With 160 Illustrations. Crown 8vo, 3_s._ 6_d._




                          MEDICINE AND SURGERY.


  _ASHBY._--NOTES ON PHYSIOLOGY FOR THE USE OF STUDENTS PREPARING FOR
      EXAMINATION. By HENRY ASHBY, M.D. Lond., F.R.C.P., Physician to the
      General Hospital for Sick Children, Manchester; formerly
      Demonstrator of Physiology, Liverpool School of Medicine. Sixth
      Edition, thoroughly Revised. With 141 Illustrations. Fcp. 8vo.,
      5_s._

  _ASHBY AND WRIGHT._--THE DISEASES OF CHILDREN, MEDICAL AND SURGICAL. By
      HENRY ASHBY, M.D., Lond., F.R.C.P., Physician to the General
      Hospital for Sick Children, Manchester; and G. A. WRIGHT, B.A.,
      M.B. Oxon., F.R.C.S., Eng., Assistant Surgeon to the Manchester
      Royal Infirmary, and Surgeon to the Children's Hospital. Enlarged
      and Improved Edition. With 192 Illustrations. 8vo., 25_s._

  _BENNETT._--Works by WILLIAM H. BENNETT, F.R.C.S., Surgeon to St.
      George's Hospital; Member of the Board of Examiners, Royal College
      of Surgeons of England.

    CLINICAL LECTURES ON VARICOSE VEINS OF THE LOWER EXTREMITIES. With 3
      Plates. 8vo., 6_s._

    ON VARICOCELE; A PRACTICAL TREATISE. With 4 Tables and a Diagram.
      8vo., 5_s._

    CLINICAL LECTURES ON ABDOMINAL HERNIA: chiefly in relation to
      Treatment, including the Radical Cure. With 12 Diagrams in the
      Text. 8vo., 8_s._ 6_d._

  _BENTLEY._--A TEXT-BOOK OF ORGANIC MATERIA MEDICA. Comprising a
      Description of the Vegetable and Animal Drugs of the British
      Pharmacop[oe]ia, with some others in common use. Arranged
      Systematically, and Especially Designed for Students. By ROBERT
      BENTLEY, M.R.C.S. Eng., F.L.S. With 62 Illustrations on Wood. Crown
      8vo., 7_s._ 6_d._

  _BRODIE._--THE ESSENTIALS OF EXPERIMENTAL PHYSIOLOGY. For the Use of
      Students. By T. G. BRODIE, M.D., Lecturer on Physiology, St.
      Thomas's Hospital Medical School.                  [_In the press._

  _CABOT._--A GUIDE TO THE CLINICAL EXAMINATION OF THE BLOOD FOR
      DIAGNOSTIC PURPOSES. By RICHARD C. CABOT, M.D. With 3 
      Plates and 28 Illustrations in the Text. 8vo., 16_s._

  _CLARKE._--Works by J. JACKSON CLARKE, M.D. Lond., F.R.C.S., Assistant
      Surgeon at the North-west London and City Orthopædic Hospitals,
      late Senior Demonstrator of Anatomy, Demonstrator of Bacteriology,
      and Curator of the Museum in St. Mary's Hospital Medical School,
      and Pathologist to St. Mary's Hospital.

    SURGICAL PATHOLOGY AND PRINCIPLES. With 194 Illustrations. Crown
      8vo., 10_s._ 6_d._

    POST-MORTEM EXAMINATIONS IN MEDICO-LEGAL AND ORDINARY CASES. With
      Special Chapters on the Legal Aspects of Post Mortems, and on
      Certificates of Death. Fcp. 8vo., 2_s._ 6_d._

  _COATS._--A MANUAL OF PATHOLOGY. By JOSEPH COATS, M.D., Professor of
      Pathology in the University of Glasgow. Third Edition. Revised
      throughout. With 507 Illustrations. 8vo., 31_s._ 6_d._

  _COOKE._--Works by THOMAS COOKE, F.R.C.S. Eng., B.A., B.Sc., M.D.,
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    APHORISMS IN APPLIED ANATOMY AND OPERATIVE SURGERY. Including 100
      Typical _vivâ voce_ Questions on Surface Marking, etc. Crown 8vo.,
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    DISSECTION GUIDES. Aiming at Extending and Facilitating such
      Practical work in Anatomy as will be specially useful in connection
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  _DAKIN._--A HANDBOOK OF MIDWIFERY. By WILLIAM RADFORD DAKIN, M.D.,
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      George's Hospital, etc. With 394 Illustrations. Large crown 8vo.,
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  _DICKINSON._--Works by W. HOWSHIP DICKINSON, M.D. Cantab., F.R.C.P.,
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    ON RENAL AND URINARY AFFECTIONS. With 12 Plates and 122 Woodcuts.
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    THE TONGUE AS AN INDICATION OF DISEASE; being the Lumleian Lectures
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      7_s._ 6_d._

    OCCASIONAL PAPERS ON MEDICAL SUBJECTS, 1855-1896. 8vo., 12_s._

  _DUCKWORTH._--THE SEQUELS OF DISEASE: being the Lumleian Lectures
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      Observations on Prognosis in Disease. By Sir DYCE DUCKWORTH, M.D.,
      LL.D., Fellow and Treasurer of the Royal College of Physicians,
      etc. 8vo., 10_s._ 6_d._

  _ERICHSEN._--THE SCIENCE AND ART OF SURGERY; a Treatise on Surgical
      Injuries, Diseases, and Operations. By Sir JOHN ERIC ERICHSEN,
      Bart., F.R.S., LL.D. Edin., Hon. M.Ch. and F.R.C.S. Ireland, late
      Surgeon Extraordinary to H.M. the Queen. Illustrated by nearly 1000
      Engravings on Wood. 2 vols. Royal 8vo., 48_s._

  _FOWLER AND GOODLEE._--THE DISEASES AND INJURIES OF THE LUNGS AND
      PLEURA. By JAMES KINGSTON FOWLER, M.A., M.D., F.R.C.B., Physician
      and Lecturer on Pathological Anatomy, Middlesex Hospital, etc.; and
      RICKMAN J. GOODLEE, B.A. Lond., M.B., F.R.C.S., etc. With
      Illustrations.                                     [_In the press._

  _GARROD._--Works by SIR ALFRED BARING GARROD, M.D., F.R.S., etc.,
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    A TREATISE ON GOUT AND RHEUMATIC GOUT (RHEUMATOID ARTHRITIS). Third
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      8vo., 21_s._

    THE ESSENTIALS OF MATERIA MEDICA AND THERAPEUTICS. The Thirteenth
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      and Therapeutics in King's College, London, etc. Crown 8vo., 12_s._
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  _GRAY._--ANATOMY, DESCRIPTIVE AND SURGICAL. By HENRY GRAY, F.R.S., late
      Lecturer on Anatomy at St. George's Hospital. The Fourteenth
      Edition, re-edited by T. PICKERING PICK, Surgeon to St. George's
      Hospital, Inspector of Anatomy in England and Wales, late Member of
      the Court of Examiners, Royal College of Surgeons of England. With
      705 large Woodcut Illustrations, a large proportion of which are
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      Nerves yellow. The attachments of the muscles to the bones, in the
      section on Osteology, are also shown in  outline. Royal
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  _HALFORD._--THE LIFE OF SIR HENRY HALFORD, Bart., G.C.H., M.D., F.R.S.,
      President of the Royal College of Physicians, Physician to George
      III., George IV., William IV., and to Her Majesty Queen Victoria.
      By WILLIAM MUNK, M.D., F.S.A., Fellow and late Vice-President of
      the Royal College of Physicans of London. With 2 Portraits. 8vo.,
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  _HALLIBURTON._--Works by W. D. HALLIBURTON, M.D., F.R.S., M.R.C.P.,
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    A TEXT-BOOK OF CHEMICAL PHYSIOLOGY AND PATHOLOGY. With 104
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    ESSENTIALS OF CHEMICAL PHYSIOLOGY. 8vo., 5_s._

      *** This is a book suitable for medical students. It treats of the
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      be performed by each member of the class.

  _LANG._--THE METHODICAL EXAMINATION OF THE EYE. Being Part I. of a
      Guide to the Practice of Ophthalmology for Students and
      Practitioners. By WILLIAM LANG, F.R.C.S. Eng., Surgeon to the Royal
      London Ophthalmic Hospital, Moorfields, etc. With 15 Illustrations.
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  _LIVEING._--HANDBOOK ON DISEASES OF THE SKIN. With especial reference
      to Diagnosis and Treatment. By ROBERT LIVEING, M.A. and M.D.,
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  _LONGMORE._--Works by Surgeon-General Sir T. LONGMORE C.B., F.R.C.S.,
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    THE ILLUSTRATED OPTICAL MANUAL; OR, HANDBOOK OF INSTRUCTIONS FOR THE
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    GUNSHOT INJURIES. Their History, Characteristic Features,
      Complications, and General Treatment; with Statistics concerning
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  _LUFF._--TEXT-BOOK OF FORENSIC MEDICINE AND TOXICOLOGY. By ARTHUR P.
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  NEWMAN.--ON THE DISEASES OF THE KIDNEY AMENABLE TO SURGICAL TREATMENT.
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  OWEN.--A MANUAL OF ANATOMY FOR SENIOR STUDENTS. By EDMUND OWEN, M.B.,
      F.R.S.C., Senior Surgeon to the Hospital for Sick Children, Great
      Ormond Street, Surgeon to St. Mary's Hospital, London, and
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  POOLE.--COOKERY FOR THE DIABETIC. By W. H. and Mrs. POOLE. With Preface
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  QUAIN.--A DICTIONARY OF MEDICINE; Including General Pathology, General
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  _QUAIN._--QUAIN'S (JONES) ELEMENTS OF ANATOMY. The Tenth Edition.
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      and Histology in University College, London; and GEORGE DANCER
      THANE, Professor of Anatomy in University College, London.

      *** The several parts of this work form COMPLETE TEXT-BOOKS OF
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      VOL. I., PART I. EMBRYOLOGY. By E. A. SCHÄFER, F.R.S. With 200
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      VOL. II., PART II. ARTHROLOGY--MYOLOGY--ANGEIOLOGY. By G. D. THANE.
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      VOL. III., PART I. THE SPINAL CORD AND BRAIN. By E. A. SCHÄFER,
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      VOL. III. PART II. THE NERVES. By G. D. THANE. With 102
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      VOL. III., PART III. THE ORGANS OF THE SENSES. By E. A. SCHÄFER,
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      VOL. III., PART IV. SPLANCHNOLOGY. By E. A. SCHÄFER, F.R.S., and
      JOHNSON SYMINGTON, M.D. With 337 Illustrations. Royal 8vo., 16_s._

      APPENDIX. SUPERFICIAL AND SURGICAL ANATOMY. By Professor G. D.
      THANE and Professor R. J. GODLEE, M.S. With 29 Illustrations. Royal
      8vo., 6_s._ 6_d._

  _RICHARDSON._--VITA MEDICA: Chapters of Medical Life and Work. By Sir
      B. W. RICHARDSON, M.A., LL.D., F.R.S. 8vo., 16_s._

  _SCHÄFER._--THE ESSENTIALS OF HISTOLOGY. Descriptive and Practical. For
      the Use of Students. By E. A. SCHÄFER, F.R.S., Jodrell Professor of
      Physiology in University College, London; Editor of the
      Histological Portion of Quain's 'Anatomy'. Illustrated by more than
      300 Figures, many of which are new. Fourth Edition, Revised and
      Enlarged. 8vo., 7_s._ 6_d._ (Interleaved, 10_s._)

  _SCHENK._--MANUAL OF BACTERIOLOGY. For Practitioners and Students. With
      especial reference to Practical Methods. By Dr. S. L. SCHENK,
      Professor (Extraordinary) in the University of Vienna. Translated
      from the German, with an Appendix, by W. R. DAWSON, B.A., M.D.,
      Univ. Dub.; late University Travelling Prizeman in Medicine. With
      100 Illustrations, some of which are . 8vo., 10_s._ net.

  _SMALE AND COLYER._ DISEASES AND INJURIES OF THE TEETH, including
      Pathology and Treatment: a Manual of Practical Dentistry for
      Students and Practitioners. By MORTON SMALE, M.R.C.S., L.S.A.,
      L.D.S., Dental Surgeon to St. Mary's Hospital, Dean of the School,
      Dental Hospital of London, etc.; and J. F. COLYER, L.R.C.P.,
      M.R.C.S., L. D.S., Assistant Dental Surgeon to Charing Cross
      Hospital, and Assistant Dental Surgeon to the Dental Hospital of
      London. With 334 Illustrations. Large Crown 8vo., 15_s._

  _SMITH_ (_H. F._). THE HANDBOOK FOR MIDWIVES. By HENRY FLY SMITH, B.A.,
      M.B. Oxon., M.R.C.S. Second Edition. With 41 Woodcuts. Crown 8vo.,
      5_s._

  _STEVENSON._--WOUNDS IN WAR: the Mechanism of their Production and
      their Treatment. By Surgeon-Colonel W. F. STEVENSON (Army Medical
      Staff), A.B., M.B., M.Ch. Dublin University, Professor of Military
      Surgery, Army Medical School, Netley. With 86 Illustrations. 8vo.,
      18_s._

  _TIRARD._--DIPHTHERIA AND ANTITOXIN. By NESTOR TIRARD, M.D. Lond.,
      Fellow of the Royal College of Physicians; Fellow of King's
      College, London; Professor of Materia Medica and Therapeutics at
      King's College; Physician to King's College Hospital; and Senior
      Physician to the Evelina Hospital for Sick Children. 8vo., 7_s._
      6_d._

  _WAKLEY._--THE LIFE AND TIMES OF THOMAS WAKLEY, Founder and First
      Editor of the _Lancet_, Member of Parliament for Finsbury, and
      Coroner for West Middlesex. By SQUIRE SPRIGGE, M.B. Cantab. With 2
      Portraits. 8vo., 18_s._

  _WALLER._--Works by AUGUSTUS D. WALLER, M.D., Lecturer on Physiology at
      St. Mary's Hospital Medical School, London; late External Examiner
      at the Victorian University.

    AN INTRODUCTION TO HUMAN PHYSIOLOGY. Third Edition, Revised. With 314
      Illustrations. 8vo., 18_s._

    LECTURES ON PHYSIOLOGY. First Series. On Animal Electricity. 8vo.,
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    EXERCISES IN PRACTICAL PHYSIOLOGY. Part I. Elementary Physiological
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      net. Part II. in the press. Part III. Physiology of the Nervous
      System; Electro-Physiology. 8vo., 2_s._ 6_d._ net.

  _WEICHSELBAUM._--THE ELEMENTS OF PATHOLOGICAL HISTOLOGY. With Special
      Reference to Practical Methods. By Dr. ANTON WEICHSELBAUM,
      Professor of Pathology in the University of Vienna. Translated by
      W. R. DAWSON, M.D. (Dub.), Demonstrator of Pathology in the Royal
      College of Surgeons, Ireland, late Medical Travelling Prizeman of
      Dublin University, etc. With 221 Figures, partly in Colours, a
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  _WILKS AND MOXON._--LECTURES ON PATHOLOGICAL ANATOMY. By Sir SAMUEL
      WILKS, Bart., M.D., F.R.S., President of the Royal College of
      Physicians, and Physician Extraordinary to H.M. the Queen, and the
      late WALTER MOXON, M.D., F.R.C.P., Physician to, and some time
      Lecturer on Pathology at, Guy's Hospital. Third Edition, thoroughly
      Revised. By Sir SAMUEL WILKS, Bart., M.D., LL.D., F.R.S. 8vo.,
      18_s._




                        VETERINARY MEDICINE, ETC.


  _STEEL._--Works by JOHN HENRY STEEL, F.R.C.V.S., F.Z.S., A.V.D., late
      Professor of Veterinary Science and Principal of Bombay Veterinary
      College.

    A TREATISE ON THE DISEASES OF THE DOG; being a Manual of Canine
      Pathology. Especially adapted for the use of Veterinary
      Practitioners and Students. With 88 Illustrations. 8vo., 10_s._
      6_d._

    A TREATISE ON THE DISEASES OF THE OX; being a Manual of Bovine
      Pathology. Especially adapted for the use of Veterinary
      Practitioners and Students. With 2 Plates and 117 Woodcuts. 8vo.,
      15_s._

    A TREATISE ON THE DISEASES OF THE SHEEP; being a Manual of Ovine
      Pathology for the use of Veterinary Practitioners and Students.
      With  Plate and 99 Woodcuts. 8vo., 12_s._

    OUTLINES OF EQUINE ANATOMY; a Manual for the use of Veterinary
      Students in the Dissecting Room. Crown 8vo., 7_s._ 6_d._

  _FITZWYGRAM._--HORSES AND STABLES. By Major-General Sir F. FITZWYGRAM,
      Bart. With 56 pages of Illustrations. 8vo., 2_s._ 6_d._ net.

  _SCHREINER._--THE ANGORA GOAT (published under the auspices of the
      South African Angora Goat Breeders' Association), and a Paper on
      the Ostrich (reprinted from the _Zoologist_ for March, 1897). By S.
      C. CRONWRIGHT SCHREINER. 8vo.

  '_STONEHENGE._'--THE DOG IN HEALTH AND DISEASE. By 'STONEHENGE'. With
      78 Wood Engravings. 8vo., 7_s._ 6_d._

  _YOUATT._--Works by WILLIAM YOUATT.

    THE HORSE. Revised and Enlarged by W. WATSON, M.R.C.V.S. With 52 Wood
      Engravings. 8vo., 7_s._ 6_d._

    THE DOG. Revised and Enlarged. With 33 Wood Engravings. 8vo., 6_s._




             PHYSIOLOGY, BIOLOGY, BACTERIOLOGY, AND ZOOLOGY.

                    (And see _MEDICINE AND SURGERY_.)


  _ASHBY._--NOTES ON PHYSIOLOGY, for the Use of Students Preparing for
      Examination. By HENRY ASHBY, M.D. With 141 Illustrations. Fcp.
      8vo., 5_s._

  _BARNETT._--THE MAKING OF THE BODY: a Children's Book on Anatomy and
      Physiology, for School and Home Use. By Mrs. S. A. BARNETT. With
      113 Illustrations. Crown 8vo., 1_s._ 9_d._

  _BIDGOOD._--A COURSE OF PRACTICAL ELEMENTARY BIOLOGY. By JOHN BIDGOOD,
      B.Sc., F.L.S. With 226 Illustrations. Crown 8vo., 4_s._ 6_d._

  _BRAY._--PHYSIOLOGY AND THE LAWS OF HEALTH, in Easy Lessons for
      Schools. By Mrs. CHARLES BRAY. Fcp. 8vo., 1_s._

  _CURTIS._--ESSENTIALS OF PRACTICAL BACTERIOLOGY. By HENRY J. CURTIS,
      M.D. With numerous Illustrations.                  [_In the press._

  _FRANKLAND._--MICRO-ORGANISMS IN WATER. Together with an Account of the
      Bacteriological Methods involved in their Investigation. Specially
      designed for the use of those connected with the Sanitary Aspects
      of Water-Supply. By PERCY FRANKLAND, Ph.D., B.Sc. (Lond.), F.R.S.,
      and Mrs. PERCY FRANKLAND. With 2 Plates and Numerous Diagrams.
      8vo., 16_s._ net.

  _FURNEAUX._--HUMAN PHYSIOLOGY. By W. FURNEAUX, F.R.G.S. With 218
      Illustrations. Crown 8vo., 2_s._ 6_d._

  _HUDSON AND GOSSE._--THE ROTIFERA, or 'WHEEL-ANIMACULES'. By C. T.
      HUDSON, LL.D., and P. H. GOSSE, F.R.S. With 30  and 4
      Uncoloured Plates. In 6 Parts. 4to., 10_s._ 6_d._ each Supplement
      12_s._ 6_d._ Complete in 2 vols., with Supplement, 4to., £4 4_s._

  _LEUMANN._--NOTES ON MICRO-ORGANISMS PATHOGENIC TO MAN. By
      Surgeon-Captain B. H. S. LEUMANN, M.B., Indian Medical Service.
      8vo., 3_s._

  _MACALISTER._--Works by ALEXANDER MACALISTER, M.D., Professor of
      Comparative Anatomy and Zoology, University of Dublin. With 41
      Diagrams.

    AN INTRODUCTION TO THE SYSTEMATIC ZOOLOGY AND MORPHOLOGY OF
      VERTEBRATE ANIMALS. 8vo., 10_s._ 6_d._

    ZOOLOGY OF THE INVERTEBRATE ANIMALS. With 59 Diagrams. Fcp. 8vo.,
      1_s._ 6_d._

    ZOOLOGY OF THE VERTEBRATE ANIMALS. With 77 Diagrams. Fcp. 8vo.,
      1_s._ 6_d._

  _MORGAN._--ANIMAL BIOLOGY: an Elementary Text-Book. By C. LLOYD MORGAN.
      With 103 Illustrations. Crown 8vo., 8_s._ 6_d._

  _SCHENK._--MANUAL OF BACTERIOLOGY, for Practitioners and Students, with
      Especial Reference to Practical Methods. By Dr. S. L. SCHENK. With
      100 Illustrations, some . 8vo., 10_s._ net.

  _THORNTON._--HUMAN PHYSIOLOGY. By JOHN THORNTON, M.A. With 267
      Illustrations, some . Crown 8vo., 6_s._




                                 BOTANY.


  _AITKEN._--ELEMENTARY TEXT-BOOK OF BOTANY. By EDITH AITKEN, late
      Scholar of Girton College. With 400 Diagrams. Crown 8vo., 4_s._
      6_d._

  _BENNETT AND MURRAY._--HANDBOOK OF CRYPTOGAMIC BOTANY. By ALFRED W.
      BENNETT, M.A., B.Sc., F.L.S., Lecturer on Botany at St. Thomas's
      Hospital; and GEORGE MURRAY, F.L.S., Keeper of Botany, British
      Museum. With 378 Illustrations. 8vo., 16_s._

  _CROSS AND BEVAN._--CELLULOSE: an Outline of the Chemistry of the
      Structural Elements of Plants. With Reference to their Natural
      History and Industrial Uses. By CROSS and BEVAN (C. F. Cross, E. J.
      Bevan, and C. Beadle). With 14 Plates. Crown 8vo., 12_s._ net.

  _CURTIS._--A TEXT-BOOK OF GENERAL BOTANY. By CARLTON C. CURTIS, A.M.,
      Ph.D., Tutor in Botany in Columbia University, U.S.A. With 87
      Illustrations. 8vo., 12_s._ net.

  _EDMONDS._--Works by HENRY EDMONDS, B.Sc., London.

    ELEMENTARY BOTANY, Theoretical and Practical. With 319 Illustrations.
      Crown 8vo., 2_s._ 6_d._

    BOTANY FOR BEGINNERS. With 85 Illustrations. Fcp. 8vo., 1_s._ 6_d._

  _KITCHENER._--A YEAR'S BOTANY. Adapted to Home and School Use. By
      FRANCES A. KITCHENER. With 195 Illustrations. Crown 8vo., 5_s._

  _LINDLEY AND MOORE._--THE TREASURY OF BOTANY. Edited by J. LINDLEY,
      M.D., F.R.S., and T. MOORE, F.L.S. With 20 Steel Plates and
      numerous Woodcuts. Two parts. Fcp. 8vo., 12_s._

  _McNAB._--CLASS-BOOK OF BOTANY. By W. R. MCNAB. Two Parts.

    MORPHOLOGY AND PHYSIOLOGY. With 42 Diagrams. Fcp. 8vo., 1_s._ 6_d._

    CLASSIFICATION OF PLANTS. With 118 Diagrams. Fcp. 8vo., 1_s._ 6_d._

  _SORAUER._--A POPULAR TREATISE ON THE PHYSIOLOGY OF PLANTS. For the
      use of Gardeners, or for Students of Horticulture and of
      Agriculture. By Dr. PAUL SORAUER. Translated by F. E. WEISS, B.Sc.,
      F.L.S. With 33 Illustrations. 8vo., 9_s._ net.

  _THOMÉ AND BENNETT._--STRUCTURAL AND PHYSIOLOGICAL BOTANY. By OTTO
      WILHELM THOMÉ and by ALFRED W. BENNETT, M.A., B.Sc., F.L.S. With
       Map and 600 Woodcuts. Fcp. 8vo., 6_s._

  _TUBEUF._--DISEASES OF PLANTS INDUCED BY CRYPTOGAMIC PARASITES.
      Introduction to the Study of Pathogenic Fungi, Slime Fungi,
      Bacteria and Algæ. By Dr. KARL FREIHERR VON TUBEUF, Privatdocent in
      the University of Munich. English Edition by WILLIAM G. SMITH,
      B.Sc., Ph.D., Lecturer on Plant Physiology, University of
      Edinburgh. With 330 Illustrations. Royal 8vo., 18_s._ net.

  _WATTS._--A SCHOOL FLORA. For the use of Elementary Botanical Classes.
      By W. MARSHALL WATTS, D.Sc. Lond. Crown 8vo., 2_s._ 6_d._




                              AGRICULTURE.


  _ADDYMAN._--AGRICULTURAL ANALYSIS. A Manual of Quantitative Analysis
      for Students of Agriculture. By FRANK T. ADDYMAN, B.Sc. (Lond.),
      F.I.C. With 49 Illustrations. Crown 8vo., 5_s._ net.

  _COLEMAN AND ADDYMAN._--PRACTICAL AGRICULTURAL CHEMISTRY. By J. BERNARD
      COLEMAN, A.R.C.Sc., F.I.C., and FRANK T. ADDYMAN, B.Sc. (Lond.),
      F.I.C. With 24 Illustrations. Crown 8vo., 1_s._ 6_d._ net.

  _COOKE._--THE FOUNDATIONS OF SCIENTIFIC AGRICULTURE. By SAMUEL COOKE,
      A.M., A.M.I.C.E., F.I.C., F.G.S., Cor.M.R.H.S., Principal and
      Professor of Chemistry and Geology, College of Science, Poona. With
      85 Illustrations and a Series of Examination Questions. Crown 8vo.,
      4_s._ 6_d._

  _WEBB._--Works by HENRY J. WEBB, Ph.D., B.Sc. (Lond.); late Principal of
      the Agricultural College, Aspatria.

    ELEMENTARY AGRICULTURE. A Text-Book specially adapted to the
      requirements of the Science and Art Department, the Junior
      Examination of the Royal Agricultural Society, and other Elementary
      Examinations. With 34 Illustrations. Crown 8vo., 2_s._ 6_d._

    AGRICULTURE. A Manual for Advanced Science Students. With 100
      Illustrations. Crown 8vo., 7_s._ 6_d._ net.




              WORKS BY JOHN TYNDALL, D.C.L., LL.D., F.R.S.


  FRAGMENTS OF SCIENCE: a Series of Detached Essays, Addresses, and
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                    *       *       *       *       *




Transcriber's note:

The half-title has been deleted.

Obvious misprints were corrected.

Inconsistant hyphenation was kept as it is.



***