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DR. I.E. SMITH
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(A LIBRARY OF
UNIVERSAL LITERATURE
IN F O U R PAR T S
f Comprising Science, Biography, Fiction ]
V and the Great Orations
PART ONE— SCIENCE
The Principles of Chemistry
(PART TWO)
BY
D. MENDELEEFF
520957
t .4 si
NEW YORK
P. F. COLLIER AND SON
• M C M I •
26
BOARD OF EDITORS
SCIENCE
ANGBLO HEILPRIN, author of "The Earth and Its Story," etc.;
Curator Academy of Natural Sciences of Philadelphia.
JOSEPH TORREY, JR., Ph.D., Instructor in Chemistry in Harvard
University.
RAY STANNARD BAKER, A.B., author of "The New Prosperity,"
etc.; Associate Editor of McClure's Magazine.
BIOGRAPHY
MAYO W. HAZELTINE, A.M., author of "Chats About Books," etc.;
Literary Editor of the New York Sun.
JULIAN HAWTHORNE, author of "Nathaniel Hawthorne and His
Wife," "History of the United States," etc.
CHARLBS G. D. ROBERTS, A.B., A.M., author of "A History of
Canada"; late Professor of English and French Literature,
King's College.
FICTION
RICHARD HENRY STODDARD, author of "The King's Bell," etc.;
Literary Editor of the New York Mail and Express.
HENRY VAN DYKE, D.D., LL.D., author of "Little Rivers," etc.;
Professor of English Literature at Princeton University.
THOMAS NELSON PAGE, LL.D., Litt.D., author of "Red Rock," etc.
ORATIONS
HON. HENRY CABOT LODGE, A.B., LL.B., author of "Life of Daniel
Webster," etc.; U. S. Senator from Massachusetts.
HON. JOHN R. PROCTOR, President U. S. Civil Service Commission.
MORRIS HICKEY MORGAN, Ph.D., LL.D., Professor in Latin, Har-
vard University.
PEINCIPLES OF CHEMISTBY
(PART TWO)
PEINCIPL.ES OF CHEMISTBY
CHAPTER VII
MOLECULES AND ATOMS. THE LAWS OP GAY-LUSSAC AND
AVOGADRO-GEBHARDT
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 decom-
posed 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, «kc., 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,
1 If the weight be indicated by P, the density by D, and the volume by V, then
D = *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
299
800 PRINCIPLES, OF CHEMISTRY"
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 ^, and therefore
it is seen without direct experiment that water contains two volumes
of hydrogen for every one volume of oxygen. So also, knowing {hat
nitric oxide contains fourteen parts of nitrogen and sixteen parts of
oxygen, and knowing that the specific jjravities 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 NOa only becomes constant and equal to twenty?
three (referred to hydrogen) above 1 35°, 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. N02, 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.,, 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
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),Jf 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 on« part by weight of hydrogen, is taken
as unity, and all volumes are expressed according to these nnita.
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 deter*
mining 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 weighty- 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 28) 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 easjly
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
MOLECULES ANJD ATOMS 801
reaction or entering into the composition of a compound, may be also
determined.
•odium 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 varions pressures are given in Chapter II., Note 27. (2) With respect to tempera-
tures above 800° (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 calcu-
lated temperature i) by means of substances boiling at a high temperature. Thus, for
i&stance, at the ordinary atmospheric pressure the temperature t of the vapour of
sulphur is about 445°, of phosphorus pentasnlphide 518°, of tin chloride 606°, of cad-
mium 770°, of zinc 930° (according to Violle and others), or 1040° (according to Deville),
&c. (S) 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 pres-
sure which is read on the barometer ; but in the case of substances which are volati-
lised 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 deter-
mination 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 decom-
posable, 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 phosphorio chloride, PC15, in admixture with the vapour of
phosphorous chloride, PC15. (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 atmo-
spheres. 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 Oay-Lussac (for the expansion of gases
302
PRINCIPLES OF CHEMISTRY
Such an investigation (either direct, or by calculation from the densi-
ties and composition) of every chemical reaction, resulting in the forma*
with f). 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.
fio. 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 glone, the latter
is sealed up and weighed. The capacity of the globe is
then measured, and in this manner the volume occu-
pied by a known weight of vapour at a known tem-
perature is determined.
PlO, 88.— Deville and Troost's apparatus
for determining the vapour densities,
according to Dumas' method, of sub*
stances which boil at. high temperatures.
A porcelain globe containing the sub-
stance whose vapour density is to be
determined is heated in the vapour of
mercury (360°), sulphur (410°), cad*
tnium (860°), or zinc (1,040°). The glob*
Is sealed up in an oxyhydrogen flame.
mentioned is essentially volumetric, because a known weight of a substance is taken,
•ltd the volume of the air displaced by the vapour at a given t and h is determined.
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. 62 and 58, 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 boil*
ing point of the liquid : this gives a vapour which displaces the air, and fills the sphe-
rical 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
MOLECULES AND ATOMS
803
tion of definite chemical compounds, shows that the volumes of the react-
ing substances in a gaseous or vaporous state are either equal or are in,
PlO. 54.— Hoftnann's apparatus tor
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
(he 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.
FlO. 66.— Victor Meyer's apparatus fop
determining vapour densities. The
tube 6 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 </. The air displaced is
collected in the cylinder e, iti the
trough/
weighing, and, if quite full of the liquid, breaks when heated in a vacuum) is intro-
duced 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 (.
The method of displacement (c) proposed by Victor Meyer is based on the fact that a
space 6 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 quan-
804 PRINCIPLES OF CHEMISTRY
simple multiple proportion.8 This forms the first law of those discovered
by Gay-Lussac. It may be formulated as follows : The amount* of sub-
stances 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, NH4C1, there enter into reaction
17 parts by weight of ammonia, NH3, which is 8'5 times denser than
hydrogen, and 36'5 parts by weight of hydrogen chloride, whose vapour
density is i8'25 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, NH3, is equal to two, and so also the
volume of hydrogen chloride. Hence the volumes of the compounds
which here combine together arer 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
tity 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. 65.
9 Vapours and gases, as already explained in the second chapter, are subject to the
same law*, 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 dearly 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) forthe variation of v with p and t (for example,
van der Waals' formula, Chapter II., Note 83), so also a -more exact expression of the
relation between the composition and the density of vapours and gases is also possible.
Bat 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.
MOLECULES AND ATOMS 80S
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 spacei
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 ofi
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 con-
tain £ v of oxygen and § 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 waterj
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 f v — that is, it is equal to the volume of the hydrogen contained
MOLECULES AND ATOMS 807
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 j 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.
Comparisons of various results made by the aid of direct observa-
tions 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 cf 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.8
When a compound is formed 'from two 6r 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
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 O, D, Sec., 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. There-
fore 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.
808 PRINCIPLES OF CHEMISTRY
compound. The reverse is naturally observed in the case of decom-
positions, 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 pro-
duced ; while those reactions in which the volumes in a gaseous or
vaporous state remain constant (the volumes being naturally com-
pared 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 sub-
stitution, 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 contrac-
tion occurs, the- properties of the, resultant compound are very dif
ferent from the properties of the substances of which it is composed.
Thus ammonia bears no resemblance in its physical or chemical pro-
perties 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 sub-
stances there is formed a liquid substance, or at any rate one which is
feasily liquefied. For this reason nitrous oxide formed by the conden-
sation 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
MOLECULES AND ATOMS 809*
not destroyed, or is only altered to a slight extent ; hence, notwith-
standing the contraction (compression) involved in its formation,
nitrons oxide supports combustion.
The preceding laws were deduced froni 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 equa.1 ; 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 Ampere (1815)
for the sake of simplifying .all kinds of physico-mathematical concep-
tions respecting gases. But Avogadro and Ampere'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 j 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
6 It mast 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 conse-
quences exactly agreeing with facts; others only take root by slow degreess and there
are many which are destined to be. related owing to their consequences being found to
be at variance with foots.
310 PRINCIPLES OF CHEMISTRY
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 tem-
peratures 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 im-
portant, and therefore, before developing the law and its consequences,
we will consider the chemical phenomena from which the law is deduced
or which jt serves to explain.
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, NH3, reacts directly with
hydrochloric acid, HC1, forming sal-ammoniac, NH4C1, 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,
C2H4, combines with chlorine, C12, in only one proportion, forming
ethylene dichloride, C2H4C12, 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, HC1, and in this case equal volumes in-
teract 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, CH4, and
carbonic anhydride, CO2, and in the proportions in which they are
formed from acetic acid they occupy equal volumes. Also from
phthalic acid, C8H6O4, there may be obtained benzoic acid, C7H6O2.
and carbonic anhydride, CO2, and as all the elements of phthalic acid
enter into the composition of these substances, it follows that, although
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 whoio volume was the greater.
MOLECULES AND ATOMS 811
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 benzole acid, C7H^O2) is itself
composed of benzene, C0H6, and carbonic anhydride, CO2, which also
occupy equal volumes.8 There is an immense number of similar
examples among those organic substances to whose study Gerhard*
consecrated his whole life and work, and he did not allow such facts aa
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 sub-
stances acting on each other occupy equal volumes as well as each of.
the two resultant substances. Thus, in general, reactions of substi-
tution 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, corre-
sponding 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.
But the question arises, What is the relation of volumes if the
reaction of two substances takes place in more than one proportion, acoord-
ing to the law of multiple proportions ? A definite answer can only bo
given in cases which have been very thoroughly studied. Thus chlorine,
in acting on marsh gas, CHA, forms four compounds, CH3C1, CH2Clj,
CHC13, and CC10 and it may be established by direct experiment that
the substance CH3C1 (methylic chloride) precedes the remainder, and
that the latter proceed from it by the further action of. chlorine. And
this substance, CH3C1, is formed by the reaction of equal volumes of
8 Let os demonstrate this by figures.' From 122 grama of benzoic acid there ace
obtained (a) 78 grama of benzene, whose density referred to hydrogen •= 39, hence the
relative volume — 9 ; and (6) 44 grams of carbonic anhydride, whose density •> 22, and
hence the volume = 2. It- is the same in other cases.
* A large number of such generalised reactions, showing reaction by equal
volumes, occur in the case of the hydrocarbon derivatives, because many of these com-
pounds, are volatile. The reactions of alkalis on acids, or anhydrides on water, &o.,
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, HLjSOj,' breaks up into the anhydride, SO3) and water, H2O, which exhibit an equality
of volumes. Let us take another example where three substances combine in equal
volumes : carbonic anhydride, CO2, ammonia, NH3, and water, ELjO (the volumes of all an
equal to 2), form acid ammonium carbonate, (NH4)HCOj.
812 PRINCIPLES OF CHEMISTRY
marsh gas, CH4, and chlorine, C12, according to the equation CH4 + C12
= CH3C1 4- HC1. 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 clear-
ness, 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 N2O3 and NO2. Such a sequence in the combination of nitro-
gen with oxygen can only be supposed on the basis of the fact that NO
forms N2O3 and NO2 directly with oxygen. If it be admitted that NO
(and not N2O or NO2) 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 de-
composed 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 per-
oxide of hydrogen, because up to now such a reaction has not been
observed, whilst the formation of H20 from H2O2 is very easily re-
produced.10
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 un-
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 gad by taking photographs, which showed the
periods of combustion and the waves of explosion, which should be taken into con-
sideration in the theory of this subject. Aa the formation of H2O2 from O2 and H2
corresponds with a less amount of heat than the formation of water from H2 and O, it
may be that the temperature of the flame of detonating gas depends on the pro-formation
of hydrogen peroxide.
MOLECULES AND ATOMS 818
equal volumes, although, in this case, it is often possible to discover a
preceding reaction between equal volumes.11
The law of Avogadro-Gerhardt may also be_eHsily 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
11 The possibility of reactions between unequal volumes, notwithstanding the general
p plication of the law of Avogadro-Gerhardt, may, in addition to what has been said
abo-e, depend on the fact that the participating substances, at the moment of reaction,
• undergo a preliminary modification, decomposition, isomeric (polymeric) transformation,
&c. Thus, if NO2 seems to proceed from N2O4,^if O2 is formed from Os, 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 No, or H5 from H2. In this manner it is obviously possible, by means
of a series of hypotheses, to explain the cases of the formation of ammonia, NH-, 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 N2H4 (Chapter VI. Note 20 bl5) is formed and the imide N2H2 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 tbat diamide gives nitrogen and
ammonia (8N2H4 = N2 + 4NH3) under the action of sparks, heat, or the silent discharge,
<fec., then it will be possible to admit that it is formed before ammonia. And perhaps
the still less stable imide N2H2, which may also decompose with the formation of
ammonia, is produced before the amide N2H4.
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 hypo-
thesis, consequences must be deduced from it, and if by then- means clearness and con-
sistency 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 the1 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 sub-
etances. 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 co'mplexity and uniformity in diversity.
And this is the highest problem of science.
314 PRINCIPLES OF CHEMISTRY
of vapour, according to the law, a volume equal to that occupied by the
molecules of other bodies, be indicated by the letters M,, M2 ............
or, in general, M, and if the letters D,, D2, ............ 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 con-
ditions of temperature and pressure, then the law requires that
M, _ Mj _ M c
D, D2 ' D
where 0 is a certain constant. This expression shows directly that the
volumes corresponding with the weights M,, M2 ............ 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 in
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 H20, for which M = 18, if H=l,
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
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 loeight 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,
NH3, the weight of the molecule or quantity of the reacting sub-
stance, 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, N20, is 22, of nitric acid 15, and of nitric peroxide 23.
In the case of nitrous anhydride, N2O3, as a substance which dissociates
into NO + N02, the density should vary between 38 (so long as the
MOLECULES AND ATOMS
N.2O3 remains unchanged) and 19 (when NO -f N0.2 is obtained).
There are no figures of constant density for H2O2, NHO3, N2O4, and
many similar compounds which are either wholly or partially decom-
posed in passing into vapour Salts and similar substances either have
no vapour density because they do not pass into vapour (for instance,
potassium nitrate, KNO3) without decomposition, or, if they pass
into vapour without decomposing, their vapour density is observed
with difficulty only at very high temperatures. The practical de-
termination 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
theto 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
M
evidently be of two kinds : the fraction — 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 — 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
M
of =- be several times, and in general a multiple, greater or less than 2.
816 PRINCIPLES OF CHEMISTRY
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 -^
is greater than 2, or the density obtained by experiment is less than is
in accordance with t/te 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
M
quotient ^- 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 — <fec. 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-
12 As the density of aqueous vapour remains constant within the limits of experi-
mental 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 bo 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 vapc ur density
of water does not give us the means of ascertaining the amount of its dissociation.
MOLECULAK AND ATOMS 817
ammoniac, NH4C1, 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 NH4C1, because it is formed
from the molecules NH3 and HC1, and contains single atoms of nitro-
gen 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
M
the observed density (about 13'4) ; hence — =;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 other-
wise, 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 NH3 and HC1, which should be
nearly 14, because the density of NH3=8'5 and of HC1=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 decora •
position 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, NH3, is 8|, and of hydrochloric
15 This explanation of the vapour density of sal-ammoniaCj 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.
818 PRINCIPLES OF CHEMISTRY
Acid, HC1, ISty, if the density of hydrogen = 1. Consequently, in a
mixture of NH3 and HC1, the ammonia will penetrate much more
rapidly through a porous mass, or a fine orifice, than the heavier hydro-
chloric 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-Gechardt.15 bi5
So also the fact of a decomposition may be proved in the other
instances where =<- proved greater than 2, and hence the apparent
•deviations appear in reality as an excellent proof of the general applica-
tion and significance of the law of Avogadro-Gerhardt.
16 It is important that the tubes, asbestos, and sal-ammoniac should be dry, as other-
wise the moisture retains the ammonia and hydrogen chloride.
14 b" Baker (1894) showed that the decomposition of NH4C1 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 PjOj) 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 playfe 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).
MOLECULES AND ATOMS 319
In those cases where the quotient ~ 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 CH2, the
density was found by experiment to be 14, and in the case of amy-
-, lene, whose composition is also CH2, the density proved to be 35, and
i consequently the quotient for ethylene = 1, and for amylene = f. II
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 C2H4, and of the second C6HIO, and for
( M
; both of them-- 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-jSO^ Br2, or HI, <fec., ethylene com-
.bines in a quantity C,H4, and amylene in a quantity C3H10, and not
•CH2. On the other hand, etbylene is a gas which liquefies with diffi-
culty (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
\r
smaller quotient — . is an indication of polymerisation, as the larger
^quotient is of decomposition. The difference between the densities of
toxygen 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
ias = 200, but the vapour density = 100, and the quotient = 2. Con-
isequently 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 mole-
cule 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
*2
320 PRINCIPLES OF CHEMISTRY
molecule must contain two atoms, O2) which might already be Concluded''
by comparing its density with that of ozone, whose molecule contains)
O3 (Chapter IV.) So also the molecule of hydrogen equals H2, ofl
chlorine C12, of nitrogen N2, &c. If chlorine react with hydrogen, the
volume remains unaltered after the formation of hydrochloric acid,
H2 + C12 = HC1 + HC1. 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 instante, sulphur,
SG — although, by heating, the density is reduced to a third, and S2 isi
formed. Judging from the vapour density of phosphorus (D = 62)'
the molecule contains four atoms P4. 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, a&' otherwise its non-volatility and
other properties cannot be understood. And if compounds are de-
composed 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 N2O4 passes into NO2, or ozone, O3, into ordinary
oxygen, O2, then we might expect to find the splitting-up of the com-
plex molecules of elements into the simplest molecule containing a
single atom only — that is to say, if 02 be obtained from 03, 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 I2. 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 1C 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 I2 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 N2O4 and the
formation of a large proportion of NO2( with a decrease of pressure. The decomposition
of I2 into I + 1 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
MOLECULES AND ATOMS 821
Besides these cases of apparent discrepancy froffl the law of
Ayogadro-Gerhardt there is yet a third, which is the last, and is very
instructive. In ihe 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 import-
ant significance of the law it becoines 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, S2OjCl2,
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 S03 and S02C12, or any other two products, and
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,
1 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 0'008C7, 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 Clj is lighter (s=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 corre-
sponds 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.
822 PRINCIPLES OF CHEMISTRY
thus the abnormal density of S2O5C12 remained unexplained until D P
Konovaloff (1885) showed that the previous investigators were working
with a mixture (containing SO3HC1), 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 deter-
mination 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 learn-
ing— physics and chemistry — in the most intimate manner. Besides
which, the law of Avogadro-Gerhardt places the conceptions of mole-
cules 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 physi-
cal 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 know-
ledge 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.
The chemical particle or molecule must be considered as the
19 And so it was in the fifties. Some took O=8, others O = 111. Water in the first
CMC would be HO and hydrogen peroxide HO.,, and in the second case, as is how gene*
rally accepted, water H?O and hydrogen- peroxide HnOo or HO. Disagreement and con-
fusion reigned. In I860 the chemists of the whole world met at Carlsrnhe 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 con-
sequences 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
•eventiea they were already in general use.
MOLECULES AND ATOMS 323
quantity of a substance which enters into chemical reaction with other'
molecules, and occupies in a stale of vapour the same volume a$ two
parts by weight of hydrogen.
The molecular weight (which has been indicated by M) of a suV
stance 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
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, <fec.) 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
19 A bubble of gas, a drop of liquid, or the smallest crystal, presents an agglomera-
tion 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.jO.j and N0.>. 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. Tho
laws of chemical equilibrium proper to dissociated systems cannot ,be regarded in any
other light.
80 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 o!
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 wheu our
knowledge of the internal structure of matter will be defined and established.
824
PRINCIPLES OF CHEMISTRY
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.
Atoms are the smallest quantities or chemically indivisible masses
of (he 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 exist-
ence, their incapacity of being converted into each other, <fcc., but also
by the weight of their atoms.
Chemical and physical properties depend on the weight, composi-
tion, 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 Avog'adro-
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 mole-
cules 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 : —
H2O
40
JNJO
NO,
Mqleculai1
Weight
18
44
30
46
Amount of
Molecular
Amount of
Oxygen
Weight
Oxygen
16
HN03
63
48
16
CO
28
16
16
CO2
44
32
32
The numbef of substances taken might be considerably increased,
but the result would be the same— that is, the molecules of the com-
pounds of oxygen would never be found to contain less than 1 6 parts
by weight of this element, but always nlQ, where h is a whole number.
MOLECULES AND ATOMS 825
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 decom-
poses above its boiling point) cannot be accurately determined, but tho
fact of its containing one part by weight of hydrogen, and all its pro-
perties 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
conf pounds. It may, for instance, be easily proved that less than w!2
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 bxygen, 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 BeCl2 — 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 Be2O3,
and the chloride by BeCl3.21 It was then found that the vapour density
vof beryllium chloride was approximately 40, when it became quite clear
that its molecular weight was 80, and as this satisfies the formula
BeCl2, but does not suit the formula BeCl3, it therefore became neces-
sary to regard the atomic weight of Be as 9 and not as 13^.
" If Be = 9, and beryllium chloride be BeCl2, then for every 9 parts of beryllium
there are 71 parts of chlorine, and the molecular weight of BeCl2 = 80 ; hence the vapour
density should be 40 or n40. If Be = 18'5, and beryllium chloride be BeClj, 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 «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 formulae 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 = l, 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 con-
crete 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).
326 PRINCIPLES OF CHEMISTRY
With the establishment of a true conception of molecules and
atoms, chemical formulae became direct expressions, not only of com-
position,22 but also of molecular weight or vapour density, and conse-
quently of a series of fundamental chemical and physical data, inasmuch
as a number of the properties of substances are dependent on their
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 : — HC1 shows
that hydrochloric acid contains 85'5 of chlorine and 1 part of hydrogen. Hence, in 86'5
parts of hydrochloric acid there is 1 part by weight of hydrogen, consequently 100 part*
by weight of hydrochloric acid will contain as many more units of , hydrogen as 100 ia
greater than 86'5; therefore, the proportion is as follows — x 1 " 100 86'5 or # = _!:?
*56'5
=• 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. 6'88 of hydrogen and 94'12 of
oxygen, it is easy to find its formula; jt 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 peroxfde 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 per'
centage amounts by tjie 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 AxBy. From this formula we know that our compound
contains xA parts by weight of the first element, and 2/B of the second. In 100 parts of
our compound there will be (by proportion) __£ of the first element, and — — '^ ®
' zA + i/B
of the second. Let us divide these quantifies, expressing the percentage amounts by the
corresponding combining weights ; we then obtain • for the first element and
XA+yo
— - — ~^ for the second element. And these numbers are in the ratio x : y — that is, in
a-A + yB
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 CO2 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,
S02, expresses the same as barium dioxide, Ba02, but sulphurous anhydride indicates
the acid properties of S02. Probably in time one harmonious chemical language will
succeed in embracing both advantages.
MOLECULES AND ATOMS 827
vapour density, or molecular weight and composition. The vapour
M
density D= , For instance, the formula of ethyl ether is C4H10O,
Si
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 formulae,
because the molecular weights in grams of all oilier vapours at 0° and
760 mm. occupy the same volume, 22 '3 litres Thus, for example, in the
case of carbonic anhydride, C02, 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 T97 gram. By combining the
laws of gases. — Gay-Lussac's, Mariotte's, and Avogadro-Gerhardt's — we
obtain23 a general formula for gases
6200s(273-M)=M/>
where « 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 mer-
cury) 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
«=0-0024.24
}I This formula (which is given in my work on ' The Tension of Gases,' and in a
somewhat modified form in the ' Comptes Bendus,' 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 s0 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 « of a cubic centimetre of a gas at a temperature t and under a pressure p
(in centimetres) is equal to «0p/76 (1 + at). Therefore, s0—s.76' (1 + a£)p; hence D =•
76.« (1 + at) '0'0000898p. whence M = 152s (1 + of)/0'0000898p, which gives the above expres-
sion, because I/a =278, and 152 multiplied by 373 and divided byO'0000898 is nearly 6200.
In place of «, mfv may be taken, where m is the weight and v the volume of a vapour.
S4 The above formula may be directly applied in order to ascertain the molecular
weight from the data ; weight of vapour m grins., its volume v c.c., pressure p cm., and
temperature t°; for,-»=the weight of vapour m, divided by the volume v, and conse-
quently M= 6,200rw (273 + t)/j>v Therefore, instead of the formula (sec Chapter II.,
Note 84), j>v=R(27S + <)i where R varies with the mass and nature of a gas, we may
apply the formula pv •* 6,200(m/M) (273 + 0- These formulae simplify the calculations
in many cases. For example, required the volume v occupied by 5 grms. of aqueous
vapour at a temperature 1 = 127° and under a pressure j>"=76 cm. According to the
formula M = 6,200 m (273 f t)lpv, we fiud that v = 9,064 c.c., as in the case of water
M = 18, m in this instance = 5 grms. (These formula, however, like the laws of gases,
are only approximate.)
328 PRINCIPLES OF CHEMISTRY
As the molecules of many elements (hydrogen, oxygen, nitrogen,
chlorine, bromine, sulphur — at least at high temperatures) are of uni-
form composition, the formulae of the compounds formed by them
directly indicate the composition by volume. So, for example, the
formula HNO3 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 vajwurs ami gases is alike. For
it is proved by mechanics that the vis viva of a moving mass=£ mv*t
where m is the mass and v the velocity. For a molecule, ?n=M, or
the molecular weight, and the velocity of the motion of gaseous
inolecules=a constant which we will designate by C, divided by the
square root of the density of the gas 25=C/N/D, and as D=M/2,
the vis viva of molecules=CJ — that is, a constant for all molecules.
Q.E.D.™ 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 Jieat, which also proves to be
in connection with the molecular weight. The observed latent heats
K Chapter I., Note 84.
*6 The velocity of the transmission of sound through gases and vapours closely
bears on this. It = J EpglT> (1 + at), where E. 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, a=0'00367, and t the temperature. Hence, if K
be known, and as D can be 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 A". The relative velocities of sound in two gases can be easily determined (Kundt).
If a horizontal glass tube (about 1 metre long and closed at both ends) be full of(»
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.
MOLECULES AND ATOMS 829
of carbon bisulphide, CS2=90, of ether, C4H10O, = 94, of benzene,
C6H6, = 109, of alcohol, C2H6O, = 200, of chloroform, CHC13, = 67,
4c., 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 be referred to the weight of the
molecule. For carbon bisulphide the formula CS2 expresses a weight
76, hence the latent heat of evaporation referred to the molecular
quantityCS2 = 76 x 90=6,840, for ether=9,656,forbenzene=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
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, <fec.), 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 dis-
tance) and the masses of substances which take effect, while those
peculiarities of a substance which are expressed in chemical trans-
formations, 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,
17 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 obser-,
vation. 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 thermo-dynamics the latent heat
of evaporation is equal to (n1 — n) 2* 18'69, where t is the boiling point, n' the
E (iT
specific volume (i.e. the volume of a unit of weight) of the vapour, and « the specific
volume of the liquid, dp'dT the variation of the tension with a rise of temperature per 1°,
and 18'89 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.
830 PRINCIPLES OF CHEMISTRY
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 dis-
persed 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 mole-
cules 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 ex-
plaining 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 com-
position, and even give the means of determining the weight of mole-
cules 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, electro-
conductivity, index of refraction, ifec.) we will select one — the 'depres-
sion ' 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
MOLECULES AND ATOMS 881
also many experimental discrepancies which cannot as yet be ex-
plained by theory.271"1*
» tu 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 deter-
minations it separates out in the form of a solid body.
Van't Hoft, starting from the second law of thermo-dynamics, showed that the
dependence of the rise of pressure (dp) upon a rise of temperature (dl) ia determined by
the equation dp = (limp /2T-) dT, where k is the latent heat of evaporation of the solvant,
m its molecular weight, p the tension of the saturated vapour of the solvent at T, and T
the absolute temperature (T = 278 + <), while Raoult found that the quantity (p— p') Ip
(Chapter I., Note 50) or the measure of the relative fall of tension (p the tension of the
solvent or water, and^' 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') lp = Cn/(T$ + 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 vahie of N is very much
greater than "n), and then, judging' from experiment, C is nearly unity — hence:
rfp/p = w,'N or dp = np,'N, and on substituting this in the above equation we have
(kmp /2T8) dT = 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 = nM, 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 = ~£r~' *na* '8> ^y taking a solution of q
grins, of a substance in 100 grms. of a solvent, and determining by experiment the rise
of the boiling point dT, we find the molecular weight M of the substance dissolved,
because the fraction 0-02 T* Ik is (for a given pressure and solvent) a Constant ; for water
at 100° (T = 878°) when A' = 534 (Chapter I., Note 11), it is nearly 5'2,for ether nearly 81,
for bisulphide of carbon nearly 24, for alcohol nearly ITS, &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, HgCl2, was taken
in the quantity q = 8'978 and 4'258 grms:, the rise in the boiling point dT was = Oc>179 and
0°'084, whence M = 261 and 268, and when alcohol was the solvent, q = 10'878 and 8'765
and dT = 0°-471 and 0°-880, whence M = 266 and 265, whilst the actual molecular weight of
corrosive sublimate =271, which is very near to that given by this metho.d. In the
same manner for aqueous solutions of sugar (M = 842), when q varied from 14 to 2-4, and
the rise of the boiling point from 0°'21 to 0°'085, M was found to vary between 889 and
864. For solutions of iodine I.j in ether, the molecular weight was found by this method
to be between 255 and 262, and Io = 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 Barpni (1898) 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 HgClj (from 255 to 834 instead of 271), KNO5 (57-66 instead of 101), AgNO3 (104-107
Instead of 170), KaSO4 (55-89 instead of 174), sugar (328-848 instead of 342), &c. On the
contrary the calculated value of M decreases as the concentration increases, for solu-
tions of KC1 (40-89 instead of 74'5), NaCl (33-28 instead, of 58'5), NaBr (60-49 instead
of 108), &c. In this case (as also for LiCl, Nal, CoH3NaO.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
'332 PRINCIPLES OF CHEMISTRY
If 100 gram-molecules of water, i.e. 1,800 grms, be taken and n
'gram-molecules of sugar, C^H^O,, ,i.e. 7*342 grms., be dissolved in
[greater. For example, according to Schlamp (1894), for LiCl, with a variation of from
I'l to 6'7 grm. LiCl per 100 of water, i varies from 1'63 to 1'89. But for substances
of the first series (HgCl2, &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 HgCU, MgSO4, &c., i exhibits a similar variation;
-thus, for HgClj the value of M is found to vary between 255 and 334 ; that is, i (as
the molecular weight = 271) varies between T06 and 0'81. Hence I do not believe that
' the difference between i and unity (for instance, for CaCl2) 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 = l + 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 Jc 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 contain-
ing 5'67 grms. CaCla 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 1'88. 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 formulae
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' . "?. N=a constant C
p am
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 in 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' Id give, for water T08, for alcohol 1'02, for ether T04, for
bisulphide of carbon I'OO, for benzene 1'02, for acetic acid 1'63. 3. By applying the
principles of thermodynamics and calling Lj the latent heat of fusion and Tt the
absolute ( = t + 273) temperature of fusion of the solvent, and Lj and T4 the corresponding
values for the boiling point, Van't Hoff in 1886-1890 deduced :—
Depression of freezing point LJJ T|*
Rise of boiling point LI ' T8*
Depression of freezing point = 'T- * *
LlMl
Rise of boiling point = A^£?
L?M|
where A = D'01986 (or nearly 0 03 M we took it above), a is the weight in grms. of the
MOLECULES AND ATOMS 33$
them, then the depression d, or fall (counting from 0°) of the tempera-
ture of the formation of ice will be (according to Pickering)
« = 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 ap-
proximately (estimating the possible errors of experiment at ±0°-005)
-nl-10, because then d=Q°, 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 dis-
solving in 100 molecules of water give a depression of about l°'l n.
Similar data for acetone (Chapter I., Note 49) give a depression of
l°-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 100H40 is nearly nl°-l to nl°*0 (ether, for instance,
gives the last number), and consequently in dissolving in 100 grms. o£
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, ace-
tone, 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 Kn. 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, <tc.) for
all dilute solutions. Thus, having found a convenient solvent for A
given substance and prepared a definite (by weight) solution (i.e. know-
ing 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 deter-
mine the molecular weight of the s'ubstance 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 per 100 grins, of the solvent, Mt 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] . The experimental data
and theoretical considerations upon which these formula; are bused will be found in text-
books of physical and theoretical chemistry.
884 PRINCIPLES OF CHEMISTRY
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 = and X is the mole-
fc
cular weight of the substance dissolved. Hence X = — -i- , which
rd
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 H202 or H3O3, ike. (H2O2 is
obtained). Moreover, attention should be drawn to the fact that
a great many substances taken afc 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 H2O.28 A similar pheno-
menon which repeats itself in the osmotic pressure, vapour tension
of the solvent, <fec. (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
18 A similar conclusion respecting the molecular weight of liquid water (i.e. that ita
molecule in a liquid state is more complex than in a gaseous state, or polymerized into
H6O4, HtfO; or in general into nH^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 pas-
sage 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 Eotvos (1886) and Ramsay and Shields (1898) on the variation
of the surface tension N with the temperature (N = the capillary constant a3 multiplied
by the specific gravity and -divided by 2, for example, for water at 0° and 100° the value
of a*=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 waa necessary
from all the data obtained, and calling this temperature T, it is found that
AS = AT, where S is the- surface of a gram-molecule of the liquid (if Mis its weight
in grams, s its sp. gr., then its sp. volume = M /», and the surface 8 =$/ (M/«)'), A the
surface tension (determined by experiment at T), and k a constant which is inde-
pendent of the • composition of the molecule. The equation A9 = kT ia in complete
agreement with the well-known equation for gases t>^j = 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 CS2, ether, benzene, and of many other
substances, have the same value as in a state of vapour, whilst with other liquids this ie
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^ to 8J, for water from 2J 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 obervations 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.
MOLECULES AND ATOMS 885
show (according to Pickering's data 1892) for solutions of NaCl and
CuSO4 in water. For
n = O01 0-03 0-05 0-1 0-5
molecules of NaCl the depression is
rf = 0°-0177 0°-0598 0°-0992 0°-1958 0»-9544
which corresponds to a depression per molecule
K=l-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 I'l n. For CuSO«
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 100H2O,
when d = Kn, for very dilute solutions of CaClj is nearly 2 '6, for
Ca(NO3)2 nearly, 2'5, for HNO3, KI and KHO nearly 1-9— 2'0, for
borax Na.2B407 nearly 3-7, <fcc., while for sugar and similar substances
it is, as has been already mentioned, nearly I'O — I'l. Although these
figures are very different 28bl8 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 mole-
cular weight of a substance in solution.29
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
Je bis Their variance is expressed in the same manner as was done by Van't Hofl
(Chapter I., Notes 19 and 49) by the quantity t, taking it as «= 1 when ft = i'05, in that case
for KI, i is nearly 2, for borax about 4, &c. -
M 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 8,600 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
weight
Sp. gr. at 15°
Molecular
weight
Sp. gr. at 15
HC1
86-5
1-0041
CaCL,
111
1-0236
NaCl
58-5
1-0106
NiClg
180
1-0828
KC1
74-6
1-0121
ZnClg
186
1-0331
BeCl2
80'
1-0138
BaClj
208
1-0489
MgCl,
95
1-0208
836 PRINCIPLES OF CHEMISTRY
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, <fec.) 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 composi-
tion, inasmuch as dense elements (in a free state) and compounds are
only met with among substances containing elements with large atomic
wejghts, 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 com-
position and the properties of the component elements.29 bis The history
28 bb 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 Frannhofer (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 ; (6)
or Cauchy's formula is used, showing the relation between the coefficient of refraction and
T»
dispersion to the wave length n = A + — , where A and B are two constants varying
for every substance but constant for all rays of the spectrum, and A 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 mole-
cular 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 O'OOOOO, and their atomic weights, O = 16, H = l ; hence their refraction equivalents
are 8 and 1-5. Water contains H2O, consequently the sum of the equivalents of
refraction is (2 *. 1-5) + 8 = 6. But as the coefficient of refraction of water =? 1-881,
1U refraction equivalent =5'958, or nearly 6. Comparison shows that, approxi-
mately, the sum of the refraction equivalents of the atoms forming compounds
'MOLECULES AND ATOMS SSf •
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 enter-
ing 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.
(or mixtures) is equal to the refraction equivalent of .the compound. According to the
researches of Gladstone, Landolt, Hagen, Briihl and others, the refraction equivalents of
the elements are — H = 1'8, Li = 8.8, B = 4'0, C = 5'0, N = 4'l (in its highest state of oxida-
tion, 5-8), O = 2-9, F = 1-4, Na = 4-8, Mg = 7'0, Al = 8'4, Si = 6'8, P = l«'8, S = J6'0, 01 = 99,
K = 8-l, Ca=10-4, Mn = 12-2, Fe = 12'0 (in the salts of its higher oxides, 20-1), Co = 10'8,
Cu = ll-6, Zn = 10-2, As = 15'4, Bi = 15'3, Ag=15'7, Cd = 18'6, 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. F6r instance,
the composition of carbon bisulphide is CSo = 76, and from its density, 1'27, we find its
coefficient of refraction to be TCIS (because the refraction equivalent =5 + 2x16 = 87),
which is very near the actual figure. It is evident that in the above representation com-
pounds are looked on as simple mixtures of atoms, and the physical properties of a com-
pound 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 refrac-
tion of different substances could hardly have arisen. For further details on this subject,
see works on Physical Chemistry.
838 PRINCIPLES OF 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 in-
soluble. 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 char-
coal, 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
CARBON AND THE HYDROCARBONS 339'
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, I
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
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
incrasting matter has been deposited. The tissue has the composition C6H10O5, 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 86 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 '765 when in a fresh
state ; when in a dry state birch weighs 0'622, pine 0'550, fir 0'356, lime 0*480, guaiacurn
1-342, ebony T226. 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
6,000 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 850J black charcoal is
produced. As the hydrogen contained in wood requires for its combu etion 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 corre-
spond with its percentage ratio — that is forty parts per 100 parts of wood. Generally,
however, the quantity produced is far less, not more than 80 p.c., because part of the
carbon is given off as gas, tar, &c. If wood has to b» transported great distance*, or if
340
PRINCIPLES OF CHEMISTRY
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.
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.
The composition of the growing parts of plants, the leaves, young branches, shoots, &o.,
differs from the composition of the wood in that these vital parts contain a considerable
K
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, 18 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 30 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
CARBON AND THE HYDROCARBONS 841
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 pro-
ducts of transformation of the organic matter are subjected to conden-
sation (for example, if the door of a stove be opened), an imperfect\
combustion takes place, and smoke, with charcoal or soot, is formed.2/
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 sub-
stances, 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, CH4O, and acetic acid, C2H4Oo. In such cases, the dry dis-
tillation 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 pro-
ducts 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.
* 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 forma-
tion 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 com--
bust ion, and to combat against such a ruinous practice as the imperfect combustion of i
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
la 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 thiflj
842
PRINCIPLES OF CHEMISTRY
y
^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 hjdrogen
is separated in a^ gaseous state, another part in combjnation__with
oxygen, and a third part separates in combination witfr carbon, and
sometimes in combination with cajrbon jmd_ jxsygen 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-
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 con-
tact of the red-hot burning charcoal. An air supply through the fire grate, an equal dis-
tribution 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 with-
out smoke, because they do not contain hydrogenous substances which furnish the pro-
ducts of dry distillation, but imperfect combustion may occur with them also ; in that
case the smoke contains carbonic oxide.
5 Under the action of air, organic substances are capable of oxidising to such an extent
that all the carbon and all the hydrogen the y'contain will be transformed into carbonic anhy-
dride 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 aapply^of air h« limited, t.hprp can be no complete transformation into water anci
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 u
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 wator — 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 decom-
posing agent. The annexed table shows, according to the data of Violette, those changes
i
CARBON AND THE HYDROCARBONS
848
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 hydro
gen, and also substances which ure capable of taking up water, can
also separate carbon from (or char) most organic substances.
Wood charcoal is prepared in large quantities in a similar manner
— that is, by the partial combustion of wood.4 In nature a similar
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 charcoal
of alder wood
C
H
OandN
Ash
150°
lOO'O
47-5
6-1
46-3
o-i
850°
29-7
76-6
4-1
18-4
0-6
1082°
18-7
61-9
2-3
14-1
1-6
1500°
17-8
95-0
0-7
8-8
1-7
4 The object of producing pharcoal 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 manu-
factured 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 ia
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 than
necessary to stop the air holes, in order as much as possible to prevent unnecessary com-
bustion. 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*
J'ourth of its weight of charcoal.
344 PRINCIPLES OF CHEMISTRY
process of carbonisation of vegetable refuse takes place in its trans- /""\
formation under water, as shown by the marshy vegetation which S
forms peat.6 In this manner doubtless the enormous masses of
coal were formed 6 which, following the example set by England, are
* 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
ft 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 dis-
solves 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 sub-
stances 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 Dokou-
chaeff's works.
If those substances which produce humus undergo decomposition under water, less
;carbonic anhydride is formed, a quantity of marsh gas, CH4, 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 Bussia it is likewise found in large quantities, especially in
the North- West districts. The old hard forms of peat resemble in composition and -pro-
perties 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.o. of
carbon, 4 p.c. of hydrogen, 1 p.c. of nitrogen, and 28 p.c. of oxygen. Its heating powe*
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
|>eat, 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 primaeval 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, pre-
served 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
CARBON AND THE HYDROCARBONS
345
now utilised everywhere as the principal material for heating steam
component parts (peat and lignites, even in their last condition, still continue to evolve
nitrogen, carbonic anhydride, and marsh gases) fonn coal. Coal is a dense homogeneous
mass, black, with an oily or glassy lustre, or more rarely dull without any evident vege-
table 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 bitu-
minous coal) to 1'G (anthracite, flameless), and even reaches I'D in the very dense variety
of coal found in the Olonetzky government (termed thungite), 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,
ts as follows: 84 parts of carbon, 6 parts of hydrogen, 1 part of nitrogen, 8 parts of
oxygen, 2 of sulphur. The quantity of ash is on an average 6 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 com-
position of coal we saw that for 5 parts of hydrogen there were 8 parts of oxygen ;
therefore 4 parts by weight of the 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, 8 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 = l'5. Therefore the anthracites contain
but little hydrogen capable of combining with the carbon to fonn hydrocarbons which
burn with a flame. Anthracites are the oldest forms of coal. The newest and least trans-
formed 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- Griiner 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.
o ?
846 PRINCIPLES OF CHEMISTRY
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
During the imperfect combustion of volatile substances containing
purposes (see Note 8). Besides coke, the dry distillation- of coal produces gas (see further,
illuminating gas, p. 861), 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 800
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 con-
snmes about 500 million tons of coal yearly. Besides household purposes, coal is
chiefly used as fuel for steam-engines. As every horse-power ( = 75 kilogramme! res 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 *uy 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' &
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.
CARBON AND THE HYDROCARBONS
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 I
depends on the temperature to which it has been submitted. Charcoal V
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,
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 tartario 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 char-
coal 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 char-
coal 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
348
PRINCIPLES OF CHEMISTRY
capability 'Of absorbing the luminous rays (hence its blackness and
opacity), and many other qualities, are familiar from everyday ex-
perience.11 The specific gravity of charcoal varies from l-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 imme-
diately 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 com-
plete 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,
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.
t'avre.
Heat emitted
Boxwood Charcoal
Cocoauut Charcoal
per gram of gas
NH3 90
172 vols.
494 units
C02 35
97 „
158 „
N,0 40
99 „
169 „
HC1 85
165 „
274 „
The quantity of gas absorbed by the charcoal increases with the pressure, and 18
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 solu-
tions are filtered through charcoal in order to purify them. The bleaching power of
charcoal maybe tested by using various coloured 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.
CARBON AND THE HYDROCARBONS 849
and at an ordinary temperature does not combine ivith 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 12bis 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, CO2>.
whilst in the vapours of sulphur, carbon bisulphide, CS2, is formed,,
and wrought iron, when acted on by carbon, becomes cast iron.
At the great heat obtained by passing the galvanic current through- J
carbon electrodes, charcoal combines with hydrogen, forming acetylene, I
C2H.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
i* 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 wit\i charcoal or coke, because at ordinary tem-
peratures it resists the action of even the strongest acids.
i> 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 BaC2, SrC2, CaC2. All these-Componnds reacti<wlith
water forming acetylene, for example :
§aC^ + 2H2O = C8H2 + 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
terreau furnace (BaCO3 + 8Mg + C = SMgO + BaC2). One hundred grams of BaC2«volve
6,200 to 5,400 c.c. of acetylene, mixed with about 2-8 p.c. of hydrogen.
The relation of acetylene, C2H2, to these metallic carbides is evident from the fact
that these metals (Ca, Sr, Ba) replace 2 atoms of hydrogen, and therefore CoBa corre-
sponds to C2H2, so that they may be regarded as metallic derivatives of acetylene.
Moissan (1893) obtained similar cm-bides directly from the oxides by subjecting them to
the action of the voltaic arc, in the presence of carbon, for instance, BaO + 8C = CO + C2Ba,
although at a furnace heat carbon has no action, on the oxides CaO, BaO, SrO. Con-
cerning AJ4C5, see Chapter XVII. Note 88.
850 PRINCIPLES OF CHEMISTRY
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 ab-
sorbs 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 allo-
tropic 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 con-
trast ; 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 com-
bustibility. 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
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
13 When subjected to pressure, charcoal loses heat, hence the densest form stands to
the less dense as a solid to a liquid, or aa 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
0-21, 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*, and between 250° and 1000°, = 3-54 + 0'00246« (see
Chapter _XIV1 NoteJ).
becomes strongly incandescent, partially volatilises, an'd is deposited in
the form of graphite. If sugar be placed in a charcoal crucible and
a powerful galvanic current passed through J», 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 com-
pact masses, sometimes permeating rocky formations like the schists
or slates, and in fact is met with in those places which, in all proba-
bility, 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
The diamond is a crystalline and transparent form of carbon. It is
1 ' 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 macs by disintegra-
ting, 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 Aliberoftsky graphite is particularly renowned j
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, SidoroS 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 ^ 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, CnH4Os. 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 amor-
phous 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.
352 PRINCIPLES OF CHEMISTRY
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. I5bis 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 in-
applicable 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 de-
posited on the negative electrode, when a weak galvanic current is
passed through liquid chloride of carbon.16 bls
Moissan (Paris, 1893) produced diamonds artificially by means of
the high temperature attained in the electrical furnace 17 by dissolving
is bis Diamonds are found in a particular dense rock, known by the name of
itacolumnite, and are dug out of the debris produced by the destruction of the
itacolunmite by water. When the dibria 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 debris 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 Erofeeff found (1887) diamond powder in a meteoric stone which fell in the Govern-
ment 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.
10 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
etones, and also for rock boring and tunnelling.
16 bit 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.
17 The electrical furnace is an invention of recent times, and gives the possibility of
obtaining a temperature of 8,600°, which is not only not obtainable in ordinary f tfrnaces,
but even in the oxyhydrogen flame, whose temperature does not exceed 2,000°. The elec-
trical 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 amperes a tem-
perature of 8,000° is easily obtained. At a temperature of 2,500° (100 amperes and 40
CARBON AND THE HYDROCARBONS $53
carbon in molten cast iron, and allowing the solution with an
excess of carbon, to cool under the powerful pressure exerted by
rapidly cooling the metal.17 bu K. Chroustchoff attained the same end
by means of silver, which dissolves carbon to the extent of 6 p.c.
volte) not only do all metals melt, but even lime and magnesia (when placed in the space
between the carbon electrodes, i.e. in the voltaic arc) become soft and crystallise on
cooling. At 8,000° lime becomes very fluid, metallic calcium partially separates
out and a carbon compound, which remains liquid for a long time. At this tem-
perature oxide of uranium is reduced to the suboxide and metal, zirconia and rock
crystal fuse and partially volatilise, as also does alumina; platinum, gold, and even
carbon distinctly volatilise ; the majority of the metals form carbides. At such a tem-
perature also cast iron and carbon give graphite, while according to Rousseau, between
2,000° and 3,000° the diamond passes into graphite and conversely graphite into the
diamond, so that this is a kind of reversible reaction.
17 bis Moissan first investigated the solution of carbon in molten metals (and the
formation of the carbides) such as magnesium, aluminium, iron, manganese, chromium,
uranium, silver, platinum, and silicon. At the same time Friedel, owing to the discovery
of the diamond in meteoric iron, admitted that the formation of the diamond is depen-
dent upon the influence of iron and sulphur. With this object, that is to obtain the
diamond, Friedel caused sulphur to react upon samples of casb iron rich in carbon, in a
closed vessel at a maximum temperature of 500°, and after dissolving the sulphide of iron
formed, he obtained a small quantity of a black powder which scratched corundum, i.e.
diamond. Moissan's experiments (1898) were more successful, probably owing to his-
having employed the electrical furnace. If iron be saturated with carbon at a tem-
perature between 1,100° and 8,000$, then at 1,100°-1,200° a mixture of amorphous-
carbon and graphite is formed, while at 8,000° 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 (8) 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 thau 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 8 and 8'5. The transparent
specimens had a greasy appearance and seemed to be, as it were, surrounded by an
envelope of carbon. At 1,050° tljey 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
854 PRINCIPLES OF CHEMISTRY
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.,, and possessed an exceptional hardness, but only
in the form of a fine powder
Judging from the fact that carbon forms a number of gaseous bodies
(carbonic oxide, carbonic anhydride, methane, ethylene, acetylene, <fec.)
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= 1 6, and that the compounds CO (carbonic oxide) and N2C? (cyanogen)
are gases, it may be argued that if carbon formed the molecule C2, like
N2 and O2, it would be a gas. And as'through polymerism or the com-
bination of like molecules (as O2 passes into O3 or NO2 into N204) the
temperatures of ebullition and fusion rise (which is particularly clearly
proved with the hydrocarbons of the CnH2n series), it ought to be con-
sidered 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 C5 . . . C10 . . . C20 • • • C30, &c., in
general Cn, where n may be very large, and in none of the other ele-
ments 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,, where n is a large quantity.
Charcoal and those complex non-volatile organic substances which
represent the gradual transitions to charcoal l9 and form the principal
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.
18 The existence of a molecule Sg is known (up to 600°), and it must be held that this
accounts for the formation of hydrogen persulphide, H2S5. Phosphorus appears in the
molecule P4 and gives P4H2. 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., CnH^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 CnH.^mOm — for example
in the humic compounds (Note 5) — the transition o'f 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
CARBON AND THE HYDROCARBONS 855
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
Hfr othq^ lynatomfgt^jare capable of combining together in
such variety as carbon and hydrogen. The hydrocarbons of the
CftH2m 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
C12H6O3. 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, C6H6, and the
wide diffusion and facility of formation of the carbohydrates containing C3 (for example,
cellulose, C6H10Os, glucose, C6Hi.jO6) give reason for thinking that the group C6 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
\ euch as in carbohydrates.
80 When charcoal burns, the complex molecule Cn is resolved into the simple mole-
cules nCOjj, and therefore part of the heat — probably no small amount — is expended in
the destruction of the complex molecule Cn. 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 Cn into separate atoms.
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 solu-
tions 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. solu-
tion 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 com-
parative 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 CH2), so does the degree of viscosity also rise. The exten-
sive series of investigations referring to this subject still await the necessary generalisa-
tion. 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 fore-
most. Many observations have already been made, bat not much has yet been dome
856 PRINCIPLES OF CHEMISTRY
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 tempera-
ture 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 physi-
cally 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 con-
sidered (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^ofhydrQ-
carbons do not combine with the oxygen of the air or oxidise at ordi-
nary 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, hydro-
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.
23 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,
C2HS, 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 CC14, are not decom-.
posed by water as is the case with phosphorus chloride and even silicic chloride and borio
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(C2H5)2, which corre-
sponds with ethyl hydride or ethane, C^Hg, in which two atoms of hydrogen have been
exchanged for one of zinc.
K 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, C2H2, benzene,
C6H6, naphthalene, C10H8, &c.
24 Wagner (1888) showed that when uusaturated hydrocarbons are shaken with a
i
CARBON AND THE HYDROCARBONS 857
carbons barn, and, according to the amount of carbon they contain,
their combustion is attended more or less with a separation o£ 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 car-
bonic 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«H2m where
n and m are whole numbers. This fact is known as the law of even
numbers. Hence, the simplest possible hydrocarbons ought tot be :\ i^
CH2, CH4, CH6 . . . C2H2, C2H4, C2H6, C2H8 ... but they do not ] T"
all exist, since the quantity of H which can combine with a certain Jl
amount of carbon is limited, as we shall learn directly.
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
CnH2ll+2. These so-called saturated hydrocarbons are incapable of
combination.25 The hydrocarbons CH6, C2H8, C3H10, &c. . . do not
exist. Those containing the maximum amount of hydrogen will be
represented byCK^n = 1, 2n + 2 = 4), C2H6 (n = 2), C3H8 (n = 3), I -/"
C4H10, &c. This maybe 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 formulae C,,H2n+2, CBH2n, CBH2)t_a,
&c. . . Those hydrocarbons which belong to any one of the series
weak (1 p.c.) solution of potassium permanganate, KMnO4, at ordinary temperatures,
ihey form glycols — for example, C2H4 yields C-jI^Oj.
84 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 hydro-
carbons 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.jH.4,
with chlorine, C13, was often regarded as a formation of the products of the substitution
of CjHjCl and HC1, which it was supposed were held together as the water of crystallisa-
tion 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 rX.>, gives a substance
which is saturated or else approaching the limit. The investigations of Franklaud
with many organo-metallic compounds clearly showed the limit in the case of metallic
compounds, which we shall constantly refer to later on.
$58 PRINCIPLES OF CHEMISTRY
expressible by a general formula are said to be homologous with one
another. Thus, the hydrocarbons CH4, C2H6, C3H8, C4H|0, &c. . .
are members of the limiting (saturated) homologous series CnH2,1+2.
That is, the difference between the members of the series is CH2.26
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 CHH2n+2 are not capable of forming additive compounds,
whilst those of the series CHH2n are capable of combining with chlorine,
sulphuric anhydride, &c. ; and the members of the C,,H2l,_6 group,
belonging to the coal tar series, are easily nitrated (give nit ro- 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
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
*6 The conception of hpraology has been applied by Gerhardt to all organic com-
pounds in his classical work, ' Trait6 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 Kekule*, in his beautiful investigations on the
structure of aromatic compounds, showed the presence in them all of the 'benzene
nucleus,' C6H6.
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, CnH2n + i(OH),
when n = 0, we obtain water, H(OH), which boils at 100°, and whose specific gravity ut
15° = 0-9992; when n = l, wood spirit CH3(OH), which boils at 66°, and at 15° has a
specific gravity = 0'7964 ; when w = 2, ordinary alcohol, C2H5(OH), boiling at 78°, specific
gravity at 15° = 0'7936, and with further increase of CELj the specific gravity increases.
For the glycols CnH2n (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 oxyg«n, determines many of the exceptional
properties of water (Chapter I.).
28 As. for example, in the saturated series of hydrocarbons C,,H.w+2, the lowest
member (n = 0) must be taken as hydrogen H2, 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 » = 1, 2, 8, the hydrocarbons CH4, C^H^, C3H8 are gases, more and more readily
liquefiable. The temperature of the absolute boiling point for CH4 = — 100°, and for
ethane C2H6, and in the higher members it rises. The hydrocarbon C4H10, liquefies at
about 0°. C5H12 (there are several isomers) boils at from + 9° (Lvoff) to 87°, C6H,4
from 58° to 78°, &c. The specific gravities in a liquid state at 16° are : —
C5H12 CgH14 C7H,6 C10H22 Ci<jHj4
0-63 0-66 0-70 0'75 Q'85
CARBON AND THE HYDROCARBONS 359
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 MCn 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, C6H6Br,
and ethyl bromide, C2HsBr, be heated with metallic sodium, the
sodium combines with the bromine of both compounds, forming sodium
bromide, NaBr. From the first combination the group C6H5 remains,
and from the second C2H5. Having an odd number of hydrogen atoms,
they, in virtue of the law of even numbers, cannot exist alone, and there-
fore combine together forming the compound CCH5.C2H5 or C8H,0
(ethylbenzene). Hydrocarbons are also produced by the breaking up j •
of more complex organic or hydrocarbon compounds, especially by heat- I
iiiff--tlmt iat by dry dfoffilatfoj^ For instance, gum-benzoin, contains
an acid called benzoicacid, C7H6O2, the vapours of which, when passed
through a heated tube, split up into carbonic anhydride, CO2, and
benzene, C6HC. Carbon and hydrogen only unite directly in oijejtatjo
'of_ combination — namely, to form acetylene, having the composition
CgHfr which, as compared with other hydrocarbons, exhibits a very
great stability at a somewhat high temperature.29
w If, at the ordinary temperature' (assuming therefore that the water formed will be
in a liquid state) a gram molecule (26 grams) of acetylene, C2H2, 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 2x97 + 69, or 263
thousand calories produced. It is evident, then, that acetylene in its formation absorbs
810—268, 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 Cn£C2n+2: 52'8 + 158'8» thousand calories; (2) for
Cnl^n : 17'7 + 158'ln thousand calories ; (8) according to Stohmann (1888) for liquid
saturated alcohols, CnH^^C) : 11-8 + 156'8n, and as the latent heat of evaporation = about
8'2 + 0-6n, in a gaseous state, 20'0 + 156'9n ; (4) for monobasic saturated liquid acids,
CnH.jnO.j :— 95'8 + 154-8M, and as their latent heat of evaporation is about 5'0 + l'2w, in a
gaseous form, about — 90 + 155n; (5) for solid saturated bibasic acids, CnH^-gOxl
—253-8 + 152-6n, if they,are expressed as CBH.jnCjH.j04, then 5T4 + 152'6n ; (6) for bed.,
tene and its liquid homologues (still according to Stohmann) CnH2n_n i — 158'6+ 156'8n,
»nd in a gaseous form about — 155 + 1577?; (7) for the gaseous homologues of acetylene,
CnHjn-a (according to Thomsen)^— 5+157rt, It is evident Jrom the preceding figurei
860 PRINCIPLES OF CHEMISTRY
There is one fflfrgfonce known amonp the saturated hydrocarbons
composed of 1 atom of carbon and 4 atoms of hydrogen ; this is a com-
pound containing the highest percentage of hySrogen (CH4 contains
25 per cent, of h^drogspj^and at the same time it is the only hydro-
carbon whose molecule contains but a single atom of carbon. This
saturated hydrocarbon, CH4, is called marsh gas or methane.^i3.
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.29bu 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
that the group CH2, or CH3 substituted for H, on burning gives out from 152 to
159 thousand calories. This is less than that given out by C + H2, 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 CH2. According to Stoh-
mann, for dextroglucose, CgH^Oe, it is 678'7 ; for common sugar, Ci2H22OH, 1325'7 ; for
cellulose, C6H10O5, 678-0; starch, 677'5; dextrin, 666'2; glycol, C5H602, 281'7; glycerine,
897'2, &c. The heat of combustion of the following solids (determined by Stohmann) is
expressed per unit of weight : naphthalene, CioHg) 9,621 ; urea, CN2H4O, 2,465 ; white of
egg, 5,579; dry rye bread, 4,421; wheaten bread, 4,302; tallow, 9,865; butter, 9,192;
linseed oil, 9,823. The most complete collection of arithmetical data for the heats of
combustion will be found in V. P. 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 dryuess 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.
CH4, 5 vols. CO, and 5 vols. N), about 12,000 ; (7) producer gas (see next Chapter), con-
taining 2 vols. carbonic anhydride, 80 vols. carbonic oxide, and 68 vols. nitrogen for one
part by weight of the wJiole carbon burnt, 6,800, and for one part by weight of the gas,
910, units of heat ; and (8) water gas (see next chapter) containing 4 vols. carbonic anhy-
dride, 8 vols. N2, 24 vols. carbonic oxide, and 46 vols. H2, for one part by weight of the
carbon consumed in the generator 10,900, and for one part by weight of the gas, 8,600
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 ah- (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).
tt tte Manure which decomposes under the action of bacteria gives off CO > and
CH4.
CARBON AND THE HYDROCARBONS 861
accord. The gas which is evolved consists pi-incipally of marsh gas.39
If wood, coal, or many other vegetable or animal substances are decom-
posed 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
50 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
tunnel.
11 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 wilh the receiver / (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 &ndj 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, (8) liquid hydrocarbons, (4) hydrogen sul-
phide, H2S, (5) carbonic anhydride, CO,, (6) carbonic oxide, CO, (7) sulphurous anhy-
dride, SOa, but a great part of the illuminating gas consists of (8) hydrogen, (9) marsh
gas, (10) olefiant.gas, CaH4, and other gaseous hydrocarbons. The hydrocarbons (8, 9, and
10), the hydrogen, and carbonic oxide are capable of combustion, and are useful com-
ponent parts, but the carbonic anhydride, the hydrogen sulphide, a'nd sulphurous anhy-
dride, as well as the vapours of ammonium carbonate, form an injurious admixture,
because they do not burn (CO2, SO2) and lower the temperature and brilliancy of the flame,
or else, although capable of burning (for example, HaS, CSo, 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, car-
bonic 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,
FeSO4, acts still better, because the latter, with lime, Ca(HO)2, forms ferrous hydroxide,
Fe(HO)., and gypsum, CaSO*. The suboxide (partly turning into oxide) of iron absorbs
CARBON AND THE HYDROCARBONS
363
As the decomposition of the organic matter which forms coal, is still
going on underground, the evolution of large quantities of marsh gas
H2S, forming FeS and H2O, 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 w,. 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, BO 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 arrange-
ment. 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 defiant gas, CoH,, is
fourteen times, and the vapours of benzene thirty-nine
times; heavier than hydrogen, and illuminating gas some-
times contains 15 p.c. of its volume of them. The brilliancy
of the flame of the gas increases with the quantity of
defiant 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 85 to 60 p.c. of its volume of marsh gas, from
80 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 8 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 con-
tains a great deal of carbonic anhydride, although on the
other hand there is an almost complete absence of sul-
phur compounds. Tar, oils, naphtha, and such materials
furnish a large qnantity 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 Ibs. 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. .
shaped mouthpiece, and
escapes in a Bne stream
from the platinum jet placed
at t|,e extremity of the side
tube.
364
PRINCIPLES OF CHEMISTRY
frequently occurs in coal-mines.33 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
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 con-
same 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 mouth-
piece ; or, in other forms, gas is passed through the blowpipe ; when a large, hot, smoke-
less 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
watei gas, which will be treated of in the following chapter), because complete combus-
tion may be effected without an access of air. It ia
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.
52 The gas which is set free in coal mines
contains a good deal of nitrogen, some carbonic
anhydride, and a large quantity of marsh gaa.
The best means of avoiding an explosion consists
in efficient ventilation. It is best to light coal
mines with electric lamps.
53 The Davy lamp, of which an unproved form
is represented in the accompanying figure, is used
for lighting coal and other mines where combus-
tible gas is found. The wick of the lamp is en-
closed in a thick glass cylinder which is firmly
held hi a metallic holder. Over this a metallic
cylinder and the wire gauze are placed. The pro-
ducts of combustion pass through the gauze, and
the air enters through the space between the
cylinder and the wire gauze. To ensure greater
"k*? the kmP cannot ** OP6"6* without extin-
guishing the flame.
94 In Pennsylvania (beyond the Alleghany mountains) many of the shafts flunk tor
petroleum only emitted gas, but many useful applications for it were found and it was
CARBON AND THE HYDROCARBONS 365
Tolerably pure marsh gas 38 may be obtained by heating a mixture of an f I
acetate with an alkali. Acetic acid, C2H402, on being heated is decom-
posed into marsh gas and carbonic anhydride, C2H402=CH4 + C08.
An alkali — for instance, NaHO — gives with acetic acid a salt,
C2H3NaO2, which on decomposition retains carbonic anhydride, form-
ing a carbonate, Na2C03, and marsh gas is given oft' :
C2H3NaO2 + NaHO=Na2C03 + CH4
Marsh_gas_isdifficult 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 CnH2n+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 hydro-
carbons ought to be derived from methane, CH4, 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 mole-
cule of methane, CH4, is capable of undergoing substitution in the four
following ways : — (1) Methyl substitution, when the radicle, equivalent
to hydrogen, called methyl CH3, replaces hydrogen. In CH4 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 H2 and CH2 (this radicle is called methylene),
is founded on a similar division of the molecule CH4 into two equiva>
conducted in metallic pipes to works hundreds of miles distant, principally for metal*
Inrgtcal purposes.
55 The purest gas is prepared by mixing the liquid substance called zino methyl,
Zn(CH3),2, with water, when the following reaction occurs :
Zh(CH3)2 + 2HOH = Zn(HO)2 + 2CHSH.
M Methylene, CH2, does not exist. When attempts are made to obtain it (for
example, by removing X3 from CH2X2), CaH, or CSH0 are produced— that is to say, it
undergoes polymerisation.
966 PRINCIPLES OF CHEMISTRY
lent .parts, H2 and CH2 ; (3) acetylene substitution, or the exchange
between CH on the one hand and H3 on the other ; and (4) carbon
substitution — that is, the substitution of H4 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 CH4 they arelikewise 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 re-
placed by CH3 there is an increase of CH2 ; when H2 is replaced by
CH2 there is no increase of hydrogen ; in the acetylene substitution
CH replaces H3,. therefore there is an increase of C and a decrease of
H2 ; in the carbon substitution there is a decrease of H4. 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 intro-
duced by the methyl substitution, since it leads to the addition
of CH2 ; starting from CH4 we obtain C2H6, C3H8, and in general,
CnH2n+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, CH4,
it is evident that the saturated hydrocarbon formed is C2H6 or
(CH3) (CH3).37 This is called ethane. By means of the methylene
substitution alone, ethylehe, C2H4, or (CH2) (CH2) may be directly
obtained from CH4, and by the acetylene substitution C2H2 or
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, CH4, on the one hand the substi-
tution of chlorine or iodine, CH3C1, CH3I, for the hydrogen -may be effected, and on the
other hand such metals as sodium may be substituted for the hydrogen, e.g. CHjNa.
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 :
CH3C1 + CH5Na = NaCl + C2H«.
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.
CARBON AND THE HYDROCARBONS 367
(CH) (CH), or acetylene, both the latter being unsaturated hydro-
carbons. Thus we have all the possible hydrocarbons with two atoms
of carbon in the molecule, C2H6, ethane, C2H4, ethylene, and C2H2,
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 C2H6 will
still contain hydrogen when C replaces H4) and therefore further sub-
stitutions 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
CnH2/1+2, but an unlimited number of unsaturated hydrocarbons, CHH£,,
Q.H2,,_2 .... CHH2 („_„), where m varies from 0 to n— I,88 and
where n increases indefinitely. From these facts not only does th«
existence of a multitude of polymeric hydrocarbons, differing in mole-
cular weight, become intelligible, but it is also seen that there is a possi-
bility of cases of isomerism with the same molecular weight. This
polymerism. so common to hydrocarbon compounds is already apparent
in the first unsaturated series CnH2H, because all the terms of this
series C2H4, C3H6, C4H8 .... C30H60 .... have one and the
same composition CH2, 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 preparation4.0 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, C6H6, and cinnamene, C8H8,
correspond with the composition of acetylene or to a compound of
the composition CH.41 The first boils at 81°, the second at 144°;
58 When m = n— 1, we have the series C«H2. The lowest member is acetylene, CaH3.
These are hydrocarbons containing a minimum amount of hydrogen.
59 For instance, ethylene, C2H4, combines with Br2, HI, HjS04, as a whole molecule,
as also does amylene, C5H10, and, in general, C«Hon.
•*° For instance, ethylene is obtained by removing the water from ethyl alcohol,
C2H5(OH), and amylene, C4H10, from amyl alcohol, C5H, )(OH), or in general CMH2,,, from
C«H2»+1(OH).
41 Acetylene and its polymerides have an empirical composition CH, ethylene and
its homologues (and polymerides) CH2, ethane CH3, methane CH4. 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
CjoHg} and C^H^, 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
*4
868 PRINCIPLES OF CHEMISTRY
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.
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 C2H6 and C3H8, because the former is
CH4, 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,
CH3CH3 ; 43 the same argument also applies in the case of pro-
pane, CH3CH2CH3, where one compound only can be imagined. It
can be distinguished beyond a doubt. Without Dalton's law chemistry could not have
been brought to its present condition, Ibut it cannot alone express 11 those grada-
tions which are quite clearly understood and predicted by the law of Avogadro-
Qerbardt. "
43 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 C2H5, methyl CH3, phenyl CeH5, <fec.) ; 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 com-
plex hydrocarbon compounds. Then Kekuleand 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 dis-
cover 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,
bat the details of the still imperfect knowledge in relation to this matter must be sought
for in special works devoted to organic chemistry.
<5 Direct experiment shows that however CHSX is prepared (where X = for instance
Cl, &c.) it is always one and the same substance. If, for example, in CX4, X is gradually
replaced by hydrogen until CH3X is produced, or in CH^, the hydrogen by various
CARBON AND THE HYDROCARBONS 869
is to be expected, however, that there should be two butanes, C4HIO,
and this is actually the case. In one, methyl may be considered
as replacing the hydrogen of one of the methyls, CH3CH2CH4CH3 ;
and in the other CH3 may be considered as substituted for H in
PFT
CH2, and there it will consist of CH3CH^g3. The latter may
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. C2H4 and CH2CH, are, it is evident,
identical ; but there ought to be, and are, two hydrocarbons of the
composition C3H6, propylene and trimethylene ; the first is ethylene,
CH2CH2, in which one atom of hydrogen is exchanged for methyl,
CH2CHCH3, and trimethylene is ethane, CH3CH3, with the substi-
tution of methylene for two hydrogen atoms from two methyl groups —
CH
that is, £jr2CH2,44 where the methylene introduced is united to both
the atoms of carbon in CH3CH3. 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 c 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 derivatires
of ethane, C2H6 : if two atoms of the hydrogen be exchanged for X2,
means is replaced by X, or els«, (or instance, if CH5X 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 funda-
mental 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 CHSCH)CH, or CH3CH2X), then there would be as many different forms of CHSX
as there were diversities in the atoms of hydrogen in CH.,. The scope of this work does,
not permit of a more detailed account of this matter. It is given in works on organio
chemistry.
44 The union of carbon atoms in closed chains or rings was first suggested by KekuU
as an explanation of the structure and isomeriem of the derivatives of benzene, C6H«,
forming aromatic compounds (Note 26).
870 PRINCIPLES OF CHEMISTRY
one will have the ethylene structure, CH2XCH2X, and the other an
ethylidene structure, CH3CHX2 ; such are, for instance, ethylene
chloride, CH2C1CH2C1, and ethylidene chloride, CH3CHC12. And as
in the place of the first atom of hydrogen not only metals may be substi-
tuted, but Cl, Br, T, OH (the water radicle), NH2 (the ammonia radi-
cle), NO2 (the radicle of nitric acid), <fcc., so also in exchange for two
atoms of hydrogeti 0, 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 descrip-
tion 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, CH2CH2>and
acetylene, CHCH, and a short acquaintance with petroleum as the natu-
ral source of a mass of hydrocarbons. Etliylene, or olefiant gas, C2HO
45 The following are the most generally known of the oxygenised but non-nitro-
genous hydocarbon derivatives. (1) The alcohols. These are hydrocarbons in which
hydrogen is exchanged for hydroxyl (OH). The simplest of these is methyl alcohol,
CH5(OH), or wood spirit obtained by the dry distillation of wood. The common spirits
of wine or ethyl alcohol, C2H5(OH), and glycol, C2H4(OH)2, correspond with ethane.
Normal propyl alcohol, CH3CH2CH2(OH), and isopropyl alcohol, CH3CH(OH)CH3, pro-
pylene-glycol, C3H6(OH)2, and glycerol, C3H5(OH)3 (which, with stearic and other acids,
forms fatty substances), correspond with propane, C3H3. All alcohols are capable of form-
ing water and ethereal salts with acids, just as alkalis form ordinary salts. (2) Aldehydes
are alcohols minus hydrogen ; for instance, acetaldehyde, CSH4O, corresponds with ethyl
alcohol. (8) It is simplest to regard organic acids as hydrocarbons in which hydrogen
has been exchanged for carboxyl (CO2H), 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 C2H0, and has the constitution C2H4(OH) (CO2H).
If to these products we add the haloid salts (where H is replaced by Cl, Br, I), the nitvo-
compounds containing NO2 in place of H, the amides, cyanides, ketones, and other com-
pounds, 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 com-
position of plants and animals.
CARBON AND THE HYDROCARBONS 871
is the lowest known member of the unsaturated hydrocarbon series of
the composition C,, H2n. 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, olenant gas splits up, first into acetylene and methane
(3C2H4 = 2C2H2 + 2CH4, Lewes, 1894), and at a higher tempera-
ture into carbon and hydrogen j 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 olenant 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.
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 ele-
ments of water from the alcohol, C2H5(OH), and gives olefiant gas ;
C2H6O = H2O + C2H4. 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 4 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
411 Ethylene bromide, CgH^Brg, when gently heated in alcoholic solution with finely
divided zinc, yields pure ethylene, the zinc merely taking up the bromine (Subaneyeff).
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, C2H4 = CH4 + C = 2C + 2H2.
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, C2H4 + SO2 = 2COj + 2H2O. 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.
4
872 PRINCIPLES OF CHEMISTRY
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 oletiant
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 oletiant gas, little by little, to combine with the sulphuric
acid, forming C2H4H2SO4. If, after this absorption, the sulphuric acid
be diluted with water and distilled, alcohol separates, which is produced
in this case by the oletiant gas combining with the elements of water,
C2H4 -H H2O = C2H6O. 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,
X2, ethylene gives saturated compounds, C2H4X2 or CH2XCH2X
(for example, C2H4C12), which correspond with ethane, CH3CH3 or
r w <8
C2I16.
Acetylene, C2H2 = CHCH, is a gas ; it was first prepared by Ber-
thelot (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 C2H2.48 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 C2H4Br2 ; from this the
hydrobromic acid is removed by means of an alcoholic solution of
caustic potash, leaving the volatile product C2H3Br ; 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,,H2,1_1Br, forming C,,H2,,_2.
48 The homologues of ethylene, CnH2,i, are also capable of direct combination with
halogens, &c., but with various degrees of facility. The composition of these homologues
can be expressed thus : (CH3)z(CH2)v (CH)jCr, 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 8x + e, whence « + 2r = « ; by
this means the possible isomerides are determined. For example, for butylenes, C4H8,
<CHS)2(CH)2, (CH3)2(CH2)C, (CH2) (CH2)8CH, and (CH2)4 are possible.
48 bii ge(i also method of preparing CjHj in Note 12 bis.
CARBON AND THE HYDROCARBONS 873
Acetylene is also produced in all those cases where organic sub-
etances 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 Same 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.
Acetylene, being further removed than ethylene frotn the limit
CBH2l,+2 °f hydrocarbon compounds, has a still greater faculty of combi-
nation 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 I2, HI, H2SO4, C12, Br2,
&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. lu this manner pure
acetylene may be obtained. Acetylene and its homologues also
readily react with corrosive sublimate, HgCl2 (Koucheroff, Favorsky).
Acetylene burns with a very brilliant flame, which is accounted for
by the comparatively large amount of carbon it contains.50
The formation and existence in nature of large masses of petroleum
or a mixture of liquid hydrocarbons, principally of the series C,,H2n+4
and C,,H2n is in many respects remarkable.51 In some mountainous
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.
40 Amongst the homologues of acetylene CnHZ(;-j, the lowest is CSH4; allylene,
OHjCCH, and allene, CH4CCH2, are known, but the closed structure, CHS(CH)2, is
little investigated.
41 The saturated hydrocarbons predominate in American petroleum, especially in
its more volatile parts ; in Baku naphtha the hydrocarbons of the composition CnH2» form
the main part (Lisenko, Markovnikoft, Beilstein) but doubtless (Mendelueff) it also con-
tains saturated ones, C,,H.,,, + 2. The structure of the naphtha hydrocarbons is only known
for the lower homologues, but doubtless the distinction between the hydrocarbons of the
Pennsylvania!! and Baku naphthas, boiling at the same temperature (after the requisite.
r874 PRINCIPLES OB CHEMISTRY
districts — as, for instance, by the slopes 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 O8 and the heavy
kinds up to 0'98. The former are very mobile liquids, and more vola-
tile ; 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, <fec.),
from which the hydrocarbons whose boiling points start from 0° may
be extracted — namely, the hydrocarbons C4H10, C5H,2 (which boils
at 30°), C6H14 (boils at 62°), C7H16 (boils about 90°), &c. Those
fractions of the naphtha distillate which boil above 130°, and contain
hydrocarbons with C9, C10, Cn, «fec., enter into the composition of the
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,,H._,,,, 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 with-
out doubt caused by the pressure or tension of the coihbustrble 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.
CARBON AND THE HYDROCARBONS 375
oily substance, universally used for lighting, called kerosene or photo-
gen 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 O75, 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, tfec. Those por-
tions 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 tire), 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
55 This is asp-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 80 p.c) of intermediate oil, is produced in enormous quantities, about
800 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 Kailway)
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 (spe-
cific 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 80°).
54 The substitution of Baku pyronaphtha, or intermediate oil, or Baku oil (see Note
68), 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 24s. (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
4s. (1 rouble 80 copecks), and with a pipe line to the shore of the Black Sea woulu not
cost more than 8 roubles, or 16s. Moreover, a 'ton of Baku naphtha will yield as much as
two-thirds of a ton of kerosene, Baku oil, and pyronaphtha suitable or illuminating
purposes.
876 PRINCIPLES OF CHEMISTRY
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/"'6 Much more probably its formation may
be attributed to the action of water penetrating through the crevasses
formed on the mountain slopes 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
44 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 86 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. Mendele'eff) : (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.
w As during the process of the dry distillation of wood, seaweed, and similar vege-
table dtbris, 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 Tegetable debris 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 dfbris 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 b«en 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
ft 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 ; (8) 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, Taman, Groznoe, Dagestan), I
Seveloped in 1876 the hypothesis of the mineral origin of naphtha expounded further on.
CARBON AND THE HYDROCARBONS
mountain chains,57 because water with iron carbide ought to give iron
oxide and hydrocarbons.58 Direct experiment proves that the so-called
spiegehisen (manganiferous iron, rich in chemically combined carbon)
when treated with acids gives liquid hydrocarbons59 which in com-
*7 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
6'5. As at the surface water and all rocks (sand, clay, limestone, granite, &c.) have a
density less than 8, 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 con-
siderably 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 ter-
restrial ones, often contain either dense masses of iron (for example, the Pallosovo iron
preserved in the St. Petersburg Academy of Sciences) or granular musses (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 wheu
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
oxj'gen 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.
48 The following is the typical equation for this formation :
8FemC,, + 4,,,H2O = ™Fe3O4 (magnetic oxide) + C^H^m (see Chapter XVII., Note 88).
59 Cloez investigated the hydrocarbons formed when cast-iron is dissolved in hydro-
chloric acid, and found C«H2« 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.
•w.
378 PRINCIPLES OF CHEMISTEY
position, appearance, and properties are completely identical with
naphtha.60
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 'Mid 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) j
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 dis-
tricts are full of such volcanoes, which from time to time are in a state of eruption. In
old naphtha beds (such as the Pennsylvania!}) 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 im-
pregnated 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 geo-
logical 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,.,
euch 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 con-
stantly 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 Pennsylvania^! 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 production1 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 ia used for the production of paraffin end ceresene,
tot. 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
cf practical men.
879
CHAPTER IX
COMPOUNDS OP CARBON WITH OXYGEN AND NITROGfiH
CARBONIC anhydride (or carbonic acid or carbon dioxide, CO2) 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 i* pro-
duced by the combustion of charcoal, and that in some cases it is found
Fit;. 61. — Dumas' and Stas' apparatus for determining tbc composition of carbonic anhydride.
Carbon, graphite, or a diamond is placed in the tube F 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 conrerted into carbonic acid by passing tb«
products of combustion through a tube F, containing cupric oxide heated in a furnace. The cuprio
oxide oxidises this CO into 00,, 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 n, I, J, K before and after the experiment), tb»
composition of carbonic anhydride and the equi Talent of carbon may be determined.
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 deter-
mined its composition, and showed that it only contains oxygen and
carbon. Berzelius, Dumas with Stas, and Roscoe, determined its com*
8'&0
PRINCIPLES OF CHEMISTRY
position, 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 thatduring the combustionof 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. O2 occupies
two volumes and is a molecule of ordinary oxygen ; CO2 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
eome caves. The well-known Dog grotto, near Agnano on the bay of
Baise, near Naples, furnishes the best known example of such an evolu-
tion. 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 vol-
canoes, 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, Spriidel 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 anhy-
dride 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 be found to be many grams per
day for a mouse. Such expeiiments on the respiration of animals have
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 881
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,
CaCO3, 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.
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 sepa-
rated 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 daring 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
\a 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. There-
fore 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 wonls, 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 mattei 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 libe-
rated 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 an-
hydride 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 com-
position C6H1005 being representatives of them. They may be considered like all carbo-
hydrates as compounds of carbon and' water, 6C + 5H2O. 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.
382
PRINCIPLES OF CHEMISTRY
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 CaCO3, 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 CO2 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
FIG. 62.— Apparatus for the combustion of organic substances by igniting tlurn with
oxide of copper.
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.) con-
taining 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.
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,
HC1 — because, in the first place, the substance formed, calcium
chloride. CaClj, is soluble in water and does not hinder the further
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 883
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
CaC03 + 2HC1 = CaCl2 + H2O + CO2.
The nature of the reaction in this case is the same as in the decom-
position 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, H2C03, 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 j
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 oji its salts, and this reaction takes
place at ordinary temperatures.4
* 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 ba used ;
for instance, washing-soda, Na/>CO3, 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 solu-
tions, the glucose C6H12O6 changes under the influence of the yeast organism, forming
alcohol (2C3H60), and carbonic anhydride (2CO2) which separates in the form of gas ; it
the fermentation proceeds in closed bottles sparkling wine is obtained. When carbonio
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 perman-
ganate, which absorbs organic matter and does not take up carbonic anhydride.
5 Hypochloroue acid, HC1O, and its anhydride, C12O, 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
884 PRINCIPLES OF CHEMISTRY
For the preparation of carbonic anhydride in laboratories, marble is
generally used. It is placed in a Woulfe's bottle and treated with hydro-
chloric 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
Carbonic anhydride may also be prepared by heating many of the
salts of carbonic acid ; for instance, by heating magnesium carbonate,
MgC03 (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, NaHC03, 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
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' I
Ij acid, nor even sulphuric acid nor acetic acid, acts on limestone except in presence of I
i -water. We shall refer to this later on.
J! 5 The direct observations made (1876) by Me.ssrs. Bogouski and Kayander lead to the'
""" conclusion that the quantity of carbonic anhydride evolved by the action of acids on i
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 tho 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 0J02 gram of carbonic anhydride will be
evolved. If the litre contains n grams of hydrochloric acid, then by experiment the
amount will be n x 0'02 of carbonic anhydride. Therefore, if the litre contains 86'5
(=HC1) grams, about 0'78 gram of carbomc anhydride (about half a litre) would be
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 ( = HNO3) grams of nitric acid, or 81
(=HBr) grams of hydrobromic acid, the quantity of carbonic anhydride evolved will still
be 0'78 gram. Spring, in 1890, made a series of similar determinations,
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 carbomc anhy- '
dride 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 com- j
mences, as this latter is heavier than the soap-bubbles filled with air. Naturally, after a ;
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 885
It is an example of those gaseous substances which have been
long ago transformed into all the" three states. In order €b obtain
liquid carbonic anhydride, the gas must be submitted to a pres-
sure of thirty-six atmospheres at 0°.7 Its absolute boiling point
ss + 320.8 Liquid carbonic anhydride is colourless, does not mix with
water, but is soluble in alcohol, ether, and oils ; at 0° its specific gravity
is 0-83.8bis 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 Ws
The capacity which carbonic anhydride has of being liquefied stands
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 im-
proved by Thilorier and Natterer, whose apparatus is given in Chapter VI. in describ-
ing N2O. 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
eolid 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 trans-
ported and preserved safely for a long time. It is used, for instance, in breweries.
9 bli 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 pro-
duces a severe frost bite similar to a burn. If the snowlike solid be mixed with ether, &
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
I
386 PRINCIPLES OF CHEMISTRY
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
atmospfieres, are now prepared artificially, because water saturated
with this gas promotes digestion and quenches thirst. For this pur-
pose 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 effer-
vescing drinks and artificially effervescing wines are thus prepared.
The presence of carbonic anhydride has an important significance
iu 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, tfcc., from them. This process
has been going on and continues on an enormous scale. Rocks
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.)
Bleekrodc, by compressing solid CO2 in a cylinder by means of a piston, obtained a
semi-transparent stick, which contained as much as 1'S and even 1*6 gram of CO2 per
cubic centimetre. In this form the CO2 slowly evaporated, and could be kept for a long
time.
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 oaves
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 im-
portant part in the nourishment of plants, because all plants contain both lime and
phosphoric acid.
COMPOUNDS OF CARBON WITH OXYGEN AND NITKOGEN 387
contain silica and the oxides of various metals ; amongst others,1
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 car-
bonic anhydride in solution in water is essential to the nourishment
and growth of water plants.
Although carbonic anhydride is soluble in water, yet no definite
hydrate is formed ; ' ' 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
K2CO3, Na2C03, HNaCOg, <fec., therefore carbonic acid ought to have
the composition H2C03 — that is, it ought to contain C02 + H2O.
Whenever this substance is formed, it decomposes into its component
parts — that is, into water and carbonic anhydride. The acid properties
of carbonic anhydride11 bls are demonstrated by its being directly ab-.
sorbed by alkaline solutions and forming salts with them. In distinction
from nitric, HNO3, 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 (NaNO3, AgN03), and with bivalent 12 metals (such as calcium,
barium, lead) salts containing two acid groups— for example, Ca(NO3)2,
Pb(NO3)2 —carbonic acid, H2CO3, in bibasic, that is contains two atoms
of hydrogen in the hydrate or two atoms of univalent metals in their
salts : for example, Na2CO3 is washing soda, a normal salt ; NaHCO3 is
the bicarbonate, an acid salt. Therefore, if M' be a univalent metal,
its carbonates in general are the normal carbonate M'2C03 and the
11 The crystalluhydrate, COa,8H,O of Wroblewski (Chapter I., 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
HjCOj which should occur, judging from the composition of the salts.
11 bu 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.
11 The reasons for distinguishing the uni-, bi-, tri-, and quadri-valent metals will be
explained hereafter on passing from the univalent metals (Na, K, Li) to the bivalent
(Mg. Ca, Ba), Chapter XIV
388 PRINCIPLES OF CHEMISTRY
tacid carbonate, M'HCO3 ; or if M" -be a bivalent metal (replacing H2) its
normal carbonate will be M''CO3 ; these metals do not usually form acid
i salts, as we shall see further on. The bibasic character of carbonic acid
is akin to that of sulphuric acid, H2SO4,13 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 con-
sidered a weak acid. This conception must, however, be taken as only
comparative, as up to this time theie is no definitely established rule for
measuring the energy u of acids. The feeble acid properties of carbonic
13 Up to the year 1840, or thereabout, acids were not distinguished by their basicity.
Graham, while studying phosphoric acid, HjPOj, and Liebig, while studying many organic
acids, distinguished mono-, bi-, and tri-basic 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, KHO, &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, NaHCOs, ^d normal salts, NaoCOj, it is evident that
the hydrate is H._,CO3, 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 j expressed the composition, but not the molecular weight, of carbonic anhydride ; and
the composition of the normal salt would be Na.jC.jOe or NaCO^, therefore carbonic acid
might have been considered as a monobasic acid. Then the acid salt would have been
represented by NaCOj,HCO5. Such questions were the cause of much argument and dif-
ference 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).j as containing the elements
necessary for the separation of the molecule of water, H2O. Thus H2SO4 or SO2(OH)j,
H2CO3, or CO(OH)2, and other bibasic acids decompose into an anhydride, BO, and water,
H20. But as nitrous, HXOo. iodic, HIO3, hypochlorous, HC1O, and other monobasic
acids easily give their anhydrides NaO3, I2O5, C12O, &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 HjSO j, on being evaporated and dis-
tilled, gives SO3 + H2O), and the decomposition of acids into water and anhydride pro-
ceeds particularly easily in dealing •with feebly energetic acids, such as carbonic, nitrous,
boric, and hypochlorous. Let ns add that carbonic acid, as a hydrate corresponding to
marsh gas, C(HO)4 = CO2 + 2H20, ought to be tetrabasic. But in general it does not form
such salts. Basic salts, however, such as CuCO-CuO, may be regarded in this sense,
for CCu2O4 corresponds with CH4O4) as Cu corresponds with Hg. Amongst the
ethereal salts (alcoholic derivatives) of carbonic acid corresponding cases are, however,
observed; for instance, ethylic orthocarbonate, C(C2HjO)4 (obtained by the action of
chloropicrin, C(NO2)Cls, on sodium ethoxide, C2HsONa; boiling point 158°; specific
gravity, Q'921. The name orthocarbonic acid lot CH^ is taken from ortnophosphorio
acid, PHjOf, which corresponds with PH3 (tee Chapter on Phosphorus).
14 Long ago endeavours were made to find a measure of affinity of acids and bases,
because some of the acids, each as sulphuric or nitric, form comparatively stable salts, de-
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 889
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
composed 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^O, sodium, NajO, 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.jOj, or bismuth oxide,
Bi.^Ox, 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
l>owerful 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 criticism?, although it com-
prises many interesting features. The fact is that Thomson, 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 accom-
j>anying 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 NX2, and Y is npt, 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 demon-
strates 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 Arrheniua
(1884), and applied on an extensive scale by Ostwald (who developed it in great detail in
bis Lehrliucli d. allgemeinen Chemie, v. ii., 1887), is founded on the fact that' the re-
lation of the so-called molecular electrical-conductivity of weak solutions of variou? 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
fate 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.
390 PRINCIPLES OF CHEMISTRY
many cases themselves acting as al kalis. 1A The acid salts of these alkalis,
NaHCO3 and KHCO3, have a neutral reaction on litmus, although they,
like acids, contain hydrogen, which may be exchanged for metals. The
ncid salts of such acids — as, for instance, of sulphuric acid, NaHS04 —
have a c'early defined acid reaction, and therefore carbonic acid is un-
able to neutralise the powerful basic properties of su^h alkalis as potash
or soda. Carbonic acid does not even combine at all with feeble bases,
such as alumina, A12O3, and therefore if a strong solution of sodium
carbonate, Na2CO3, be added to a strong solution of aluminium sulphate,
A12(S04)3, although according to double saline decompositions alu-
minium carbonate, A12(CO3)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 : A]2(C03)3 •+- 3H20= A12(OH)G
+ 3C02. Thus feeble bases are unable to retain carbonic ecid even at
ordinary temperatures. For^the same reason, in the case of bases of
medium energy, although they form carbonates, the latter are compara-
tively easily decomposed by heating, as is shown by the decomposition of
copper carbonate, CuCO3 (see Introduction), and even of calcium carbo-
the annexed figures, in which the energy of hydrochloric acid is taken as equal to
100:—
I II III
Hydrochloric acid, HC1 . . 100 100 100
Hydrobromic acid, HBr , . 101 98 105
Nitric acid, HNO3 . . 100 100 96
Sulphuric acid, H3S04 65 49 74
Formic acid, CHa6.2 2 4 1
Acetic acid, C.jirLjO.; 1 2 1
Oxalic acid, C.,H.,04 20 24 18
Phosphoric acid, PH3O4 7 6
The coincidence of these figures, obtained by so many various methods, presents a
jnosl*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 HKo consideration. On this account, until the theory of solution
is nfore 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^COj. 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.
COMPOUNDS OF CABBON WITH OXYGEN AND NITROGEN 891
nate, CaC03. 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, NaHCO3 — decompose
even on heating their solutions (2NaHC03 SB NajC03 + H2O + COa),
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 satu-
rated (up to the formation of a normal salt) with sulphuric or nitrio
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"X8, insoluble pre-
cipitates of carbonates, M2CO3 or M"CO3. Thus a solution of barium
chloride gives with sodium carbonate a precipitate of barium carbonate,
BaCO3. For this reason rocks, especially those of aqueous origin, very
often contain carbonates ; for example, calcium, ferrous, or magnesium
carbonates, &c.
Carbonic anhydride — which, like water, is formed with the develop-
ment of a large amount of heat — is very stable. Only very few sub-
stances 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 ; COa + H2 = CO + HaO.
16 Although carbonic acid is reckoned among the feeble acids, yet there are evi-
dently 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 sub-
stances. Thus, in solution, sodium silicate, NajSiOj, with carbonic anhydride will to some
extent form sodium carbonate and silica, SiO.> ; 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 car-
bonic 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 (sec Chapter X.)
*5
892 PRINCIPLES OF CHEMISTRY
But only a portion of the carbonic acid gas undergoes this change, wn^
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 ex-
ceedingly stable, still on being heated it partially decomposes into car*
bonic oxide and oxygen. Deville showed that such is the case if car-
bonic 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 C02 give two volumes of CO and
one volume of 0. 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 th^
carbonic anhydride be removed, then the mixture explodes when a.
spark is passed and forms carbonic anhydride.-17 bls If from an identical
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, be-
cause during the combustion of hydrocarbons the hydrogen burns first. Some1 idea of
this similarity of affinity may be formed by the quantity of heat evolved. Gaseous
hydrogen, H», on combining with an atom of oxygen, O = 16, develops 69,000 heat-unita
if the water formed be condensed to a liquid state. If the water remains in the form of
K 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 Oj = 82 develops about
97,000 calories, forming gaseous C02. If it were gaseous like hydrogen, and only
contained Cn 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 — thai
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 oi
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 wiB
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 know*
ledge.
17 bis The degree or relative magnitude of the dissociation of C05 varies with the
temperature and pressure — that is, it increases with the temperature and as the pressure
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 89$
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 absorp-
tion 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 sub-
stances composing the mixture. This is one of the instances of the
influence of mass.
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 in-
termediate 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 anhy-
dride (in many ways resembling carbonic anhydride) under the action
of light (and also of heat) forms sulphur and sulphuric anhydride, S08,
and in the presence of water, sulphuric acid. But no similar decompo-
sition has been obtained directly with carbonic anhydride, although it
forms an exceedingly easily decomposable higher oxide— percarbonic
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.. is decomposed when the tempera-
ture 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.; will be
considerable even at comparatively moderate temperatures, but at the temperature ol
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
CO2 (i.e. the amount of heat required to raise 44 units of weight of CO.; 1°), according
to the determinations and calculations of Mallard and Le Chatelier, for a constant
volume Cv = 6'26 + 0'0087< ; for a constant pressure Cp = C» +2 (see Chapter XIV., Note 7),
i.e. the specific heat of CO? 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.
894 PRINCIPLES OF CHEMISTRY
acid ; I8 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, CH2O,
and from it organic acids, and that these latter in their final trans-
formation form all the other complex organic substances of the plants.
Many organic acids are found in plants in considerable quantity ; for
instance, tartaric acid, C4H6O6, found in grape-juice and in the acid
juice of many plants ; malic acid, C4HCO6, found not only in unripe
apples but in still larger quantities in mountain ash berries ; citric
acid, C6H8O7, found in the acid juice of lemons, in gooseberries,
cranberries, &c. ; oxalic acid, C^B.tO4, 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, C4H5KO6. In sorrel we find the
so-called salts of sorrel, or acid potassium oxalate, C2HKO4. 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, C2H4O2) the
acid of vinegar, be passed in the form of vapour through a heated tube,
18 Percarbonic acid, H2CO4 ( = H2CO3 + O) is supposed by A. Bach (1898) 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: 8H2C05 = 2H.2CO4 + CH;>O. This
substance CH2O 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 CO2. 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, CO3.
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 (Schbne,
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 H2O.2.
Besides which Bach (1894) showed that, in the first place, traces of formic aldehyde
and oxidising agents (CO3 or HoCy are formed under the simultaneous action of CO2
and sunlight upon a solution containing a salt of uranium (which is oxidised), and diethyl-
aniline (which reacts with CH2O), and secondly, that by subjecting BaO2) shaken up tn
water, to the action of a stream of CO., in the cold, extracting (also in the cold) with
ether, and then adding an alcoholic solution of NaHO, crystalline plates of a sodium
ealt may be obtained, which with water1' 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.
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEM 895
it splits up into carbonic anhydride and marsh gas = CO2 + CH4. 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 CH3Na is obtained,
this directly absorbs carbonic anhydride, forming a salt of acetic acid,
CH3Na + CO2 = C2H3NaO2 ; 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. CH4 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 rela-
tion between carbonic acid, H2CO3, 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 rela-
tion 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 ACO2 will be a monobasic
organic acid, A2CO2 a bibasic, A3CO2 a tribasic, and so on — that is,
each molecule of CO2 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 HCO2, or carboxyl. If the addition of CO2 raises
the basicity, the removal of CO2 lowers it. Thus from the bibasic
ox'alic acid, C2H2O4. or phthalic acid, C8H6O4, by eliminating CO,
(easily effected experimentally) we obtain the monobasic formic acid,
CH2O2, or benzoic acid, C7H6O2, 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 CO2 is CH4 with the exchange of H4 for O2, and that
the hydrate of carbonic anhydride, H2CO3, is CO(OH)2, that is,
methane, in which two parts of hydrogen are replaced by two parts of
the water radicle (OH, hydroxyl) and the other two by oxygen.
Therefore the group CO(OH), or carboxyl, HCO2, is a part of carbonic
acid, and is equivalent to (OH), and therefore also to H. That is, it
896 PRINCIPLES OF CHEMISTRY
;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 con-
tains 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.19 bis If in the saturated hydrocarbons, CnH2n+2, one part
of hydrogen is replaced by carboxyl, the monobasic saturated (or fatty)
acids, CnH2n+1(CO2H), will be obtained, as, for instance, formic acid,
HCO2H, acetic acid, CH2CO2H, . . . stearic acid, C17H35 COjH, &c.
The double substitution will give bibasic acids, CnH2n(CO2H)(C02H) ;
for instance, oxalic acid n = 0, malonic acid n = 1, succinic acid n = 2,
<fcc. To benzene, C6H0, correspond berxzoic acid, C6H,,(CO2H), phthalic
auid (and its isomerides), C0H4(CO2H)2, up to inellitic acid, C6(CO2H)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.
Carbonic Oxide. — This gas is formed whenever the combustion of
organic substances takes place in the presence of a large excess of
is bis jf CO., ig the anhydride of a bibasic acid, and carboxyl corresponds with it, re-
placing the hydrogen of hydrocarbons, and giving them the character of comparatively
feeble acids, then SOs is the anhydride of an energetic bibasic acid, and sulphoxyl,
SO2(OH), corresponds with it, being capable of replacing the hydrogen of hydrocarbons,
and forming comparatively energetic sulphur oxy acids (sulphonic acids); for instance,
CBHr,(COOH), benzoic acid, and C6H0(SO.OH), benzenosulphonic acid, are derived from
C6H6. As the exchange of H for methyl, CHj, is equivalent to the addition of CH.>, the.
exchange of carboxyl, COOH, is equivalent to the addition of CO2 ; so the exchange of H
for sulphoxyl is equivalent to the addition of SOj. The latter proceeds directly, for
instance: C6H6 + SO3=C6H5(SO2OH).
As, according to the determinations of Thomsen, the heat of combustion of the vapours
of acids RCO> is known where R is a hydrocarbon, and the heat of combustion of the
hydrocarbons R themselves, it may be S3en that the formation of acids, RCO2, from
R + CO2 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= H2 CH4 CoH§ CgHg
68-4 212 870 777
RCOo= 69-4 225 887 76G
•Thus H2 corresponds with formic acid, CH2O2 ; benzene, C6H6, with benzoic acid, CjHgOj.
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.
COMPOUNDS OF CAEBON WITH OXYGEN AUD NITBOGEN 397
'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, C0.2 + C = 2CO. By this reaction
carbonic oxide is prepared by passing carbonic anhydride through char-
coal 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, com
bustion 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 char-
19 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 ab-
sorb heat) are not as capable of giving a
high temperature in ordinary furnaces —
for instance, fir cones, peat, the lowei
kinds of coal, &c. — the same gas is ob-
tained 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 ac
companying drawing. The fuel lies on
the fire-bars O, the air enters through FM.M^Oas.nnid|loer for the formfttion of carbon
them and the ash-hole (drawn by the monoxide for heating purposes,
draught of the chimney of the stove where
the gas burns, or else forced by a blowing apparatus), the quantity of tax being exactly
regulated by means of valves. The gasjs formed are then led by the. tuba V, provided
398 PRINCIPLES OF CHEMISTRY
coal 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 £ by volume) and
nitrogen (§ 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.80 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 regene-
rative furnace?* By means of this process not only may the high
temperatures indispensable in many industries be obtained (for instance,
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 com-
bustible material and covered with a lid.
*° 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-ban, 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°Ms 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).
** Regenerative furnaces were introduced by the Brothers Siemens about the year
1660 in many industries, and mark a most important progress in the use of fuel, espe-
cially 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 com-
bustion 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 tempera-
ture 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
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 899
glass- working, steel-melting, &c.), but great advantage also23 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.
The transformation of carbonic anhydride, by means of charcoal,
into carbonic oxide (C + CO2 = 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 6tted, 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
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
K 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 an 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 con-
cerning combustion.
44 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 (lesa
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 CO2, and
that the formation of CO2 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 CO2 + H2=CO + H2O.
Rathke (1881) showed that at no temperature whatever is the reaction as expressed by the
equation CO2 + C = 2CO2, 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
untransfonned into carbonic oxide, even after the action has been continued for several
hours. The endothermal reactions, C -t- 2HSO = CO2 + 2H2, and CO + H2O = CO2 + H2,
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,
400 PRINCIPLES OF CHEMISTRY
of electric sparks also decomposes carbonic oxide into carbonic anhydride
and carbon, and if the carbonic anhydride be removed by alkali com-
plete 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..O
= H2 + 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 (H2 + CO)
are formed. This mixture of combustible gases is called water gas.™
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 pub-
lished 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 O'OOl atmosphere. At the atmospheric pressure, not more than
O'Oo 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.
'•>* b'* 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 hydiogen) is able,
when heated even to 850°, to completely decompose carbonic oxide into CO-i and carbon,
which remains with the nickel and is easily removed from it by heating in a stream of
air. Here 2CO = CO..> + 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).
w A molecular weight of this gas, or 2 volumes CO (28 grams), on combustion
{forming CO..>) gives out 68,000 heat units (Thomsen 67,960 calories). A molecular weight
of hydrogen, H_, (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. Char.
coal, resolving itself by combustion into the molecular quantity of CO3 (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., = 2CO absorbs 89,000 heat units; (8) C + H2O = H., + 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 + CO2) ; (4) C + H..,O = C0.> + 2H.2 absorbs (if the
water be in a state of vapour) 19,000 heat units; (5) the reaction 00 + 1120 = 002 + 111
develops 10,000 heat units if the water be in the state of vapour ; and (6) the decomposi-
tion expressed by the equation 2CO = C + CO2 (Note 24 bis) is accompanied by the evolu-
tion of 89,000 units of heat.
Hence it follows that 2 volumes of CO or Ho burning Into CO2 or H^O develop
almost the same amount of heat, just as also the heat effects corresponding with the
equations
C+CO|«GO.+CO
*re nearly equal.
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 401
But aqueous vapour (and only when strongly superheated, otherwise
it cools the charcoal) only acts on charcoal to form a large amount of
i I 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 4- 2H2O = CO2 + 2H2. Besides
this, carbonic oxide on splitting up forms carbonic anhydride, and
therefore water gas always contains a mixture26 in which hydrogen
predominates, the volume of carbonic oxide being comparatively less,
K 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 nitrogenJrom 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 car-
bonic anhydride could be as readily obtained in as pure a state as water, then CO might be
prepared directly from COo + 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 byvthe
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 deter-
mined 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 + H2O = CO + HU.
402 PRINCIPLES OF CHEMISTKY
whilst the amount of carbonic anhydride increases as the temperature
of the reaction decreases (generally it is more than 3 per cent.)
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 carbo-
nates (MgC03 or BaC03) 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, CH202, on being heated to 200°, easily decomposes into carbonic
oxide and water, CH202 = 00 + H2O.27 bl» Usually, however, car-
bonic oxide is prepared in laboratories, not from formic but from oxalic
acid, C2H204, 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, <fcc. ; 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 C2HaO4 = HaO + C02 + CO. This decomposition is
generally practically effected by mixing oxalic acid with strong sul-
n The so-called yellow prussiate, I^FeCeNg, on being heated with ten parts of strong
sulphuric acid forms a considerable quantity of very pure carbonic oxide quite free from
carbonic anhydride..
17 bb 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.
K 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 Hjin which one or both halves of the hydrogen
are exchanged for carboxyl ; therefore they are equal to H3 + CO2 and Hs+ 2CO2 ; but
H2 reacts with CO3, as has been stated absve, forming CO and H2O. From this it is also
evident that oxalic acid on losing C03 forms formic acid, and also that the latter may
proceed from CO + H20, as we shall see further on.
COMPOUNDS OF CAKBON WITH OXYGEN AND NITROGEN 403
phuric 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.MM*
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
wi>is Greshoff (1888) showed that with a solution of nitrate of silver, iodoform, CHIj,
forms CO according to the equation C HI3 + 3 AgN O3 + H2O = 8 Agl + 3HNO3 + CO. The
reaction is immediate and is complete.
19 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 proceed* and spreads
•lowly unless there be perfect dry ness. I think that this may be 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 dry ness, many chemical transformations (for
example, the formation of ozone from oxygen, the decomposition of AgO, KClOj 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 S03 does not act upon dry CaO or CnO ; (2)
perfectly dry sal-ammoniac does not give NH3 with dry CaO, but simply volatilises ; (8)
dry NO and O do not react ; (4) perfectly dry NH3 and HC1 do not combine ; (5) perfectly
dry sal-ammoniac does not dissociate at 350° (Chapter V II., Note 15 bis) ; and (6) perfecfly
dry chlorine does not act upon metals, <fec.
90 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).
404 PRINCIPLES OF CHEMISTRY
incomplete combustion of diarcoal 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 hydro-
gen, 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 H2, 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, COC12 •
that is to say, the molecule of hydrogen, H2, 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 car-
bonyl chloride. This characterises the so-called diatomic or bivalent re-
actions of radicles or residues. H is a mohatomic residue or radicle,
like K, Cl, and others, whilst carbonic oxide, CO, is an indivisible (un-
decomposable) bivalent radicle, equivalent to H2 and not to H, and
therefore combining with X2 and interchangeable with H2. This,
distinction is evident from the annexed comparison .
HH, hydrogen. CO, carbonic oxide.
HC1, hydrochloric acid* COC12, carbonyl chloride.
HKO, potash. CO(KO)2, potassium carbonate.
HNH2, ammonia. CO(NH2)2, urea.
HCH3, methane. CO(CH3)2, acetone.
HHO, water. CO(HO),, carbonic acid.
Such monatomic (univalent) residues, X, as H, Cl, Na, NO2, NH4,
CH3, C02H (carboxyl), OH, and others, in accordance with the law
of substitution, combine together, forming compounds, XX , and with.
COMPOUNDS OF CARBON WITH' OXYGEN AND NITROGEN 405
oxygen, or in general with diatomic (bivalent) residues, Y— for instance,
O, CO, OH2, S, Ca, <fec. forming, compounds XX' Y ; but diatomic
residues, Y, sometimes capable of existing separately may combine
together, forming YY' and with X2 or XX', as we see from the transi-
tion of CO into CO2 and COC12. This combining power of carbonic
oxide appears in many of its reactions. Thus it is very easily ab-
•sorbed by cuprous chloride, CuCl, dissolved in fuming hydrochloric
acid, forming a crystalline compound, COCu2Cl2,2H2O, decomposable
by water ; it combines directly with potassium (at 90°), forming
(KCO)n 3I with platinum dichloride, PtCl2, with chlorine, C12, <fec.
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 cars
bonic 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 po'tash 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 CHKO2 is obtained by this absorption ; it cor-
responds with an acid found in nature — namely, the simplest organic
(carboxyl'.c) acid,/brwiic acid, CH2O2. 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(CHO2), carbonic acid, HO(CHO2), and oxalic acid,
(CHO2)2, are the simple organic or carboxylic acids, R(CHO'2) cor-
?' 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
K2C202 is formed, and probably K^C^On), -because products containing C10 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 1878 in a colourless form by means of
NHSK, described in Chapter VI., Note 14) have greatly extended our knowledge of thi»
compound, much still remains unexplained. It probably exists in various polymeric and
isomeric forms, having the composition (KCO)» and (NaCO)*.
406 PRINCIPLES OF CHEMISTRY
responding 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 X2, that is of forming COX2. 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, CH4,
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 com-
pounds of carbon of those compounds which contain nitrogen. By
way of an example, we will take one of the carboxyl acids, R(CO2H),
where R is a hydrocarbon radicle (residue). Such an acid, like all
others, will give by combination with NH3 an ammoniacal salt,
R(CO2NH4). 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(CO2NH4), forming by the loss of one molecule of water, amides,
RCONH2, and by the loss of two mofecules 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 recog-
nised 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, HCO2H, carbonic, OHCO2H,
and oxalic, (CO2H)2. Cyanogen, like carboxyl, is evidently a mon-
atomic residue and acid, similar to chlorine. As regards the amides
RCONH2, corresponding to the carboxyl acids} they contain the
ammoniacal residue NH2, and form a numerous class of organic com-
pounds met with in nature and obtained in many ways,33 but not
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, (CNH2O)2, is obtained in
the form of an insoluble precipitate on adding a solution of ammonia to an alcoholic
solution of ethyl oxalate, (CO2C2H5)2, which is formed^ by the action of oxalic acid on
alcohol: (CHO2)2 + 2(C2H5)OH = 2H6H + (CO2C2H5)2. As the nearest derivatives of
ammonia, the amides treated with alkalis yield ammonia and form the salt of the acid.
The nitrites do not, however, give similar reactions so readily. The majority of amides
corresponding to acids have a composition RNH2, and therefore recombiue with water with
great ease even when simply boiled with it, and with still greater facility in presence of
distinguished by such characteristic peculiarities as the cyanogen com-
pounds.
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(NH4)CO3, and
normal carbonate, (NH4)2C03 ; they represent compounds of one oj
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, NaHCO3), 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(NH4)HCO>
«=s HaO + C02 + (NH4)2CO3 ; the latter, however, decomposes in solu-
tion, and can therefore only be obtained in crystals, (NH4)2CO3,HaO, at
low temperatures, and from solutions containing an excess of ammonia
as the product of dissociation of this salt : (NH4)2CO3 = NH$
+ (NH4)HC03. But the normal salt,34 according to the general type, is
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 th»
amide was derived is formed ; RNH.> + KHO = RKO + NH3.
The same reaction takes place with acids, only an ammoniacal salt of the acid is of
coarse formed whilst the acid held in the amide is liberated: RNH.2 + HC1 + H2O
= RHO + NH4C1.
Thus in the majority of cases amides easily pass into ammoniacal sails, bat 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 vola-
tilise without alteration and frequently are volatile crystalline substances which may bo
easily sublimed. Such, for instance, are the amides of benzoic, formic, and many
other organic acids.
54 The acid salt, (NI^JHCOj, on losing water ought to form the carbdmic acid,
OH(CNHgO) ; but it is not formed, which is accounted for by the instability of the acid
ealt itself. Carbonic anhydride is given .off and ammonia is produced, which give*
t^nimomum_cj,rbamate.
408 PRINCIPLES OF CHEMISTRY
capable of decomposing with separation of water, and forming ammonium
carbamate, NH4O(CONH2) = (NH4)2CO3 - H2O ; this still further
complicates the chemical transformations of the carbonates of am-
monium. 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, 2NH4C1 + CaC03 = CaCl2 + (NH4)2CO3. 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 NH4O(CONH2) + 2OH(CO2NH4)
•=4NH3 + 3CO2 + 2H2O. This salt, in parting under various con-
ditions 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 ca/bonate 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.
Each of the two ammoniacal salts of carbonic acid has its corre-
sponding amide. That of the acid salt should Ve acid, if the water given
offtakes up the hydrogen of the ammonia, as it should according to the
common type of formation of the amides, so that OHCONH2, or
carbamic acid, is formed from OHCO3NH.i.. This acid is not known in
a free state, but its corresponding ammoniacal salt or ammonium car-
bamate is known. The latter is easily and immediately formed by
mixing 2 volumes of dry ammonia with 1 volume of dry carbonic anhy-
dride, 2NH3 -1- CO2 = NH4O(CONH2) ; 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
55 In the normal salt, 2NH3 + CO., + H20, in tbe acid salt, NHj + C02 + H2O, but in the
commercial salt only 2H8O to 3CO2.
36 Naumann determined the following dissociation tensions of the vapour of ammonium
carbamate (in millimetres of mercury) : —
-10° 0° +10° 20° 80° 40° 60° 60°
5 12 80 62 124 248 470 770
Horstmann and Isambert studied the tensions corresponding to excess of NH3 or CO2,
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.
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 409
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 car-
bonate 3r — as in an anhydrous state ammonia and carbonic anhydride
only form one compound, C022NH3.38 As the normal ammonium car
bonate 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(NH2)2. This must be termed
carbamide. It is identical with urea, CN2H40, which, contained in the
urine (about 2 per cent, in human urine), is for the higher animals
(especially the carnivorous) the ordinary product of excretion39 and
oxidation of the nitrogenous substances found in the organism, If
ammonium carbatnate be heated to 140° (in a sealed tube, Bazaroff),
or if carbonyr chloride, COC12, 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, CN2H40 + H2O = CO2 4- 2NH3.
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 GN2H4O,HNO3),
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,
CUHO,39tlis which is a nitrile of carbonic acid — that is to say, is a
"V 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 CaCl2, probably because calcium carbamate is soluble in water, and all the
(NH3)2C02 is not converted by dissolving into the normal salt, (NJH4Q)2COS.
58 It must be imagined that the reaction takes, place at first between equal volumes
(Chapter VII.).; but then carbamic acid, HO(CNH20), -is produced, which, as an acid,
immediately combines with the ammonia, forming NH40(CNH.iO).
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 SO 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.
59 ">'» Its polymer, CsNjHjOs, is formed together with it. Cyanic acid Is a. very
unstable, easily changeable liquid, while cyanuric acid is a crystalline solid which is very
gtable at the ordinary temperature.
410 PRINCIPLES OF CHEMISTRY
'cyanogen compound, corresponding to the acid ammonium carbonate,
OH(CNH402), which on parting with 2H2O ought to form cyanic acid,
CNOH. Liquid cyanic acid, exceedingly unstable at the ordinary
temperatures, gives its stable solid polymer cyanuric acid, O3H3C3N3.
Both have the same composition, and they pass one into another at
, different temperatures. If crystals of cyanuric acid be heated to a tem-
perature, 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 vapodr contains cyanic acid, and, if it be rapidly cooled, it con-
denses into a mobile volatile liquid (specific gravity at 0°= 1*14). If
the liquid cyanic acid be gradually heated, it passes into a new amor-
phous 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 temperatui'e,\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.cyauuric acid.
The properties .of cyanic acid which we have described were principally
observed by Wohler, and clearly show the faculty of polymerisation of
cyanogen compounds. This is observed in many other cyanogen deriva-
tives, and is to be regarded as the consequence of the above-mentioned
explanation of their nature. All cyanogen compounds are ammonium
salts, B(CNH4O2), deprived of water, 2H2O ; therefore the molecules,
RON, ought to possess the faculty of combining with two molecules of
water or with other molecules in exchange for it (for instance, with
H2S, or HC1, or 2H2, &c.), and are therefore capable of combining to-
gether. The combination of molecules of the same kind to form more
complex ones is what is meant by polymerisation.40
40 Just as the aldehydes (such as C2H4O) are alcohols (like C2H«O) which have
lost hydrogen and are also capable of entering into combination with many substances,
and of polymerising, forming slightly volatile polymerides, which depofymerise
on heating. Although there are also many similar phenomena (for instance, the trans-
formation 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.
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 411
Besides being a substance very prone to form polymerizes, 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 + H20 = C02 -f NH3. A
solution of ammonium cyanate, CN(NH4)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 sub-
stances is identical, CN2H4O, 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, NH4, 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 bo 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(NH2)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 Wbhler in
1828.41 Formamide, HCONH2, and hydrocyanic acid, HCN", as a
nitrile, correspond with formic acid, HCOOH, and therefore ammonium
formate, HCOONH4, and formamide, when acted on by heat and by
substances which take up water (phosphoric anhydride)f orm 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
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 NH4OCN into CO(NH2)2
is the best example of the passage of one system of equilibrium of atoms into another
more stable system.
& 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 hydro-
cyanic acid one atom of hydrogen is under the influence of two acid-forming elements.
Acetylene, C2H2, which contains but little hydrogen, presents acid properties in certain
respects, for its hydrogen is easily replaced by metals. HydronitrouB acid, HNj,
which contains little hydrogen, also has the properties of an acid.
412 PRINCIPLES OF CHEMISTRY
— 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, potas-
sium cyanide (KHO + HCN = H?O + KCN) in solution— have a
strongly alkaline reaction.43 If ammonia be passed over charcoal at
a fed hea,t, 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.43bis ' Potassium cyanide is much
used in the arts, and is obtained, as above stated, under many circum-
stances— 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 K4FeC6NG + 2H2O. The name
of cyanogen (KVOVOS) is derived from the property which this yellow
prussiate possesses of forming, with a solution of a ferric salt, FeX3,
the familiar pigment Prussian blue. The yellow prussiate is manu-
43 Solutions of cyanides — for-'instance, those of potassium, or barium — are decom-
posed 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 hydro-
cyanic 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^CN^ is formed, which
is a soluble double salt.
45 bi» The'c^pnyersibn oi 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, noi 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 each as ammonia, which is of great importance in
agriculture.
COMPOUNDS OF CARBON WITH OXYGEN AND NITROGEN 413
factured on a large scale, and is generally used as the source of the
other cyanogen compounds.
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 decom-
poses, 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 i815 Gay-
Lussac finally settled its properties and showed that it contains only
hydrogen, carbon, and nitrogen, CNH. If the distillate (a weak solu-
tion of HCN) be redistilled, and the first part collected, the anhy-
drous 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
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
** 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 orde'r 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 be 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)j.
If a strong solution of hydrochloric acid be poured upon these crystals, and the mix-
ture 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 con-'
densed. 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 + H2S = 2HCN + HgS.
414 PRINCIPLES OF CHEMISTRY
proportions in water, alcohol, and ether weak aqueous solutions are
used in medicine.45
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 con-
sider 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 HgC2N2, 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, (GN)a
or CNCN, corresponds to hydrocyanic acid in the same manner as free
chlorine, C12 or C1C1, corresponds to hydrochloric acid. This composition,
iudging 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
45 A weak (up to 2 p.c.) aqueous solution of hydrocyanic acid is obtained by the dis-
tillation 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 :
C10H,TNOn + H.jO = C7H«O + CNH tC&itQe
Amygdalin in Water Bitter Hydrocyanic Glucose
bitter almonds almond oil acid
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.
46 This salt will be described in Chapter XIII.
COMPOUNDS OF CAKBON WITH OXYGEN AND NITROGEN 415
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 : HgC2N2 = Hg
+ C2N2.47 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^ 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
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 prus-
siate and mercuric chloride may be used, then double decomposition and the formation
of mercuric cyanide take place in the retort. Silver oyanide also disengages cyanogen,
on being heated.
48 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 pheno-
menon 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 = 580° 581° 600° 685°
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 850°, 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.
*6
416 PRINCIPLES OF CHEMISTRY
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.
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)/9
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,'0 the
nitriles (or cyanogen compounds, RON), and nitro-compounds (con-
taining the radicle of nitric acid, RN02), 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 : C2N2 + 2NaHO = NaCN + CNNaO + H2O. But the
latter salt decomposes relatively easily, and moreover part of the cyanogen liberated by
heat from its compounds undergoes a more complex transformation.
40 If , in general, compounds containing the radicle NH2 are called amides, some of the
amines ought to be ranked with them ; namely, the hydroc'arbons CnH2m, in which part of
the hydrogen is replaced by N H2 ; for instance, meth ylamine, CHjNH2, aniline, CeHsNH^,
fee. In general the amines rnay be represented as ammonia in which part or all of the
hydrogen is replaced by hydrocarbon radicles — as, for example, trimethylamine, N(CH3)j.
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.
417
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 com-
binations 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, H2O, „ „ H2R.
Ammonia, H3N, „ „ H3R.
Marsh gas, H4C, „ „ H4R.
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 composi-
tion 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
H2M, H3M, HnMm. then we must conclude, on the basis of the law of
substitution, that this element will give compounds M2O, M;,N, MHO,
MH3C, <tc. Chlorine is an example of this kind. If we know that
another element, R, like oxygen, gives with hydrogen a molecule H2R,
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
418 PRINCIPLES OF CHEMISTRY
to foretell, if not the properties (for example, the acidity or basicity),
at any rate the composition,1 of some of their compounds. This forma
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), R2O, RC1, R3N, R4C, «fcc. Oxygen,
in that form in which it gives water, is the representative of the
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 CO2, COX2, CH4, 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 atomr
city to a greater degree than most of the other elements. Nitrogen in NH3, NH2(OH),
N2O3, and even in CNH, must be considered as trivalent, but in NH4C1, NO^(OH), and
in all their corresponding compounds it is necessarily pentavalent. In N2O, if the
atomicity of oxygen = 2, nitrogen has an uneven atomicity (1, 8, 5), whilst in NO it is
bivalent. *• If sulphur be bivalent, like oxygen, in many of its compounds (for example,
H2S, SC12, KHS, &c.), then it could not be foreseen from this- that it would form SO2,
SO3, SC14, SOClj, and a series of similar compounds in which its atomicity must be
'acknowledged as greater than 2. Thus SO2," Sulphurous anhydride, has many
points in common with CO.,, and if carbon be quadrivalent then the S in SO? ia
quadrivalent. Therefore the principle of atomicity (valency) of the elements cannot be
considered established as the basis for the study of the elements, althcmgh 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 con-
sideration of the properties of the elements : 1. Such univalent elements as H, Cl, &c.,
appear in a free state as molecules H2, C12, &c., and are consequently like the nnivalent
radicles CHj, OH, CO2H, &c., which, as might be expected, appear as C2H6, O2H2,
Ci^Hj (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 ih&tfree affinities may exist. Granting
this, nothing prevents, the assumption that free affinities exist in all unsaturated com-
pounds; for example, two free affinities in NH3. 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^H — where uni-
valent elements are combined in molecules which are more complex than B2, and form
molecules, Rj, R4, &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 be considered as a fundamental property.
4. Crystallo-hydrates (for instance, NaCl,2H2O, or NaBr,2H2O), double salts (such as
PtCl4,2KCl,H2SiFe, &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. There-
fore 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— BERTHOLLET'S LAWS 419
bivalent elements, forming RH2, RO, RC12, RHC1, R(OH)C1, R(OH)2,
R2C, RON, &o. Nitrogen in ammonia is the representative of the
trivalent elements, giving compounds RH3, R203, R(OH)3, RClj, RN,
RHC, (fee. In carbon are exemplified the properties of the quadrivalent
elements, forming RH4, RO2, RO(OH)2, R(OH)4, RHN, RC1<, RHC13,
«fec. 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 quan-
titative analogies or resemblances, which are foretold by the law of
substitution (Chapter VI.), there exist among the elements qualita-
tive 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 an.d 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 sub-
stance, and of hydrochloric acid, and of the base sodium hydroxide,
formed by the non-metal chlorine and the metal sodium, which corre-
spond with it.
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 forma-
tions, 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
' The primary formations art 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.
420 PRINCIPLES OF CHEMISTRY
also for the formation of other masses of workable salt, such as rock
salt, and of saline springs and lakes. 2bi*
The extraction of saltyVom 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
ad vantage 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 Baume'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 arti-
ficial evaporation, about one part of tolerably pure table salt at the
very commencement of the operation ; the total amount held in solu-
tion being about 2^ p.c. The remaining* portion separates out inter-
z bls 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. Miintz (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 0'31 milligram of chloride
of sodium, while a litre of rain collected from the valley contained 2'5-7'6 milligrams.
G 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 ther"e by natural means and only requiring a treat-
ment and refining, which is also necessary in addition for the ' summer salt ' obtained
•from sea-water
SODIUM CHLORIDE— BERTHOLLET'S LAWS 421
mixed with the bitter salts of magnesium which, owing to their solu-
bility 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, CaSO42H2O, 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 evapora-
tion the chlorides of potassium and magnesium begin to separate
in a state of combination, forming the double salt KMgClg^HjO,
-which occurs in nature as carnallite.* 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 proportion6 of magnesium salts
(when it acquires a bitter taste) the remaining liquor is run back into
the sea.
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 ; up
heavals 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
4 The double salt KCl,MgCL> is a crystallohydrate of KC1 and MgCl2) and is only
formed from solutions containing an excess of magnesium chloride, because water de-
composes 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 S04, CO3), which are disposed in
every possible kind of grouping ; for instance, K as KC1, KBr, K2S04, Mg as UgClj,
MgBr2, MgSO4) 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 NaoSO^lOH^O is first deposited because it is the first to arrive at
saturation at low temperatures. Consequently this solution contains MgCl2 and Na-jSO.),
besides MgSO4 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.
422 PRINCIPLES OF CHEMISTRY
in those localities where there are deposits of gypsum. But inasmuch1
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 dis-
tricts 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 sub-
soil 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 im-
pervious 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
SODIUM CHLORIDE-BERTHOLLET'S LAWS 423
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 rook 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 llis
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 bb ij, Western Europe, deposits of rock salt have long been known at Wieliczkft,
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 froip 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
424
PRINCIPLES OF CHEMISTRY
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.
If a saline spring, or the salt water pumped from bore-holes, con-
tains but little salt, then the first concentration of the natural solution
FIG. 64. — Graduator for the evaporation of the water of saline spriugs.
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
the Government of Erivan in the Caucasus; (d) the'Katchiezmansky deposit in the
province of Kara ; (e) the Krasnovodsky deposit in the Trans-Caspian province; and (/)
the Bardymkulsky saltmines in Kokhand.
SODIUM CHLORIDE— BERTHOLLET'S LAWS 425
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 evapora-
tion 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 Ibs. 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 pro-
cesses, 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, <fec., and of chlorine
compounds (bleaching powder and hyrdochloric acid). The yearly pro-
duction of salt in Europe amounts to as much as 1\ million tons.
Although certain lumps of rock salt and crystals of bay salt some-
times consist of almost pure sodium chloride, still the ordinary com-
mercial salt contains various impurities, the most common of which are
426 PRINCIPLES OF CHEMISTRY
magnesium salts. If the salt be pure, its solution gives no precipitate
with sodium carbonate, Na2CO3, showing the absence of magnesium
salts, because magnesium carbonate, MgCO3, 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 car*
ried 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 solu-
tion of hydrochloric acid), while the impurities remain in solution. By
repeating the operation and fusing the salt (when adhering hydro-
chloric acid is volatilised) a pure salt is obtained, which is again
crystallised from its solution by evaporation.9
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-
8 The fracture of rock salt generally shows the presence of interlayers of impurities
which are sometimes very email 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 850 interlayers for the whole thickness (about 85 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 800 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.
10 According to the determinations of Klodt, the Briansk rock salt withstands a
pressure of 840 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.
SODIUM CHLORIDE— BERTHOLLET'S LAWS 427
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 ordinafry salts, be considered as non- volatile,
although as yet no exact experiments have been made in this direction.
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 crystal-
lisation ; u but if the crystals are formed at a low temperature,
11 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 coolei;
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 =216, 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 pkce 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 chmes) and at various depths.
13 By combining the results of Poggiale, Miiller, and Karsten (they are evidently
more accurate than those of Gay-Lussac and others) I found that a saturated solution at
<°, from 0° to 108°, contains 85'7 + 0'024< + 0-0002<* grams of salt per 100 grams of water.
This formula gives a solubility at 0° = 85'7 grams ( = 26'8 p.c.), whilst according to Kar»
sten it is 86-09, Poggiale 85-5, and Miiller 35-6 grams.
M 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
428 PRINCIPLES OF CHEMISTRY
especially from a saturated solution cooled to — 12*, then they present
a prismatic form, and contain two equivalents of water, NaCl,2H2O.
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,10H20 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.
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
air, according to Stas. (In the Briansk mines, where the temperature throughout the
whole year ia about + 10°, it may be observed, as Baron Klodt informed me, that in the
summer daring 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 hygro-
scopic. If it contain any magnesium chloride, it partially deliquesces in a damp atmo-
sphere. The crystallised and not perfectly pure salf 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,2H2O also crystallises either in plates or prisms.
Sodium bromide also crystallises with 2H.jO 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 con-
sidered 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
= -0'6c to c = 10, according to Rosetti = — 0'649c to c = 8'7, according to De Coppet
(to c = 10)= -0-56 c - 0'006c», according to Karsten (to c = 10) -0-762c + 0'0084c?, 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 —
t= 0-649 x 5815 = 2'05.
18-5
Pickering (1893) gives for c=l-0'603, for c = 2-l'220; that is (c up to 2'7) about
- (0-600 + 0'005c)c.
The data for strong solutions are not less contradictory. Thus with 20 p.c. of salt, ice
is formed at -14-4° accprding to Karsten, -17° according to Guthrie, — 17'6° according
to De Coppet. Riidorff states that for strong solutions the temperature of the formation
of ice descends in proportion to the contents of the compound, NaCl,2H2O (per 100 grams
of water) by 0°'342 per 1 gram of salt, and De Coppet shows that there is no proportior
ality , in ft strict sense, for either a percentage of NaCl or of NaCl,2H2O.
SODIUM CHLORIDE-BERTHOXLET'S LAWS 429
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 SI5 = 9991-6 + 7l'I7p + 0'2140/>2.
For instance, for a solution 200H2O + NaCl, in which case p = 1'6,
S15 = 1-0106. It is seen from the formula that the addition of water
produces a contraction. l8 The specific gravity I9 at certain temperatures
and concentrations in vacuo' referred to water at 4° = 10,000 20 is here
given for
0°
15°
30°
100°
: 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 Tiaum^'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
" 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 dis-
tances of the globules (that is, the length of a column limited by layers of definite con-
centration) 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 — Ksr = log. (A — q),
where A, B, and K are constants ; that at a given moment the rate of diffusion of tho
different layers is proportional to their depth, &c.
18 If So be the specific gravity of water, and S the specific gravity of a solution con-
taining p p.c. of salt, then by mixing equal weights of water and the solution, we shall
obtain a solution containing fa of the salt, and if it be formed without contraction, then its
211
specific gravity x will be determined by the equation — =— •+• • , because the volume is
x So S
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 St, then as a cubic centimetre of air under the usual
conditions weighs about 0-0012 gram, the sp. gr. in a vacuum S = Si + 0'0012 (Sl — 1),.
if the density of water = 1.
"° If the sp. gr. Sj 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 S2 by the sp. gr. of water at the tem-
perature of observation.
430 PRINCIPLES OF CHEMISTRY
solution. Common salt is somewhat soluble in alcohol,21 but it is
insoluble in ether and in oils.
Common salt gives very few compounds " (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 elec-
trode 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 Ms 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 + HC1
= NaCl + H20.
21 According to Schiff 100 grams of alcohol, containing p p.c. by weight of C2HaO,
dissolves at 16° —
p = 10 20 40 60 80
28-5 22-6 18-2 5'9 1'2 grams NaCl.
83 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 hydro-
chloric acid until chlorine ceases to be liberated, is a remarkable one. Its composition is
NaIOs,NaCl,14H2O. 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 + HoO = NaHO + HC1). 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 anodo 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 chlorne and caustic soda is an
application of this method which has been already worked in England on an industrial
scale.
SODIUM CHLORIDE-BERTHOLLET'S LAWS 431
With respect 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 + MCI
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 hydro-
chloric 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
roust be taken in food. Not only do animals consume those small
amounts of sodium chloride which are found in drinking water or in plants
432 PRINCIPLES OF CHEMISTRY
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 + H2SO4 HC1 + NaHS04
Sodium chloride Sulphuric acid Hydrochloric acid Acid sodium 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 + H2SO4 = 2HC1 + Na-jSO,
Table salt Sulphuric acid Hydrochloric acid Sodium sulphate
This double decomposition is the result of the action of the acid
salt, NaHSO4, first formed, on sodium chloride, for the acid salt,
since it contains hydrogen, itself acts like an acid, NaCl + NaHSO4
= HC1 + Na2SO4. 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 sulpTiuric acid. The sulphuric acid in
each case displaces the other, volatile, acid.
SODIUM CHLORIDE— BERTHOLLET'S LAWS 433
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 tempera-
ture 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 con-
siderable 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
** If MX and NY represent the molecules of two salts, and if there be >u> third
tubitancc 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 + H2SO4) where Cl is a univalent haloid and SO4 is bivalent, then the
matter may be complicated by the formation of other compounds besides MY and NX, and
434 PRINCIPLES OF CHEMISTRY
the double decomposition of salts in which M and N 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 -f 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, KNO3, and well-dried sodium
acetate, C2H3NaO2, 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, NaN03,
and potassium acetate, C2H3KO2, both of which are highly hygro-
scopic.24 Ws
"When Berthollet enunciated his doctrine the present views of atoms
and molecules had yet to be developed, and it is now necessary to sub-
mit 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, ^Ca or ^Mg (bivalent elements) replace hydrogen.
And since, according to Berthollet's doctrine, when wMX of one
salt comes into contact with nNY of another salt, a certain quantity
wMY and o:NX is formed, there remains in — 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 nMY + wNX + (m — w)MX — that is, a portion of one
only of the -salts taken would remain unchanged because the reaction
could only proceed between wMX and wNY If x were actually equal
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 con-
sider 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.
2« 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 820°
SODIUM CHLOKIDfi-BEBTHOLLErS LAWS 485
to n, the mass or the salt MX would not have any influence ou the
modus operandi of the reaction, which is equally in accordance with
the teaching of Bergmann, who supposed double reactions to be inde-
pendent 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
M
their ratio - less than unity — that is, the decomposition will be ex-
n
pressed by the equation, mMX + wNY = (m — a;)MX + (n — as)NY
+ xMY + «NX ; (2) by increasing the mass m we increase the de-
composition— that is, we increase x and the" ratio / — - — -, until with
\n — x)
an infinitely large quantity m the fraction — will equal 1, and the de-
n
composition 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 — «)NY + a;MY + a;NX. These direct consequences of Ber-
thollet'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 KC1 + NaN03
there might arise KNO3 + 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. Experi-
ment proves the latter hypothesis to be the true one. But before
citing the most historically important experiments verifying Berthollet's
436 PRINCIPLES OF CHEMISTRY
doctrine, we must stop to consider the conception of the num-of the
reacting substances. Berthollet understood by mass the actual rela-
tive 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 + H2S04 the salt is taken in one equivalent and the acid in
two. If 2NaCl + H2S04 act, then the number of equivalents are equal,
and so on. The influence of mass on the amount of decomposition
* forms the root of Berthollet's doctrine, and therefore we will first of
n
all turn our attention to the establishment of this principle in relation
to the double decomposition of salts.
About 1840 H. Rose 9* showed that water decomposes metallic sul-
phides like calcium sulphide, CaS, forming hydrogen sulphide, H2S,
notwithstanding the fact that the affinity of hydrogen sulphide, as an
acid, for lime, CaH2O2, as a base, causes them to react on each othen
forming calcium sulphide and water, CaS + 2H20. 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 an timonious -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. 2S bis
** 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 pre-
decessor in this branch of research was Berzelius, and his successor Fresenius.
25 MS 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 be 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 tho
mass of hydrochloric acid forming one of the products of the reaction.
SODIUM CHLORIDE -BERTHOLLET'S LAWS 437
Barium sulphate, BaSO4, which is insoluble in water, when fused
with sodium carbonate, Na2CO3, gives, but not completely, barium
carbonate, BaCO3, (also insoluble), and sodium sulphate, Na2SO4. 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 carbo-
nate 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 deter-
mined 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 attain-
ment 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 (ZnSO4 and
2NaCl) were taken, there were produced by exchange sodium sul-
phate 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
iaken had been decomposed. When, however, a mixture of solutions
of sodium sulphate and zinc chloride was taken, the precipitate pre-
sented the same composition as before — that- is, about 28 per pent, of ,the
salts taken had been subjected to decomposition. In a similar experi-
ment with a mixture of sodium chloride and magnesium sulphate,
SNaCl + MgSO4 or MgCl2 + Na2SO4, about half of the metals under-
went the decomposition, which may be expressed by the equation
488 PRINCIPLES OF CHEMISTRY
4NaCl + 2MgS04 = 2NaCl + MgS04 + NaiS04 +MgCJ2 = 2Na2SO4 +
2MgCl2. 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 (BaCO3 + Na2SO4) 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 ap-
plied for determining the composition of a solution without the removal
of any of its component parts. When Gladstone took equivalent quanti-
ties of ferric nitrate and potassium thiocyanate — Fe(N03)3 + 3KCNS
— only 13 per cent, of the salts underwent decomposition. On
increasing the mass of the latter salt the quantity of ferric thio-
cyanate 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 NO3
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 NO3,
on the other hand. The investigation of the variation of the fluor-
escence 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 appre-
ciated 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 trans-
formations received a very important support in the researches of
Bertbelot 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
SODIUM CHLORiDE-BERTHOLLET'S LAWS 439
differs in that it proceeds slowly at the ordinary temperature, extend-
ing 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, C2H5OH, with acetic acid, HC2H3O2, gives the follow-
ing system rapidly when heated, or slowly at the ordinary temperature,
ROH + HX + 2RX + 2H20, -whether we start from 3RHO + 3HX
or from 3RX 4- 3H2O. 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 etheritication 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 be-
tween 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 w =1) and the salts MY and NX are
*7
440 PRINCIPLES OF CHEMISTRY
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 (I — x) (1 - x) and for the second pair k'xx, and a
state of equilibrium or limit will be reached when k (1 — a)2 = k'x3t
whence the ratio k/k' = [x/(l — x)f. Therefore in the case of the
action of alcohol on an acid, when x = |, the magnitude k/k' = 4,
that is, the reaction of the alcohol on the acid is four times as fast as
thai of the ethereal salt on water. If the ratio kfk1 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 a; = 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
But it is impossible to subject the formation of salts to any process
directly analogous to that which is so conveniently effected in etherifi-
cation. 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,
CaCO3, from the mixture of solutions of calcium chloride and sodium or
potassium carbonate, and found in this case that not only was the
86 From the above it follows that an excess ot acid should influence nh« 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. ot 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.
SODIUM CHLORIDE— BERTHOLLET'S LAWS 441
rate of action (for example, in the case of CaCl2 + Na2CO3, 75 per
cent, of CaCO3 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, inas-
much as a portion of the substance is obtained or remains in the form of
a precipitate, so that the system is heterogeneous. The investigation o£
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
27 As an example two methods_may be mentioned, Thomson's and Ostwald's. Thomsen
(1869) applied a thenno-chemical method to exceedingly dilute solutions without taking"
the water into further consideration. He took solutions of caustic soda containing
lOOHjO per NaHO, and sulphuric acid containing iH,S04 + lOOH^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 ^ •. heat
(n + O'o}
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 hydro-,
chloric 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 tUe
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 HNO3 was taken per JNa.2SO4, 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 NajSO.!, 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 JNa-jS04 +.HNOS
442 PRINCIPLES OF CHEMISTRY
is much more simple, and therefore its significance in confirmation
of Berthollet's doctrine -is of particular importance. The only cases
and NaNO3-t- iH2SO.i is equal, as for ethereal salts, to 4. By taking this figure and ad-
mitting 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.
Thomsen concludes his investigation with the words : (a) ' When equivalent
quantities of NaHO, HNO3 (or HC1) and AH2S04 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 desig-
nates 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 them-
selves 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..,O5 + H2O gives +3,600 heat units, and the
solution of the resultant hydrate, 2NHOj, in.a large excess of water evolves + 14,986 heat
units. The formation of S03 + H.>O evolves + 21,308 heat units, and the solution of
H>SO4 in an excess of water 17,800 — that is, sulphuric acid gives more heat in both cases.
The interchange between Na2SO4 and '2HNO- is not only accomplished at the expense of
the production of NaNOj, but also at the expense of the formation of HoSO4, 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
SODIUM CHLORIDE— BERTHOLLET'S LAWS 44S
•which can be compared with these reactions for simplicity are those
exchange decompositions investigated by G. G. Gustavson, which
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.2SO4 there was formed a solution of NajSO4 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 HNO5 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! = 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 Q2 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! and Q3, 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
eulphuric acid of the composition SO3 + 100H2O, occupying a volume of 1,815 c.c., with a
corresponding quantity of a solution 2(NaHO + 5H2O), whose volume = 1,793 c.c., we
obtain not 8,608 but 3,633 c.c., an expansion of 25 c.c. per gram molecule of the resulting
salt, Na-jSO^ 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 con-
sidered 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 sepa-
rated from such solutions with great difficulty. After mutual saturation they form the
salt Na2SO4, 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 com-
bination ; 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.,>O has a sp. gr. 2*8, hence its molecular
volume = 22 ; the sp. gr. of SO5 is 1-9 and volume 41, hence the sum of their volumes is 63 ;
for N 8^804 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 H4SO4=53*8, that of 2NaHO = 37'4 ; there ia
produced 2H2O, volume = 36, + NaaSO^, volume = 53'6. There react 90'7 c.c., and on satu-
ration there result 89-6 c.c. ; consequently contraction again ensues, although less, and
444 PRINCIPLES OF CHEMISTRY
take place between CC14 and RBrn on the one hand, and CBr4 and
RC1,, on the other. This case is convenient for investigation inas-,
much as the RC1» and RBr,, taken (such as BC13, SiCl4, TiCl4, POC13,
and SnCl4) belong to those substances which are decomposed
by water, whilst CC14 and CBr4 are not decomposed by water ; and
therefore, by heating, for instance, a mixture of CC14 -h SiBr4 it is
possible to arrive at a conclusion as to the amount of interchange
by treating the product with water, which decomposes the SiBr4 left
unchanged and the SiCl4 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 in-
stance, 4BC13 + 3CBr4. 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 4BBr3 + 3CC14. In this last 89'97 per cent, of
bromine in the BBr3 was replaced by chlorine ; that is, there were
obtained 89'97 molecules of BC13 and there remained 10'02 molecules
of BBr3, and therefore the same state of equilibrium was reached as
that given by the system 4BCla + 3CBr4. Both systems gave one and
the same state of equilibrium at the limit, which is in agreement with
Berthollet's doctrine.28
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 2NaNO3 + H2SO4 = 2HNO; + Na7SO4, the volumes vary but slightly: in the
above example they are 2(38'8) + 53 8 and 2(41'2)+'58'6; hence 181 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 $ 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 thn
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, notwith-
SODIUM CHLORIDE-BERTHOLLET'S LAWS 445
Thus we now find ample confirmation from various quarters for the
following rules of Berthollet, applying them to double saline decom-
positions • 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 rela-
tive 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
standing its importance and the fact that the theoretical side of the subject (thanks
especially to Guldberg and Van't Hoff) has since been rapidly poshed forward. If the re-
searches 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 CBr4, and consequently the less the amount of bromine
replaced by chlorine by the action of CC14 on bromine compounds. For instance, for
chlorine compounds the percentage of substitution (at the limit) is —
BC1S SiCl4 TiCl4 AsClj SnCl4
10-1 12-5 48-6 71-8 77'6
It should be observed, however, that Thorpe, on the basis of liis 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 CBr4 be heated with RC14, then an exchange of the bromine
for chlorine takes place. But what would be the result if it were mixed with CC14 ?
Judging by the magnitude of the atomic weights, B = ll, 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 CC14
and CBr4 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 CBr4 is, within the limit, substituted by the chlorine of the CC14 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 CC14 all the atoms of 01 do not remain constantly combined with
the same atoms of carbon, and there is an exchange of atoms between different mole-
cules in a homogeneous medium also. This hypothesis may in my opinion explain
ceitain 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 Pfauudler 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.
19 Berthollet's doctrins 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
446 PRINCIPLES OF CHEMISTRY
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 solu-
tion, 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 com-
position 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 trans-
formations is applicable with great clearness to a number of re-
actions 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
email that even a large mass would still give no observable displacements. The funda-
mental 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, CC14 + 2H2O = CO2 + 4HC1), just as there are non- volatile sub-
stances. 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 conies 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 solutior
SODIUM CHLORIDE— BERTHOLLET'S LAWS 447
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
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 l>e 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 mix-
ture 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
M 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 decom-
position will only in a few cases render it possible for new metallic chlorides to be ob-
tained, 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 solu-
tion. 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,+ MgSO4 = MgCl2 + Na2SO4 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 Stassf urt in a natural form. It might be
said that this form of double decomposition is only accomplished with a change of tem-
perature ; 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
jf 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.
448 PRINCIPLES OF CHEMISTRY
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 head-
ing 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 ol 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 Na2SO4 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
51 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 boilsover 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 retard 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 FeClj, 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 bath.
Phosphoric anhydride absorbs hydrogen chloride (Bailey and Fowler, 1888;
2P.2O5 + 3HCl = POClj + 3HPO5) at the ordinary temperature, and therefore the gas
cannot be dried by this substance.
SODIUM CHLORIDE— BERTHOLLET'S I+A.WS
449
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
FIG. 85. — 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.
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
51 bta in chemical works where sulphuric acid of 60° Baiiine* (22 p.c. of water) is
employed, 117 parts of sodium chloride are taken to about 125 parts of sulphuric acid
450 PRINCIPLES OF CHEMISTRY
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.
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 dennite direction. Here it encounters a cur-
rent 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 absorp-
tion 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
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^O + rzHCl when boiled (the residue will be
almost HC1,8H3O) evolve (not water but) a solution of the composition H2O + 445n4HCl ;
for example, on distilling HC1,10H.3O, HC1,23H2O is first obtained in the distillate. As
the strength of the residue becomes greater, so also does that of the distillate, and there-
fore 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 pur-
poses (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 hydro-
chloric acid.
SODIUM CHLORIDE— BERTHOLLETS LAWS 451
fiowc 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 absorp-
tion of the hydrochloric acid is carried on in the so-called coke totvers,
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.
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.
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 hydro-
53 Thus the metallic chlorides, which are decomposed to a greater or less degree by
water, correspond with feeble bases. Such are, for example, MgCl2, AlCIj, SbClj, BiClj.
The decomposition of magnesium chloride (and also carnallite) by sulphuric acid pro-
ceeds at the ordinary temperature ; water decomposes MgCl.j 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.
54 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 se«a that the expansion of this
liquid is greater than that of gases (Chapter II., Note 84).
452 PRINCIPLES OF CHEMISTRY
chloric 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 HNO3), 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
whichalso approaches to 25 p.c. of acid, and more nearly as the tempera-
ture 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 hydro-
gen 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 IIC1,6H2O corresponds with 25-26 p.c. HC1 — 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 HC1,6H2O. Besides this
hydrate there exists also a crystallo-hydrate, HC1,2H2O,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
The mean specific gravities at 15°, taking water at its maximum
55 According to Roscoe and Dittraar 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.
57 This crystallo-hydrate (obtained by Pierre and Puchot, and investigated by Rooze-
boom) is analogous to NaCl,2H2O. The crystals HC1,2H2O at -22° have a specific
gravity 1'46 ; the vapour tension (under dissociation) of the solution having a composition
HC1,2H.20 at -24° = 760, at -19° = 1,010, at -18° = 1,057, at -17° = 1,112 mm. of mer-
cury. In a solid state the crystallo-hydrato 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 HC1,2H2O at — 18°, the temperature
of solidification falls to -25°, and the hydrate HC1,3H2O is formed (Pickering, 1898).
58 According to Roscoe at 0° one hundred grams of water a a pressure p (in millimetre*
of. mercury) dissolves —
p .- 100 200 800 500" 700 1,000
Grams HC1 65'7 70'7 73'8 78'2 81'7 66'6
SODIUM CHLORIDE— BERTHOLLET'S LAWS 458
density (4°) as 10,000, for solutions containing p per cent, of hydrogea
chloride are —
p S' p S
5 lu,242 25 11,266
10 10,490 30 11,522
15 10,744 35- 11,773
20 11,001 40 11,997
The formula 5=9,991-6- -I- 49-43/)+0'0571/»2, up to j9=25'26, •which
answers to the hydrate HC1,6H2O mentioned above, gives the specific
gravity. Above this percentage S = 9,785-1 + 65'10p — 0'240/>8. The
At a pressure of 760 millimetres and temperature t, one hundred grams of water
dissolves
t = 0 8° 16° 24° 40° 60°
Grams HC1 82'5 78'8 74-2 70'0 68'3 661
Boozeboom (1886) showed that at t° solutions containing c grams of hydrogen
per 100 grams of water may (with the variation of the pressure p) be formed together
with the crystallo-hydrate HC1,2H2O:
t = -28°-8 -21° -19° -18° -17°-7
c = 84-2 86-8 92-6 98'4 101'4
p = — 884 580 900 1,078 mm.
The last combination. answers to the melted crystallo-hydrate HC1,'2H30, 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 98-8 95-7 89'8 84*2
From these data it is seen that the hydrate HC1,2H2O 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 HC1+ wH2O (at about 80*,
taking the specific heat of water = 1) is given by the expression —
C(86'5 + m!8) = 18m- 28'89 + 140/m - 268/m»
if m be not less than 6'25. For example, for HC1 + 25H2O, C = 0'877.
According to Thomsen's data, the amount of heat Q, expressed in thousands of calorie;,
evolved in the solution of 86:6 grains of gaseous hydrochloric acid in wiHjO or 18m grams
of water is equal to —
m = 2 4 10 50 400
Q = 11-4 14-8 16-2 17-1 17'8
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 HClnHoO at 0° :—
n = 6 6-9 9-8 14 27'1 129*5
It =• 2-81 2-08 1-86 1-67 1-52 1*89
IV also appears 'that strong solutions diffuse more rapidly into dilute solutions than
into water.
454 PRINCIPLES OF CHEMISTRY
rise of specific gravity with an increase of percentage (or the differential
. ) reaches a maximum at about 25 p.c.39 The intermediate solution,
dp)
HC1,6H2O, 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<), 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 --] rises with a rise of tern-
\ at/
perature.41
p= 0 5 10 15 20
S0 — S15 = 7-2 23 38 52 64
5,5-530 = 34-1 42 50 59 67
Whilst for solutions which cdntain a greater proportion of hydrogen
chloride than HC1,6H2O, these coefficients decrease with a rise of
temperature ; for instance, for 30 p.c. of hydrogen chloride S0 — Si&
= 88 and 5,,-, — S30 =87 (according to Marignac's data). In the cas^,
of HC1,6H2O these differences are constant, and equal 76.
Thus the formation of two definite hydrates, HC1,2H2O and
HC1,6H20, 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 pot-
ash, lime, <fec. In a dry state, however, hydrochloric acid does not alter
39 If it be admitted that the maximum of the differential corresponds wit
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( = So-fcSo<, or V( = !/(!- fc<) 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 the 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 <° it = 10,580 - <(2'13 + 0'027<). Whence also may be found the coefficient
of expansion (Note 40).
SODIUM CHLORIDE— BEKTHOLLET'S LAWS 455
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, <fec.,
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 H2SO4, but which, like copper, decompose it to the
limit of SO2) displace the hydrogen from hydrochloric acid Thus
hydrogen is disengaged by the action of zinc, and even of copper and
tin.42Ms 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 ami prevents the further
action of the acid on the metal. The same is to be remarked with re-
spect 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 MCI ; for
example, NaCl, KC1, AgCl, HgCl, if the metal replaces hydrogen
equivalent for equivalent, or, as it is said, if it be monatomic or
univalent. In the cpse of bivalent metals, they have a composition
MC12 ; for example, CaCl2, CuCl2, PbCl2, HgCl2, FeCl2, MnCl2. The
•composition of the haloid salts of other metals presents a further
•variation ; for example, A1C13, PtCl4, <fec. Many metals, for instance
•Fe, give several degrees of combination with chlorine (FeCl2,FeCl3)
as with hydrogen. In their composition the metallic chlorides differ
from the corresponding oxides, in that the O is replaced, by C12, as should
'follow from the law of substitution, because oxygen gives OH2, and is
48 Thus, for instance, with feeble bases they evolve in dilute solutions (Chapter III.,
Note 58) almost equal amounts of heat ; their relation to sulphuric acid is quite identical.
They both form faming solutions as well as hydrates ; they both form solutions of con-
etant boiling point.
« w» Pybalkin (1891) found that copper begins to disengage hydrogen at 100°, and
(hat 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.
456 PRINCIPLES OF CHEMISTRY
consequently bivalent, whilst chlorine forms HC1, and is therefore
univalent. So, for instance, ferrous oxide, FeO, corresponds with
ferrous chloride, FeCl2> and the oxide Fe2O3 with ferric chloride, which
is also seen from the origin of these compounds, for FeCl2 is ob-
tained by the action of hydrochloric acid on ferrous oxide or- carbonate
and FeCl3 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 HC1 and MnCl2m without any oxygen bear the name
of haloid salts ; for instance, HC1 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, MC12, and water, is limited at
high temperatures by the reverse reaction MC12 + H2O = MO + 2HC1,
and the more pronounced are the basic properties of MO the feebler is
the reverse action, while for feebler bases such as A12O3, MgO, <fcc., 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 BaCl4
corresponding with the peroxide BaO2. 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 CaCl2 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 corro-
sive sublimate, HgCl2, is particularly volatile, whilst the oxide HgO
decomposes at a red heat. Silver chloride, AgCl, is fusible and
is decomposed with difficulty, whilst Ag2O 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 pre-
cipitates 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,
HgC2 + Cu = CuCl2 + Hg, and hydrogen at a red heat displaces silver
from silver chloride, 2AgCl -f H2 = Ag2 + 2HC1. 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
SODIUM CHLORIDE— BERTHOLLET'S LAWS 457
temperature, <fcc., as was shown in the preceding portions of this chapter,
•and as will be frequently found hereafter.
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 SO3HC1 ; 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 compo-
sition of its molecule, belongs to the monobasic acids, and does not,
therefore, give true acid salts (like HNaS04 or HNaCO3) ; nevertheless
many metallic chlorides, formed from powerful bases, are capable oi.
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 HC1 for combining with M,,Clm is much more frequently
encountered than was previously supposed. Thus, for instance, dry hydro-
chloric acid when passed into a solution of zinc chloride (containing an
excess of the salt) gives in the cold (0°) a compound HCl,ZnCl2,2H2O,
and at the ordinary temperature HCl,2ZnCl2,2H2O, just as it is able at
low temperatures to form the crystallo-hydrate ZnCl2,3H2O (Engel,
1886). Similar compoundsare obtained with CdClj.ChiCla, HgCl2,Fe2Cl6)
<fec. (Berthelot, .Ditte, Cheltzoff, Lachinoff, and others). These com-
pounds with hydrochloric acid are generally more soluble in water than
the metallic chlorides themselves, so that whilst hydrochloric acid
decreases the solubility of M,.Clm, 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 un-
saturated hydrocarbons (for instance, with turpentine, CIOH1G,2HC1) and
their derivatives. Sal-ammoniac, or ammonia hydrochloride, NH4C1
= NH3,HC1, also belongs to this class of compounds.43 If hydrogen
chloride gas be mixed with ammonia gas a solid compound consisting
45 When an unsaturated hydrocarbon, or, in genera], an unsaturated compouud)
assimilates to itself the molecules CL, HC1, SO3, ELSOj, &c., the cause of the reaction
is most simple. As nitrogen, besides the type NX3 to which NH3, belongs, gives com-
pounds of the type NXS — for example, NO2(OH) — the formation of the salts cf
458 PRINCIPLES OF CHEMISTRY
of equal volumes of each is immediately formed. The same com-
pound 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 T55. 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
ammonium should be understood in this way NH3 gives NH4C1 because NX5 is
capable of giving NX5. But as saturated compounds — for instance, S03 H2O, NaCl,
&c. — are also capable of combination even between themselves, it is impossible
to deny the capacity of HC1 also for combination. SO5 combines with H3O, and also with
HC1 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, PC13 as an atomic compound and PC15 as a molecular one, only
because it easily splits up into molecules PC13 and Cla.
44 Sal-ammoniac is prepared from ammonium carbonate, obtained in the dry distilla-
tion 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 sublima-
tion. 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 por-
tion, 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 NH4C1, at 50° about
50 parts, and at the ordinary temperature about 85 parts. This is sometimes taken
advantage of for separating NH4C1 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 87'28 41'72 46 55 61 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 + Sl-26p — O'085/i2, where pis the amount by weight
of ammonium chloride in 100 parts of solution. With the majority of salts the differen-
tial 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. titcherbakofi).
459
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, ' 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 de-
composition 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. 2bls
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.
* 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.
* bu 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 hydro-
Chloric acid with hydrogen, Lavoisier and Berthollet supposed it to be a compound with
oxyeen of OB anhydride contained in hydrochloric acid. They suppossd that hydro-
460 PRINCIPLES OF CHEMISTRY
A.n aqueous solution of hydrochloric acid is generally employed for
the 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 2HC1. forming water, H20, and setting the
chlorine free, 2HC1 -•- O (disengaged by the oxidising substances)
i=z H2O + C12. 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.8
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, <fec
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 4HC1 and Mn02, in which the manganese takes the place
of the four atoms of hydrogen, or the chlorine and oxygen exchange
places — that is, MnCl4 and 2H2O are produced. The chlorine com-
pound, MnCl4, 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, MnCl4 = MnCla
-f C12 3 bis The action of hydrochloric acid requires a temperature of
chloric acid contained water and' the oxide of a particular radicle, and that chlorine was
a higher degree of oxidation of this radicle mitnas (from the Latin name of hydrochloric
acid, acidum muriaticum). It was only in 1811 that Gay-Lussac and Thenard in
France and Davy in England arrived at the conclusion that the substance obtained by
Scheele does n6t 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 \\<ap6ft
signifying a green colour, because of the peculiar colour by which this gas is chanio
terised.
3 However, nitric acid has been proposed as a means for obtaining chlorine, but by
methods which have the drawback of being very complicated.
3bl< This representation of the process of the reaction is the most natural. However,
this decomposition is generally represented as if chlorine gave only one degree of combi-
nation with manganese, MnCl2, and therefore directly reacts in the following manner—
MnO2 + 4HC1 = ,MnCl2 + 2H20 + C12, in which case it is supposed that manganese
peroxide, MnO2, breaks up, as it were, into manganous oxide, MnO, and oxygen, both of
which react with hydrochloric acid, the manganous oxide acting upon HC1 as a base,
giving MnCl2 and at the same time 2HC1 + 0 = H2O + C12. In reality, a mixture of oxygen
THE HALOGENS 461
about 100° In the laboratory the preparation of chlorine is carried on
in flasks, heated over a water-bath, by acting on manganese peroxide
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 (Mn2O3, MnO2, MnO3, MnjO;), with the exception of man-
gftnous oxide, MnO, disengage chlorine from hydrochloric acid, because manganous
chloride, MnCl.i, 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 oxygeii by chlorine, C13 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 com-
pounds 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, SbCl5, which splits up into
chlorine "and antimony trichloride when heated. Cupric chloride, corresponding with
copper oxide, and having a composition CuCL>, 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 Cu2Cl3O. With
hydrochloric acid this substance gives cupric chloride (Cu.jCl.jO + 2HC1 = H2O + 2CuCl2),
which has only to be dried and heated jn order again to obtain chlorine. Thus, in solution,
and at the ordinary temperature, the compound CuCl2 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 CuS04 and Na^SO^. CuCl2 is then formed by the double
decomposition of the salt of copper and the hydrochloric acid; the CnCl2 liberates
chlorine, and the CuCl forms CujCljO with the oxygen of the air, which again gives CuClj.
with 2HC1, and so on.
Magnesium chloride, which is obtained from sea-water, caruallite, &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 CaClg in this manner. A solu-
tion of CaCl.i, containing a certain amount of common salt, is evaporated and oxide of
magnesium added to it. When the solution attains a density of T2445 (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-Eeychler (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-
462 PRINCIPLES OF CHEMISTRY
with hydrochloric acid or a mixture of common salt and sulphuric
acid4 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
not participate m any way in the reaction, is added to the mixture to prevent its fusing.
The reactions may be expressed by the following equations: (1) 8MgCl2 + SMnClj + 8O
= Mg5MnsO8 + 12C1 ; (2) Mg3Mn5Os + 16HC1 = SMgClj + SMnCl, + 8H2O + 4C1. 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 HNO3,
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
CaCl2+SiOi + O(air) = CaOSiOi + C12, the second on the action of the oxygen of the air
(heated) upon MgCl2 (and certain similar chlorides) MgCLj + O = MgO + C12. The remaining
MgO is treated with sal-ammoniac to re-form MgCl2 (MgO + 2N H4C1 = MgClj + H2O + 2NH3)
and the resultant NH-, again converted into sal-ammoniac, so that hydrochloric acid
is the only substance consumed The latter processes have not yet found much appli-
cation.
4 The following proportions are accordingly taken by weight : 5 parts of powdered man-
ganese peroxide, 11 parts of salt (best fused, to prevent its frothing), and 14 parts of sul-
phuric 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
eulphur, and becomes brittle under the action of the chlorine).
The reaction which proceeds may be expressed thus. MnO.j + 2NaCl + 2H2SO«
= MnSO4 + Na,2S04 + 2H.>0 + Cl2. The method of preparation of C12 from manganese
peroxide and hydrochloric acid was discovered by Scheele, and from sodium chloride by
Berthollet.
* The reaction of hydrochloric acid upon bleaching powder gives chlorine without
the aid of heat, CaCl.2O2 + 4HC1 = CaCl > + 2H2O + 2C12, 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 (Merrae", Kammerer). 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.
\ mixture of potassium dichromate and hydrochloric acid evolves chlorine perfectly
free from oxygen (V. Meyer and Langer).
6 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 con-
eists 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
THE HALOGENS 463
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
condenses7 into a liquid which has a yellowish-green colour, a
density of T3, 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^8 At 0° one volume of water dissolves about 1£ volume
of chlorine, at 10° about 3 volumes, at 50° again l£ volume.9 Such
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 man-
ganese 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 + a;CaCl2, which is found to be the best propor-
tion, 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, MnOj, and oxide of chlorine on a large
manganese, Mn2O5. This is due to the manganous oxide absorb-
ing 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 MnCl2
is the only one which is stable (see Note 8). 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.
7 Davy and Faraday liquefied chlorine in 1823 by heating the crystallo-hydrate
CljSILjO in a bent tube (as with NH3), 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 Luchvig's observations (1868), and from the fact that the coefficient of
expansion of gases increases with their molecular weight (Chapter tt., Note 26, for hydrogen
= C'867, carbonic anhydride = 0-873, hydrogen bromide = 0-886), 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 Pelonze, determined that the solubility increases between
0° and 8°-10° (from 1J to 2 vols. of chlorine per 100 vols. of water at 0° up to 8 to 2J 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 solu-
bility of chlorine in the presence of hydrogen— even in the dark. Berthelot determined
:8
464 PRINCIPLES OF CHEMISTKY
a solution of chlorine is termed ' chlorine water ; ' and is employed in
a. diluted form in medicine and as a laboratory reagent. It is pre-
pared 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, C12,8H2O, which easily splits up
into chlorine and water when heated, so that jf 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
Chlorine explodes with'hydroyen, if a mixture of equal volumes be
exposed to the direct action of the sun's rays " or brought into contact
an increase of solubility with the progress of time. Schb'nbem and others suppose that
chlorine acts on water, forming hypochlorous and hypochloric acids, (HC1O + HC1)
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 after-
wards 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 0D'24 and + 28°'7
(wheu 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 C1.2,10H2O, but Rooze-
boom (1885) showed that it is poorer in water and =C12,8H2O. At first small, almost
colourless, crystals are obtained, but they gradually form (if the temperature be below
kheir critical point 28°'7, ibove which they do not exist) large yellow crystals, like those of
potassium chromate. The specific gravity is T28. The hydrate is formed if there be
more chlorine tn 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 0° = 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 forma-
tion 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 280-7,
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 hetero-
geneous 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 coloured glass, otherwise oxygen is evolved and hydrochloric acid
formed.
11 The chemical action of light on a mixture of chlorine and hydrogen was discovered by
Gay-Lussac ajid Thenard (1809). It has been investigated by many savants, and especially
THE HALOGENS 465
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
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. 1-16) of hydro-
chloric acid, in which case the water is not decomposed — that is, no oxygen becomes
mixed with the chlorine.
l! 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 actino-
meter — 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 coloured green by a salt of copper evinces more chemical action than the colourless
flame, but the flame brightly coloured 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 hydro-
chloric 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 ot
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, ^fo of hydrogen by 38 p.c.,
f&9 of oxygen by 10 p.c., r&g of chlorine by 60 p.c., &c. According to the researches of
Kliraenko and Pekatoros (1889), the photo-chemical alteration of chlorine water is
retarded by the presence of traces of metallic chlorideti, 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 U
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.
466 PRINCIPLES OF CHEMISTRY
reaction of chlorine on hydrogen occupies (at the original temperature
and pressure) a volume equal to the sum of the original volumes , that
IB, a reaction of substitution here takes place H2 + C12 = HC1 + HC1
In this reaction twenty-two thousand heat units are evolved for one
part by weight [1 gram] of hydrogen.13
These relations show that, the affinity of chlorine for hydrogen is
very great and analogous to the affinity between hydrogen and oxygen.
Thus u 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 sun-
light, 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 H2O + C12 = 2HC1 + 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.
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 1G 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
15 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
Thomson's, and, for Na^O, Beketoff's results)
f2NaCl, 195; CaCU, 170; HgClo, 68> 2AgCl, 69.
\ Na.2O, 100; CaO, 181; HgO, 42 ; Ag2O, 6.
< 2AsCl3) 143; 2PC15, 210 ; CC14, 21 ; 2HC1, 44 (gas).
\ A8jO3, 155 ; P3O5, 870 ; CO7, 97 ; H2O, 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 HC1 and the
oxide, whilst the remainder have other properties, as is seen from the fact that they are
not formed from hydrocliloric acid and the oxide, but give hydrochloric acid with water
l* This has been already pointed out in Chapter III., Note 5.
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
18 An instructive experiment on combustion in chlorine maybe conducted as follows:
leaves of Dutch metal (used instead of gold for gilding) are placed in a glass globe, and a
THE HALOGENS 467
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 eta-
ployed 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 hydro-
chloric 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 act*
ing 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.
The majority of non-metals also react directly on chlorine ; hot
sulphur and phosphorus burn in it and0 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, PC13, and sulphurous chloride,
<fcc., do not have the properties of salts, and, as we shall afterwards see
more fully, correspond to acid anhydrides and acids ; for example, PC13
— to phosphorous acid, P(OH)3
NaCl FeCl2 SnCl4 PC13 HC1
Na(HO) Fe(HO), Sn(HO)4 P(HO)3 H(HO)
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, PtCU, is then formed, whilst this substance de-
composes 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 platiqous 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, NO2C1 (nitric acid, NO2OH, in which HO is replaced by chlorine ; see Chapter on
Phosphorus), and that of nitrous acid, NOC1 — but these do not act on gold. The
foimationof aqua regia may therefore be expressed by 4NHOs+8HCl=2NOaCl + 2NOCl
+ GHjO + 2C12. The formation of the chlorides NO.|C1 and NOC1 is explained by the fact
that the nitric acid is deoxidised, gives the oxides NO and NO2,and they directly combine
with chlorine to form the above anhydrides.
468 PRINCIPLES OF CHEMISTRY
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,
PC13 + 3H20 = P(HO)3 + 3HC)
Phosphorus ™ fc Phosphorous Hydrochloric
trichloride acid acid
whilst other chlorine compounds are formed from hydroxides and
hydrochloric acid, with the liberation of water, for example,
NaHO -i- HC1 = NaCl -«- H2O
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) + HC1 == MCI + H20
hydrate hydrochloric acid = chloranhydride * water
and if the hydrate ROH be acid, then,
RC1 + H2O = R(HO) + HC1
Chloranhydride -t water = hydrate + hydrochloric acid
The chloranhydrides MCI corresponding to the bases are evidently
metallic chlorides or salts corresponding to HC1. Tn 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, C12, 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, NO2, SO2, and other lower oxides which are
capable of combining with oxygen may also combine with a corre-
THE HALOGENS 469
spending quantity of chlorine. Thus CaCl2, NOC1, NO2C1, SO2C12>
«fec., are obtained. They correspond with the hydrates CO(OH)2,
NO(OH), NO2(OH), SO2(OH)2, &c., and to the anhyhrides CO2,
N2O3, N2O5, SO3, &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 Cl2 : 0 ; (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 PC13 and PC15, it is evident that
PC15 is the higher form of combination compared with PC13. To the
form PC15, or in general PX5, correspond PH4I, PO(OH)3, POC13, <fec.
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 CnH2n+2
very distinctly. The unsatu rated hydrocarbons are sometimes able to
combine with chlorine with the greatest ease and thus reach the limit.
Thus ethylene, C2H4, combines with C12, forming the so-called Dutch
liquid or ethylene chloride, C2H4C12, because it then reaches the limit
CnX2n+;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, C2H4(OH)2, is obtained
from C2H4C12.
Chlorine in the presence of water very often acts directly as an
oxidising agent. A substance A combines with chlorine and gives, for
example, AC12, 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 + H2O + C12 = 2HC1 + 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 hydro-
chloric 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 : S02 + 2H2O + C12 = H2SO4 + 2HC1.
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 con-
470 PRINCIPLES OF CHEMISTRY
verted 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 bleach-
ing— 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
The power of chlorine for combination is intimately connected with
its capacity for substitution, because, according to the law of substitu-
tion, 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 cMorine which combines
with the atom of hydrogen is also able to replace the atom of hydrogen.
We mention this property of chlorine ^iot 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.
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
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.
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.
THE HALOGENS 471
continues to bum for a certain time, a large amount of soot is obtained,
and hydrochloric acid is formed. In this case the gaseous and incan-
descent 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.
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 ancl 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 hydro-
chloric 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.*2 The scheme of the process is as
follows :
C.H.iX 4- C12 = CnHroMClX + HC1
Hydrocarbon Free chlorine Product of metalepsis Hydrochloric acid
Or, in general terms —
RH 4= C12, = RC1 + HCL
The conditions under which metalepsis takes place are also very
constant. In the dark chlorine does not usually act on hydrogen com-
n 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.
M 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.
472 PRINCIPLES OF CHEMISTRY
pounds, 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
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
CH4 + Olj = CH3C1 + HC1
Marsh gas Chlorine Methyl chloride Hydrochloric acid
The volume of the mixture remains unaltered. The methyl
chloride which is formed is a gas. If it be separated from the hydro-
chloric 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
25 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 Petroftsky 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, C6He, 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 operancli of this instructive reaction
is based on the property of aluminium bromide to enter into combination with hydro-
carbons and their derivatives. The details of this and all researches concerning the
metalepsis of the hydrocarbons must be looked for hi 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.
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 Q2H€ + C1J = C2H5C1 + HC1, according to the data given by
Thomson, evolves about 20,000 heat units, whilst the formation of hydrochloric acid
evolves 22,000 units
THE HALOGENS 473
subjected to a further metalepsical substitution — the second atom of
hydrogen may be substituted by chlorine, and a liquid substance,
CH2C12, called methylene chloride, will be obtained. In the same
manner the substitution may be carried on still further, and CHC13>
or chloroform, and lastly carbon tetrachloride, CC14, 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 anaesthetic ; 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 quanti-
ties 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 CnCl2m 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, CH3-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— CH2C1'COOH, CHC12'COOH, and
CC13-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 com?
plexity 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, CH3-OH is obtained from CH3C1. By the
action of metallic derivatives of hydrocarbons — for example, CH3Na —
the chlorine also gives a salt, and the hydrocarbon radicle — for instance,
474 PRINCIPLES OF CHEMISTRY
CH3 — takes the place of the chlorine. In this, or in a similar manner,
CH3-CH3, or C2H6 is obtained from CH3C1 and C6H5-CH3 from C6Hfi.
The products of metalepsis also often react on ammonia, forming hydro-
chloric acid (and thence NH4C1) and an amide ; that is, the product of
mefcalepsis, with the ammonia radicle NH2, &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 com-
plex 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
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, NH3, and calcium hydroxide,
80 With the predominance of the representation of compound radicles (this doctrine
dates from Lavoisier and Gray-Lussac) in organic chemistry, it was a very important
moment iu 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..,H5, or the radicle of
common alcohol, C.jHj'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, C2H5H = CgHg, it
was looked on as containing the same ethyl, just as methyl hydride, CH4 = CH3H, waa^
considered as existing in methane. Having obtained free methyl, CH5CH5=C.2H8, from
it, it was considered as a derivative of methyl alcohol, CH;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.2H5 = CH2CH3. In its time a still greater impetus was given by
the study of the reactions of monochloracetic acid, CHjCl'COOH, or CO(CH2C1)(OH).
It appeared that metalepsical chlorine, like the chlorine of chloranhydrides — for instance,
of methyl chloride, CH3C1, or ethyl chloride, C._,H5C1 — is capable of substitution ; for
example, glycollic acid, CH,(OH)(CO,H), or CO(CHo'OH)(OH), was obtained from it, and
it appeared that the OH in the group CH2(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. (Sen
Chapter VIII., Note 42.)
n By including many instances of the action of chlorine under metalepsis we not
only explain the indirect formation of CC14, NC1S, and C12O 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 repre-
sentations to the law of substitution we may foretell metalepsis as a particular case of a
general law.
THE HALOGENS 475
CaH.2O2 or Ca(OH)2, may be replaced by chlorine and give potassium
hypochlorite, KC1O, calcium hypochlorite, CaCl2O2, and the so-called
chloride of nitrogen, NC13. 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 + C12 = HC1 + KC1O, but also KHO + HC1 = H2O + KC1,
and the result of both simultaneous phases will be 2KHO + Clz
= H,O + KC1 + KC1O. We will here discuss certain special
cases.
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 CH4). 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, 8NH3 + 3C12 = 6NH4C1 + N4.
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 NH3 + 3C12 forms NC13 + 3HC1.59
K This may be token 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 mast be present.
29 The hydrochloric acid formed combines with ammonia, and therefore the final result
is 4NH5 + SC12=NC15+3NH4C1. 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
(he -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: NCl5 + 4NH5=N2-l-SNH4Cl. Therefore, so long as the liquid is alka-
line from the presence of ammonia the chief product will be nitrogen. The reaction
NH4C1 + SCL = NC13 + 4HC1 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
476 PRINCIPLES OF CHEMISTRY
The resulting product of metalepsis, or chloride of nitrogen, NC13>
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
(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 re-
action NC13 + 4HC1 = NH4C1 + 3C12 enabled Deville and Hautefeuille to determine the
composition of chloride of nitrogen. When slowly decomposed by water, chloride of
nitrogen gives, like af chloranhydride, nitrous acid or its anhydride, 2NC13 + SH2O
= N2O5 + 6HC1. From these obseryations it is evident that chloride of nitrogen presents,
great chemical interest, which is strengthened by its analogy with trichloride of phos-
phorus. The researches of F. F. Selivanoff (1891-94) prove that NClg may be regarded
as an ammonium derivative of hypochlorous acid. Chloride of nitrogen is decomposed by
dilute sulphuric acid in the following manner : NC15 + 3H2O + H2SO4 = NH4HS04 + 3HC10.
This reaction is reversible and is only complete when some substance, combining vrith
HC1O (for instance, succiuimide) or decomposing it, is added to the liquid. This ia
easily understood from the fact that hypochlorous acid itself, HC1O, 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 NC13 as H2O does to NH3, or as RHO to
RNH2, R2NH, and R3N — that is to say, NC13 corresponds as an ammonium derivative to
C10H and C12 in exactly the same manner as NR3 corresponds to ROH and R2. The
connection of NCIs and other similar explosive chloro-nitrogen compounds (called
chloryl compounds by Selivanoff ; for example, the C2H5NC12 of Wurtz is chloryl ethyl-
amine), such as NRC12 (as NC2H5C12), and NR,C1 (for instance, N(CH3CO)HC1, chloryl-
acetamide, and N(C2H5)2C1, chloryl diethylamine) with HC10 is evident from the fact
that under certain circumstances these compounds give hypochlorous acid, with water, for
instance, NR2C1 +. H2O = NRjH + HC1O, and frequently act (like NC13 and HC1O, or C12)
in an oxidising and chloridising manner. We may take chloryl succinimide, C2H4(CO)2NC1
for example. It was obtained by Bender by the action of HC1O upon succinimide,
C2H4(CO)oNH, and is decomposed by water with the re-formation of amide and HC1O
(the reaction is reversible). Selivanoff obtained, investigated, and classified many of
the compounds NR2C1 and NRC12, 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, NRs, where
R is fully substituted by chlorine.
29 bi> In preparing NC1S every precaution must be used to guard against an explosion,
and care should be taken that the NG13 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 pre-
pared 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 op 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 con-
taining mercury, and first a saturated solution of common salt is poured into the funnel,
THE HALOGENS 477
Chloride of nitrogen is a yellow oily liquid of sp. gr. 1-65, which
boils at 71°, and breaks up into N + C13 at 97°. The contact of
phosphorus, turpentine, india-rubber, <fec. 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
tinits for NC13, 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, NaCIO, and sodium chloride are simul-
taneously formed : 2NaHO -f C12 = NaCl + NaCIO + H2O. The
resultant solution contains NaCIO 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 + 2C12
= CaCl2O.2 -f- CaCl2 + 2H2O. 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 contains30 both these substances. It
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.
80 Quicklime, CaO (or calcium carbonate, CaC05), does not absorb chlorine when cold,
]>ut at a red heat, in a current of chlorine, it forms calcium chloride, with the evolution
of oxygen. (This was confirmed in 1898 by Wells, at Oxford.) This reaction corresponds
with the decomposing action of chlorine on methane, ammonia, and water. Slaked lime
(calcium hydroxide, CaH2O2) also, when dry, does not absorb chlorine at 100°. The
absorption proceeds at the ordinary temperature (below 40°). The dry mass thus ob-
tained contains not less than three equivalents of calcium hydroxide to four equivalents
of chlorine, so that its composition is [Ca(HO)o]3Cl4. In all probability a simple absorp-
tion 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.jO, 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
478
PRINCIPLES OF CHEMISTRY
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 2CaCla
+ Ca(C103)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).
KlO. 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.
The products of the metalepsis of alkaline hydrates, NaCIO and
Ca(ClO)2, which are present in solutions of ' Javelle salt ' and bleaching
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 formulae : From the
dry mass Ca3(HO)6Cl4 there is formed lime, Ca(HO)2, calcium chloride, CaCl2, and a
saline substance, Ca(ClO)2. Ca3H6O6Cl4 = CaH2O2 + CaCl2O2+CaCl? + 2H2O. The re-
sulting 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 evapo-
ration there remains Ca2O2Cl43H2O. 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-:
CaCl2 + CaCl2O2 + 2Cl3 = 2CaCl2+2Cl2O. Chlorine monoxide may be prepared in thfo
manner.
It is sometimes said that bleaching powder contains a substance, Ca(OH)zCl2, that is
calcium peroxide, CaO2, in which one atom of oxygen is replaced by (OH)2, and the other
by C12 ; but, judging from what has been said above, this can only be the case in the dry
etate, and not in solutions.
On being kept for some time, bleaching powder sometimes decomposes, with the
evolution of oxygen (because CaClgOj = CaCLj + Os, see p. 10*5); the same takes place
when it is heated.
THE HALOGENS 479
powder (they are not obtained free from metallic chlorides), must be
counted as salts, because their metals are capable of substitution. But
the hydrate HC1O 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
H2CO3 or HNO3) into water and the anhydride, or chlorine monoxide.
C1.,O = 2HC1O — H2O ; and, in the second place, because, in a number
of instances, it evolves oxygen with great facility ^forming hydrochloric
acid : HC1O = HC1 + O. Both hypochlorous acid and chlorine
monoxide may be regarded as products of the metalepsis of water,
because HOH corresponds with C10H and C10C1. 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 -t- NaCIO
+ 3HC1 = 2NaCl + HC1 + C12 + H2O ; (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) j
oxygen is also disengaged on heating the dry salts — for instance,
NaCl + NaCIO = 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 CaCl2) 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 HC1O), 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
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 : 2C12+ CaCOs=CO2 + CaCl^+ C12O. From
this we may conclude that, although carbonic anhydride displaces hypochlorous anhy-
dride, it may be itself displaced by an excess of the latter.
480 PRINCIPLES OF CHEMISTRY
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 + 2C12 = Hg2OCl2 -f C12O. In this case, a
compound of mercury oxide with mercury chloride, or the so-called
mercury oxychloride, is obtained : Hg2OCl2 = HgO -f- HgCl2. 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
Chlorine monoxide, which corresponds to bleaching and hypo-
chlorous 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 con-
denses 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 C12O.33 The explosion may even take place spontaneously, and also
•"-' 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, MC10, be added to a solution of mercuric salt, HgX.j, mercuric oxide is liberated,
because the hypochlorite is decomposed.
ss 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.,3O4 + RjO + C12 = NaCl + HC1O + NaHSO4. Here
the hypochlorous acid is formed, together with HC1, at the expense of chlorine and
water, for Cle + H3O = HC1 + HC10. 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 fey 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 hypochlorpus acid — i.e. chlorine mon-
oxide— is disengaged.
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.
THE HALOGENS 481
in the presence of many oxidisable substances (for instance, sulphur,
organic compounds, <feo.), but the solution, although unstable and
showing a strong oxidising tendency, does not explode.34 It is evident
that the presence of hypochlorous acid, HC1O, may be assumed in an
aqueous solution of C120, since C120 + H2O = 2HC1O.
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
C12, NaCl, HC1, hydrochloric acid,
C12O, NaCIO, HC1O, hypochlorous acid
C12O3, NaClO2, HC1O2, chlorous acid.35
C12O5, NaClO3, HC1O3, chloric acid.
C12O7, NaCIO,, HC1O4, 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 MC1O, the other contains none at all.
3MC1O = MC103 + 2MC1
hypochlorite chlorate chloride
54 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 com-
bination with the unsaturated hydrocarbons is very often taken advantage of in organic
Chemistry. Thus its solution absorbs ethylene, forming the chlorhydriu C2H4C10H.
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.
54 Chlorous add, HC1O4 (according to the data given by Millon, Brandau, and
others) in many. respects resembles hypochlorous acid, HC10, 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, HNO2 The anhydride of chlorous acid, C12O5, is not known
in a pure state, but it probably occurs in admixture with chlorine dioxide, C1O3, which is
obtained by the action of nitric and sulphuric acids on a mixture of potassium chlorate
with such reducing substances as citric oxide, arsenious oxide, sugar, &c. All that is at
present known is that pure chlorine dioxide CIO? (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, NOa (giving HNO3 and HNO2), or as a mixed anhy-
dride, 2ClO2+H2O = HClOj+HClO.2. The silver salt, AgClO2, is sparingly soluble in
water. The investigations of Garzarolli-Thumlackh and others seem to show that the
nnhydride ClgOj does not exist in a free state.
482 PRINCIPLES OF CHEMISTRY
Part of the salt — namely, two-thirds of it — parts with its oxygen in
order to oxidise the remaining third.36 From an intermediate sub-
stance, RX, two extremes, R and RX3 are formed, just as nitrous
anhydride splits up into nitric oxide and nitric anhydride (or nitric
acid). The resulting salt, MC1O3, corresponds with chloric acid and
potassium chlorate, KC1O3. It is evident that a similar salt may bo
obtained directly by the action of chlorine on an alkali if its solu-
tion be heated, because RC1O will be first formed, and then RC1O3 ;
for example, 6KHO + 3C12 = KC1O3 + 5KC1 + 3H2O. Chlorates
are so prepared ; for instance, jjoiassium chlorate, which is easily
separated from potassium chloride, being sparingly soluble in cold
water.37
58 Hydrochloric acid, which is an example of compounds of this kind, is a satu-
rated 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 CjH8, three atoms of oxygen produce an alcohol, glycerin
or glycerol, CsH5(OH)j. 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,
HOC1 — 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, HClOj and HC1O4, 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 seemj» have the power of fortning 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, NaC104, NajSO,^ NaMnO4, Na^PO^, &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, KC1 + KC1O, 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
IB 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 ia
THE HALOGENS 483
If dilute sulphuric acid be added to a solution of potassium chlorate,
chloric acid is liberated, but it cannot be separated by distilla-
tion, as it is decomposed in the process. To obtain the free acid,
sulphuric acid must be added to a solution of barium chlorate.3* 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 : 4HC1O3 = 2HC1O4 + H2O + C12 + O3. 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 hydro-
chloric acid. Hydrochloric acid gas gives chlorine with chloric acid
cold water, especially in the presence of other salts. The double decomposition taking
place is Ca(ClOj).2 + 2K.C1 = CaCL + 2KC1O3. 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. KC1O3 is always formed by the action
of an electric current on a solution of KC1, especially at 80° (Haussermann and Naschold,
1894), BO that this method is now used on a large scale.
Potassium chlorate crystallises easily in large colourless tabular crystals. Its solu-
bility 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 0° = 28
parts, 20° = 85 parts, 40° = 40 parts, 100° = 57 parts; potassium perchlorate at 0° about
1. part, 20° about If part, lOO^ about 18 parts. When heated, potassium chlorate melts
(the melting point has been given as from S35°-876° ; 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 (J 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 rap'idly,. and
therefore bursts the guns, and it also has a very strong oxidising action on their metal.
The sodium salt, NaClO.i, 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; 0°=24 parts, 20° = 37 parts, 80° = 98 parts 6f salt
per 100 of water.
18 Barium chlorate, BafClOs^I^O, is prepared in the following way : impure chloric
acid is first prepared and saturated with baryta, and the barium salt purified by crystal-
lisation. The impure free chloric acid is obtained by converting the potassium in potas-
sium 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 silico-
fluoride are very sparingly soluble in water. Chloric acid is easily soluble in water.
484 PRINCIPLES OF CHEMISTRY
(and consequently with KC103 also) acting in the same manner as it
acts on the lower acids : HC103 + 5HC1 = 3H2O + 3C12.
By cautiously acting on potassium chlorate with sulphuric acid, the
dioxide (chloric peroxide), C1O2,39 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 C102.40 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 O2, and in many instances41
therefore it acts as an oxidising agent, although (like nitric peroxide)
it may itself be further oxidised.47 In dissolving in water or alkalis
chloric peroxide gives chlorous and hypochlorous acids — 2C1O2 + 2KHO
= KG103 4- KC1O2 -f H2O — and therefore, like nitric peroxide,
the dioxide may be regarded as an intermediate oxide between the
(unknown) anhydrides of chlorous and chloric acids : 4C102 = C12O3
+ C1206.'3
As the salts of chloric acid, HC1O8, are produced by the splitting
up of the salts of hypochlorous acid, so in the same way the salts of
36 To prepare C1O2 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: 3KC10j+2H2S04=2KHSO4
+ KC1O4+2C102+H2O. The reaction may result in an explosion. Calvert and Da vies
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 2KC1O5
+ 3C2Hj,O42H8O = 2C2HKO4 + 2CO2+ 2C1O8 + SttjO. The reaction is still further facili-
tated by the addition of a small quantity of sulphuric acid. If a solution of HC1 acts
upon KClOj at the ordinary temperature, a mixture of Cl^ and C102 is formed, but if the
temperature be raised to 80° the greater part of the C1O2 decomposes, and when passed
through a hot solution of MnCL it oxidises it. Gooch and Ereider 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 C12O4 would take place, as the reactions of ClOn point to
its being a mixed anhydride of HC1O2 and HClOj.
11 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.
*'J Potassium permanganate oxidises chlorine dioxide into chloric acid (Fiirst).
45 The euchlorine obtained by Davy by gently heating potassium chlorate with hydro-
chloric acid is (Pebal) a mixture of chlorine dioxide and free chlorine. The liquid and
gaseous chlorine oxide (Note 35), which Millon considered to be C12O5, probably contains
a mixture of ClO^ (vapour density 35), CljOs (whose vapour density should be 59), and
chlorine (vapour density 85'5), since its vapour density was determined to be about 40.
THE HALOGENS ji85
perchloric acid, HC104, are produced from the salts of chloric acid,
HC1O3. But this is the highest form of the oxidation of HC1. Perchloric
acid, HC1O4, 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 2KC103
=KC1O< + KC1 + O2.
•The formation of this salt is easily observed in the preparation
of oxygen from potassium chlorate, owing to the fact that the potas-
sium 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 HC104 45 is a colourless and exceedingly caustic substance
44 If a solution of chloric acid, HC1OS, 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: 4HC1O3=2HC1O4+ C12 + 3O + H2O. Roscoe accordingly
decomposed directly a solution .of potassium chlorate by hydrofluosilicic acid, decanted
it from the precipitate of potassium silicofluoride, K2SiF6, concentrated the solution of
chloric acid, and then distilled it, perchloric acid being then obtained (see following foot-
note). 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 distilla-
tion 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 composi-
tion HC1O4,2H2O 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 HC104,H2O which
melts at 50°. On carefully heating this hydrate it breaks up into perchloric acid, which
distils over below 100°, and into the liquid hydrate HC1O4,2H2O. The acid HC104 may
486 PRINCIPLES OF CHEMISTRY
which fumes in the air and has a specific gravity 1'78 at 15° (some-
times, after being kept for some time, it decomposes with a violent ex-
plosion). 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,
C1H04,H2O, separates out. This is much more stable, but the liquid
hydrate HC1O4,2H2O is still more so. The acid dissolves in water in
all proportions, and its solutions are distinguished for their stability.48
When ignited both, the acid and its salts are decomposed, with the
evolution of oxygen.47
also be obtained by adding one-fourth part of strong sulphuric acid to potassium chlorate,
carefully distilling and subjecting the crystals of the hydrate HC1O4,HSO obtained in
the distillate io a fresh distillation. Perchloric acid, HC1O4, itself does not distil, and
is decomposed on distillation until the more stable hydrate HC1O4,HKO ia formed ; this
decomposes into HC1O4 and HC1O4,2H2O, which latter hydrate- distils without decom-
position. This forms an excellent example of the influence of water on stability, and of
the property of chlorine of giving compounds of the type C1X7) of which all the above
hydrates, C1O3(OH), C1O.,(OH)S, and C10(OH)S> are members. Probably further re-
search will lead to the discovery of a hydrate C1(OH)7.
t*6 According to Eoscoe the specific gravity of perchloric acid = 1'782 and of the
hydrate HC1O4,H.2O in a liquid state (50°) 1*811 ; hence a considerable contraction takes
place in the combination of HC1O4 with HoO.
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 de-
composing, for example, lithium chlorate LiCIO3, 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, SLiC10s=2LiCl + LiQO4 + 6O, and that towards the end the
remaining salt is decomposed thus: 5LiC10Sr»4LiCl-tLiClO4 + 10O. The phenomena
observed by Potilitzin obliged him to admit that lithium perchlorate is capable, of decom-
posing 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 spon-
taneously, 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 re-
t2tions 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 ?.KC1O3 = KC1O4 + KC1 + O2, whilst the salt by itself evolves about half
as much oxygen, in accordance with the equation, 8KC1O3 = 5KC104 + 3KC1 + 2O2. The
decomposition of potassium perclilorate in'admixture with manganese peroxide proceeds
to completion, KC1O4 = KC1 + 2O2. But in decomposing by itself the salt at first give*
potassium chlorate, approximately according to the equation 7KC1O4^2KC1O3 + 5KC1
+ HO.j. Thus there is now no doubt that when potassium chlorate is heated, the per-
chlorate is formed, and that this salt, in decomposing with evolution of oxygen, again
gives the former salt.
In the decomposition of barium hypdchlorite, 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 accom-
panied by the formation of a certain quantity of the oxides and by the evolution of
THE HALOGENS 487
On comparing chlorine as an element not only with nitrogen and
carbon bat with all the other non- metallic elements (chlorine has so little
analogy with the metals that a comparison with them would be super-
fluous), w^ 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 HC1,
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 them-
selves.
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
chlorine, the chlorine being displaced by the oxygen disengaged. Spring and Pro~t
(1889) represent the evolution of oxygen from KC1O-, as due to the salt first splitting np
into base and anhydride, thus (1) 2MC1OS = M2O + C12O5 ; (2) CljOs = C12 + O5 ; and (8)
I may further remark that the decomposition of potassium chlorate as a reactiop
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.)i when the reaction is accompanied by the evolution of heat, or
when (for instance, H2 + L>, &c.) little heat is absorbed or evolved. In these cases it is
evident thc.t 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 a 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 be formed, but their
formation cannot affect the explanation of the phenomena.
*9
488 PRINCIPLES OF CHEMISTRY
on this fundamental point, if the grouping in one family be natural.47 Ws
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 C1O2 corresponds with N02 and C02, yet the last is the
highest oxide of carbon, whilst that of nitrogen is N2O5, and for chlorine,
if there were an anhydride of perchloric acid, its composition would
be C12O7, 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, HC1, HBr, HI.
Their oxygen compounds exhibit a similar analogy. Only fluorine
does not give any oxygen compounds. The iodine and bromine com-
pounds corresponding with HC103 and HC1O4 are HBr03 and HBr04>
HI03 and HIO4. 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 com-
paratively 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
« bu se6) for example the melting point of NaCl, NaBr, Nal 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 :
BC13 17° BBr5 90°
SiClj 59° SiBr4 153°
PClj 76° PBr3 175°
SbCl3 223° SbBr3 275°
BiCl3 447° BiBr3 453°
SnCl4 606° SnBr4 619°
ZnCl2 730° ZnBr, 650°
Thus for all the more volatile compounds the replacement of chlorine by bromine
raises the boiling point, but in the case of ZnXj it lowers it (Chapter XV. Note 19).
THE HALOGENS 489
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 oLcases.
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 NH3) and quinquivalent
to oxygen (it gives N206), and C being quadrivalent to both H and 0
as it forms CH4 and C02. 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 wkh 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,
P2 + H20 = 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 liused 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, CaF2, fluorine was also liberated ; but it was mixed
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 w» 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 + F2=CaF2 + O — that is, the calcium distributed itself between the
oxygen and fluorine. MnF4, which is capable of splitting up into MnF2 and F2) is without
doubt formed by the action of a strong solution of hydrofluoric acid on manganese per-
oxide, but under the action of vater 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.
490 PBINCIPLES OF CHEMISTRY
with chlorine, and it was impossible to study the properties of the
resultant gas. Brauner (1881) also obtained fluorine by igniting cerium
fluoride, 2CeF4 = 2CeF3 + F2 ; but this, like all preening 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 decom-
posed 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 indium 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, SiF4),
boron (forming BF3), sulphur, <fec. Its density (H = 1) is 18, so that
its molecule is F2. 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, com-
pletely 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
In 1894, Brauner obtained fluorine directly by igniting the easily
49 According to Moissan, fluorine is disengaged by the action of an electric current
on fused hydrogen potassium fluoride, KHF2. 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 compara-
'tiye study in relation to oxygen and chlorine. There is particular interest in the pheno-
mena 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 H2S, HC1, HBr, CS2, and CNH with a flash ; it does not act upon
Oo, N2( CO, and CO3 ; Mg, Al, Ag, and Ni, when heated, burn in it, as also do S, Se,P (forms
PF5) ; it reacts upon H2 even in the dark, with the evolution of S66'00 units of heat. At
a temperature of — 95°, F2 still retains its gaseous state. Soot and carbon in general (but
not the diamond) when heated in gaseous fluorine form fluoride of carbon, CF4
(Moissan, 1890) ; this compound is also formed at 300° by the double decomposition of
CC14 and AgF ; it is a gas which liquefies at 10° under a pressure of 5 atmospheres.
With an alcoholic solution of KHO, CF4 gives K,CO3, according to the equation CF4 + 6KHO
= K2C03+ 4KF+ SH-jO. CF4 is not soluble in water, but it is easily soluble in CCl, and
alcohol
THE HALOGENS 491
formed 49bl" double lead salt HF,3KF,PbF4, which first, at 230°,
decomposes with the evolution of HF, and then splits up forming
3KF,PbF2 and fluorine F2, which is recognised by the fact that it
liberates iodine from KI and easily combines with silicon, forming
SiF4. This method gives chemically pure fluorine, and is based upon
the breaking up of the higher compound — tetrafluoride of lead, PbF4,
corresponding to PbO2, into free fluorine, F2, and the lower more
stable form — bifluoride of lead, PbF2, which corresponds to PbO ; that
is, this method resembles the ordinary method of obtaining chlorine
by means of MnO2, as MnCl4 here breaks up into MnCl2 and
chlorine, just as PbF4 splits up into PbF2 and fluorine.
Among the compounds of fluorine, calcium fluoride, CaF2, is some-
what widely distributed in nature as fluor spar,™ whilst cryolite, or
aluminium sodium fluoride, Na,,AlF6, 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.
Hydrofluoric acid, HF, cannot be obtained from fluor spar in glass
retorts, because glass is acted on by and destroys the acid. It is
48 bu x. Nikolukin (1885) and subsequently Friedrich and Classen obtained PbCl*
and a double ammonium salt of tetrachloride of lead (starting from the binoxide),
PbCl.,2NH4Cl ; Hutchinson and Pallard obtained a similar salt of acetic acid (1893)
corresponding to PbX4 by treating red lead with strong acetic acid ; the composition of
this salt is Pb(C2H5O..,)4 ; it melts (and decomposes) at about 175°. Braoner (1894)
obtained a salt corresponding to tetrafluoride of lead, PbF4, and the acid corresponding
to it, HiPbFg. For example, by treating potassium plumbate (Chapter XVIII. Note 55)
with strong HF, and also the above-mentioned tetra-acetate with a solution of KHFj,
Brainier obtained crystalline HK3PbFg — i.e. the salt from which he obtained fluorine.
*° 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,
SiO2T 2CaF., = -2CaO + SiF4; 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 8'1. When steam is passed over incan-
descent fluor spar, lime and hydrofluoric acid are formed : CaFa + H/) = 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 pabses over to the
potassium or sodium, and the oxygen to the calcium. In solutions — for examplei
Ca(NOs)2 + 2KF = CAF2 ( precipitate) -f2KNOs (in solution) — the formation o! calcium
fluoride takes place, owing to its very sparing solubility. 26,000 parts of water dissolve
one part of fluor spar.
492 PRINCIPLES OF CHEMISTRY
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, CaF2 + H2SO4 = CaSO4 + 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 evo-
lution of a considerable amount of heat, and gives a solution of constant
boiling point which distils over at 1 20° ; showing that the acid is able to
combine with water. The specific gravity of the compound is 1'15, and
its composition HF,2H2O.52 With an excess of water a dilute solu-
tion 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 sub-
stances, but it acts in a highly corrosive manner on metals, glass, porce-
lain, and the majority of rock substances.53 It also attacks the skin,
41 According to Gore. Fremy obtained anhydrous hydrofluoric acid by decomposing
lead fluoride at a red heat, by hydrogen, or by heating the double salt HKF2, 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. .
48 This composition corresponds to the crystallo-hydrate HC1,'2H2O. All the proper-
ties of hydrofluoric acid recall those of hydrochloric acid, and therefore the com-
parative ease with which hydrofluoric acid is liquefied (it boils at +19°, hydrochloric
acid at — 85°) must be explained by a polymerisation taking place at low temperatures,
as will be afterwards explained, H2F.>, 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.
43 The corrosive action of hydrofluoric acid on glass and similar siliceous compounds
is based upon the fact that it acts on silica, SiOs, as we shall consider more fully in
describing that compound, forming gaseous silicon fluoride, SiO2 + 4HF = SiF4 + 2H.jO.
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 de-
stroyed. 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 com-
posed 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
THE HALOGENS 498
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, W 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 hydro-
chloric, hydrobromic, nor hydriodic acid acts on sand and glass, whilst
hydrofluoric acid corrodes them, forming gaseous silicon fluoride. The
othe'r halogen acids only form normal salts, KC1, NaCl, with Na or K,
whilst hydrofluoric acid gives acid salts, for instance HKF2 (and by
dissolving KF in liquid HF, KHF22HF 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 H2F2, as Mallet (1881) showed ; but a
depolymerisation occurs with a rise of temperature, and the density
approaches 10, which answers to the formula HF.64
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
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.
51 Mallet (1881) determined the density at 80° 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 Ht>F2 is probably connected
with the property of many fluorides of forming acid salts — for example, KHF2 and
HgSiFg. We saw above that HC1 has the same property (forming, for instance, H2PtC!6,
&c., p. 457), and hence this property of hydrofluoric acid does not stand isolated from the
properties of the other halogens.
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 com-
bination of the halogens with metals maybe made by throwing a small piece (a shaving)
ef 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
494 PRINCIPLES OF CHEMISTRY
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 hypo-
chlorous 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.
Bromine and iodine, like chlorine, occur in sea water in combina-
tion with metals. However, the amount of bromides, and especially
of iodides, in sea water is so small that their presence caa 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
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 (Thorn-
sen) for analogous actions of (1) chlorine, (2) bromine, and (8) iodine, with respect to
inetals ; 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 : —
128
K3 +X, 211 191 160
Na-j + Xj 196 172 188
Ag,,-fX,, 50 45 28
Hgg + X, 83 68 48
Hg -r-Xo 63 61 84
Ca + X, 170 141
Ba +Xa 195 170
Zn +X., 97 78 49
Pb +Xa 83 64 40
Al +X, 161 120 70
"We may remark that the latent heat of vaporisation of the molecular weight Br4is about
7'2, and of iodine 6'0 thousand heat units, whilst the latent heat of fusion of Br2 is about
0-8, and of Ia about 8'0 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 Znla would give
48 + 8+8, or about sixty thousand heat units, or 1£ times less than Zn + Cla.
M 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.
THE HALOGENS 495
mineral waters (those of Kreuznach and Staro-rossiisk) 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
chlorinej 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.
Bromine is a dark brown liquid, giving brown fumes, and having a
poisonous suffocating smell, whence its name (from the Greek /3/xo//.o5,
signifying evil smelling). The vapour density of bromine shows that
its molecule is Br2. 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
w 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: SKBr+KBrOj + eHjSO^
= CKHSO4 + 8HoO + 8Br2. 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 CaBr2.
59 There has long existed a difference of opinion as to the melting point of pure
bromine. By some investigators (Regnanlt, Pierre) it was given as between —7° and —8°,
and by others (Balard, Liebig, Quincke, Baumhauer) as between — 20° and — 26°. There
is how no doubt, thanks more especially to the researches of Ramsay and Toung (1885),
that pure bromine melts at about -7°, This figure is not only established by direct ex-
496 PRINCIPLES OF CHEMISTRY
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.
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
Fwus, Laminaria, &c. The fused ashes of these sea- weeds are called
* kelp ' in Scotland and ' varech ' in Normandy. A somewhat con-
siderable quantity of iodine is contained in these sea-weeds. After
periment (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 80 85 40 45mm.
t= -16°'6 -14° -12° -10° - 8-5° - 7°
For liquid bromine —
p= 50 100 200 400 600 760mm.
t»-6°-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 + 278) corresponding with equal
tension varies for every pair of substances in rectilinear proportion in dependence upon t,
*nd, therefore, for the above" pressure p, Ramsay and Young determined the ratio of
t + 278 for water and bromine, and found that the straight lines expressing these ratios
for liquid and solid bromine intersect also at 7°'0£ ; thus, for example, for solid bromine —
p= 20 25 80 85 40 45
273 + 1 = 256-4 259 261 268 264*6 266
278+*'= 295-8 299 802'1 804'8 807'2 809*8
c= 1152 1154 1-157 1159 1'161 1*168
where t' indicates the temperature of water corresponding with a vapour tension p, and
where c is the ratio of 278 + 1' to 278 + t. The magnitude of c is evidently expressed with
great accuracy by the straight line c = 1*1708 + 0*00111. In exactly the same way we find
the ratio for liquid bromine and water to be ct = 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 L
Note 49) on the temperature of the formation of ice f — 1°*115, with 1*891 gram of bromine
in 100 grams of water) in an aqueous solution of bromine, showed that bromine is contained
in solutions as the molecule Bra. Similar experiments conducted on iodine (Kloboukoff
1689 and Beckmann 1890) show that in solution the molecule is Lj.
6. Roozeboom investigated the hydrate of bromine as completely aa the hydrate of
chlorine (Notes 9, 10), The temperature of the complete decomposition of the hydrate it
+ 6°*2; the density of Br8 10H,O = 1*49.
THE HALOGENS 497
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 car-
bonate 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 combi-
nation 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
1 5 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
aalts.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
61 In general, SHI + O = Iz •+• H2O, 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.jOS) is a higher oxide, and ferrous oxide, FeO, a lower oxide ; the former
corresponds with FeX3, and the latter with FeX2, 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 CuX2, give compounds of the snboxide CuoO, or CuX. Even sulphuric acid,
which corresponds to the higher stage SO3, is able to act thus, forming the lower oxide
BO,. 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 hydri-
odic 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 ehe 11 again have occasion to
refer. In Chili, where a large quantity of iodine is extracted in the manufacture of Chili
nitre, which contains NalOj, it is mixed with the acid and normal sulphites of sodium
In solution ; the iodine is then precipitated according to the equation 2NaIOj + SNa^SOs
+ 2NaHSO3 = 5Na,,SO4 + 12 + H5O. The iodine thus obtained is purified by sublimation,
498 PRINCIPLES OF CHEMISTRY
apparatus in tabular crystals, having a black grey colour and metallic
lustre.62
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 tempera-
ture, and is of a violet colour, whence iodine receives its name (lotiSrp,
violet). The smell of iodine recalls the characteristic smell of hypo-
chlorous 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 3-^5-^ part of iodine. A
brown solution is thus obtained, which bleaches, but much more feebly
than bromine arid 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 bi-
sulphide, 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.
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
68 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).
65 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 £n many instances un-
stable highly iodised compounds, resembling crystallo-hydrates, have been obtained from
such solutions. Thus iodide of tetramethylammoninm, N(CHS)4I, combines with I2 and !«.
Even a solution of iodine in a saturated solution of potassium iodide presents indication*
of the formation of a definite compound Klj. Thus, an alcoholic 'solution of KIj does
not give up iodine to carbon bisulphide, although this solvent takes up iodine from an
alcoholic solution of iodine itself (GiraulttJbrgensen, and others). The instability of these
compounds resembles the instability of many crystallo-hydrates, for instance of HC1,2HSO.
THE HALOGENS 499
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 pro-
perties 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 A vogadro- Gerhard t'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 rati6 of their
atomic weights. Dividing the atomic weight of chlorine (35-5) by it»
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) 26.6«
The metallic bromides and iodides are in the majority of cases, in most
respects analogous to the corresponding chlorides^65 but chlorine displace*
'the bromine and iodiue 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 displace-
ment 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,
fl 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 substi-
tution 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 16, of sodium chloride 27, of sodium bromide 32, and of sodium iodide 41.
The volume of silicon chloroform, SiHClj, 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,2O has a sp. gr. (at 15°/4°) of 1'OIOG, consequently
(he volume of the solution 8,658'5/l:0106 = 8,620, hence the volume of sodium chloride in
solution =3,620 — 8,603 (this is the volume of 200 H^O) = 17, and in similar solutions,
NaBr = 26 and Nal = 85.
** 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
(s 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.
06 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 halogen*
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 propor-
tional to their equivalence. Thus, if NaCl + Br be taken, then out of 100 parts of
chlorine, 5'64 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 MCln be taken, it gives' with nBr a percentage sub-
stitution s= 4M/n2, where M is the atomic weight of the metal. This • law was deduced
500 PRINCIPLES OF CHEMISTRY
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, K2C03 + I2 = 2KI + CO2 + O, only the reac-
tions 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.
from observations on the chlorides of Li, K, Na, Ag (w = l), Ca, Sr, Ba, Co, Ni, Hg, Pb
(w = 2), Bi (« = 3), Sn (n = 4), and Fe2 (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 65.
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, 85-0, 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 sub-
stitution 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, KC1 12 p.o. and AgCl 4J 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 mag-
nitude 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 96 190
K 79-82 87 '4 88'22 94 '21
Na 88-68 90'74 91'70 95'4» —
That is, the- conversions which are accompanied by an evolution of heat proceed
with very much greater rapidity than the reverse conversions.
THE HALOGENS 501
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 par-
tially, whilst hydrogen iodide is decomposed by light even at the
ordinary temperature and very easily by a red heat. -Hence the reaction
I2 + H2 = HI -f 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 HC1, the conversion of 2HC1 into H2 + Cl^ requires the expenditure
" The dissociation of hydriodic acid has been studied in detail by Haute feuille and
Lemoine, from whose researches we extract the following information. The decom-
position 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, I7 + H,
= SHI — 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
decompcsit.on of hydriodic acid reaches the limit at 250° in several months, and at
440° ia 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.o.
If the pressure under which 2HI passes into H2 + I2 be 4£ atmospheres, then the limit is
84 p.c. ; under a pressure of % atmosphere the limit is 29 p.c. The small influence of
pressure on the dissociation of hydriodic acid (compared with N2O.,, Chapter VL Note 46)
is due to the fact that the reaction 2HI = I2 + H2 is not accompanied by a change of
volume. In order to show the influence of time, we will cite the following figures
nferring to 350°: (1) Reaction H2 + 12; after 8 hours, 88 p.c. of hydrogen remained free ;
8 hours, 69 p.c.; 84 hours, 48 p.c.; 76 hours, 29 p.c. ; and 827 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 H2 be added, then at the limit only half
as large a mass of HI is decomposed. Therefore, if an infinite mass of hydrogen b*
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 corre-
sponds 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 2H1 = I2 + H.2 ,
which are accompanied by slight heat effects, all foreign and feeble influences may strongly
affect the progress of the action (Note 47).
502 PRINCIPLES OF CHEMISTRY
of 44,000 heat units. The decomposition of 2HBr into H2 -f Br2
only requires, if the bromine be obtained in a gaseous state, a con-
sumption of about 24,000 units, whilst in the decomposition of 2HI
into H2 + I2 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, C1HO,
CrO3, A*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<,
68 The thermal determinations of Thomsen (at 18C) gave in thousands of calories,
C1 + H= +22, HCl + Aq (that is, on dissolving HC1 in a large amount of water) = +17'8,
and therefore H + C1 + Aq= +39'3. In taking molecules, all these figures must be
doubled. Br + H = + 8'4 ; HBr + Aq = 19'9 ; H + Br + Aq = + 28'8. According to Ber-
thelot 7'2 are required for the vaporisation of Br2, hence Br2 + H2 = 16'8 + 7'2= +24,
if Br2 be taken as vapour for comparison with C12. H + I=— 6'0, HI + Aq=19'2;
H + I + Aq= +13'2, and, according to Berthelot, the heat of fusion of I2 = 8'0,and of
vaporisation 6'0 thousand heat units, and therefore I2 + H2= — 2(6'0) + 8 + 6 = — 8'0, 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 thermo-
chemistry 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 + HC1, KHO + HI and Cl + H in aqueous solutions, it was possible to find H + 1 + Aq ;
then, knowing HI + Aq, to find I+H. It is evident that unavoidable errors may
accumulate.
& One can believe, however, on the basis of Berthollet's doctrine, and the obser-
vations of Potililzin (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 comprehen-
sible. 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 inves-
tigators, 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.
THE HALOGENS 503
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 decom-
pose 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 thermocheraicaL data,
for in the formation of HX in solution (in a large excess of water)
from the gaseous elements there is evolved forHCl 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.
In consequence of what has been said above, it follows that hydro-
bromic and hydriodic acids, being substances which are but slightly
stable, cannot be evolved in a gaseous state under many of those condi-
tions 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 hydri-
odic 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, H2SO4 + 2HI
= 2H2O + SO2 + 12. 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 SO2), in which case not only the affinity of hydriodic acid for
water is brought to light but also the action of water in directing chemi-
cal 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
hydriodio 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
70 On the basis of the data in Note 68.
71 A number of similar cases confirm what has been said in Chapter X.
71 This is prevented by the reducibiUty 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 hydriodio
acids only act feebly (like phosphoric acid), or do not act at all (like boric acid).
504 PRINCIPLES OF CHEMISTKY
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 BerthoHet'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
If hydrogen sulphide be passed through water containing iodine, the
reaction H2S + I2 = 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 + 4I2 + 9H20 (according to
Bineau) does not react with H2S, 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
To obtain 75 gaseous hydrobromic and hydriodic acids it is most
73 This is in agreement with the thermocheniical 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 2'8) we have iI2+S = 4;7; H2 + C13=44; H2 + Br2 = 24, and H2 + Lj=.
—8 thousand heat units ; hence the formation of H.2S gives less heat than that of HC1 and
HBr, but more than that of HI, In dilute solutions H2 + 6 + Aq = 9'3, and consequently
less than the formation of all the halogen acids, as H2S evolves but little heat with
•water, and therefore in dilute solutions chlorine, bromine, and iodine decompose
hydrogen sulphide.
" Here there are three elements, hydrogen, sulphur, and iodine, each pair of which,
is able to form a compound, HI, H.,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.
74 bis The same essentially takes place when sulphurous anhydride, in a dilute solu-
tion, 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 every-
where distinctly seen.
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 com-
plete than it now is it would be possible to foretell all cases of preparation with every
THE HALOGENS 605
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), hydro-
bromic acid is abundantly and uniformly disengaged.76 Hydrogen
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 men*
tioned in Note 74, and the • matter is here further complicated by the possibility of the
formation of the compound PH3 with HI, as well as the production of PI2, PI3, 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 ou 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. Bro-
mide 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) propose 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 con-
taining the bromine be kept at 60°, the hydrogen takes up almost the theoretical amount
lof 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.
Hydriodic acid is obtained hi the same manner as hydrobromic. The iodine is heated
in a small fiask, 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
506 PRINCIPLES OF CHEMISTRY
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, PI3 + 3H2O
= PH3O3 + SHI."
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
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
PIj, because otherwise a portion of the iodine distils over. If less than one-tenth part
of iodine be taken, much phosphonium iodide, PHjI, 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, wheu 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 PH4I and PI2 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 jaroceeds quietly, but nevertheless requires
care.
L. Meyer showed that with an excess of iodine the reaction proceeds without the
formation of bye-products (PH4I), according to the equation P + 51 + 4H2O = PH3O4 + 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 phosphoru0 and 10 grains 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).
78 The specific. gravities of their solutions as deduced by me on the! basis of Topsoe
and Berthelot's determinations for 15°/4° are as follows :—
10 20 SO 40 50 60 p.c,,
HBr 1-071 1-156 1'258 1-874 T505 1'650
HI 1-075 1-164 1-267 T399 1'567 1'769
Hydrobromic acid forms two hydrates, HBr,2H2O and HBr,H.jO, which have been
studied by Eoozeboom with as much completeness as the hydrate of hydrochloric acid
(Chapter X. Note 87).
"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.
THE HALOGENS 507
easy decoinposability 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 CJB^n+a or even the
saturated compounds. For example, benzene, CGH6, when heated in a
closed tube with a strong solution of hydriodic acid, gives hexylene,
C6HI2. 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 there-
fore, to obtain the products of iodine substitution, either iodic acid, H1O3
(Kekule), or mercury oxide, HgO (Weselsky), is added, as they imme-
diately react on the hydrogen iodide, thus : HIO3 + 5HI = 3H2O + 3I2,
or, HgO -f 2HI =.HgI2 + H2O. From these considerations it will
be readily understood that iodine acts like chlorine (or bromine)
oil ammonia and sodium hydroxide, for in these cases the hydriodic
acid produced forms NH4I and Nal. With tincture of iodine
or even the solid element, a solution of ammonia immediately forms
a highly-explosive solid black product of metalepsis, NHI2, 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 NAgI2) : 3NH3 + 2I2
= 2NH4I +NHI2. However, the composition of the last product is
variable, and with an excess of water NI3 seems to-be formed. Iodide
of nitrogen is just as explosive as nitrogen chloride.78 bis In the
78 bu Iodide of nitrogen, NHI2 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 NHI2. A ^y temperature facilitates its formation.
NHLj dissolves in ammonia water, and when heated the solution forms HIO3 and iodine.
With KI, iodide of nitrogen gives iodine, NH3 and KHO. These reactions (Selivanoff)
are explained by the formation of HIO from NHI2 + 2H2O - NH3 + 2HIO— and then
Kl + HIO = Io+KHO. Selivanoff (see Note 29) usually observed a temporary for-
mation of hypoiodous acid, HIO, in the reaction of ammonia upon iodine, so that
here the formation of NHI2 is preceded by that of HIO— i.e. first I2 + H2O = HIO + HI,
and then not only the HI combines with NH3, but also 2HIO + NH3=NHI2 + 2H2O.
With dilute sulphuric acid iodide of nitrogen (like NClj) forms hypoiodous acid, but it
immediately passes into iodic acid, as is expressed by the equation 6HIO = 2I2 + HIOj
+ 2H.,O (first 8HIO = HI03 + 2HI, and then HI + HIO=I3+H20). Moreover. Selivanoff
508 PRINCIPLES OF CHEMISTRY
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 + 3I2 = 5NaI + 3H2O + NaIO3
(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, HIO3 = I02(OH), shows the propensity of iodine to give
compounds of the type IX5. 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
found that iodide of nitrogen, NHI.j, 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 NHI2, 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, NO2C1). The latter are tolerably stable, while (the
haloid being designated by X) NHXj, NX3, XOH, RXO (see Chapter XIII. Note 48), &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.
78 Hypoiodous acid, HIO, is not known, but organic compounds, BIO, of this type
are known. To illustrate the peculiarities of their properties we will mention one
of these compounds, namely, iodosobenzol, CgHjIO. This substance was obtained
by Willgerodt (1892), and also by V. Meyer, Wachter, and Askenasy, by the action
of caustic alkalis upon phenoldiiodochloride, C^H.JLClg (according to the equation,
C6H5IC12 + 2MOH = CgHsIO + 2MC1 + H2O). lodosobenzol .is an amorphous yellow snb- .
stance, whose melting point could cot 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 C6H5IO, they give saline compounds in which iodosobenzol appears
as a basic oxide of a diatomic metal, C6H5I. 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 CgH5l(NO3)2 (like Ca(NO3)j).
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 HC1O. The action of peroxide of hydrogen, chromic acid, and other similar
oxidising agents gives iodoxybenzol, CgH5IO2, which is a neutral substance — i.e. incapable
of giving salts with acids (compare Chapter yTTT. Note 48).
79 bis The oxidation of iodine by strong nitric acid was discovered by Connell; Millon
bowed that it is effected, although more slowly, by the action of the hydrates of nitric
acid up to HNO3,H2O, but that the solution HNO3,2H2O, 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 combinet|
directly with iodine — for instance, at 0° forming the compound ISr4NH3 — whilst iodide of.
nitrogen is only formed in presence of water.
THE HALOGENS 509
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 iodio
acid. It is disengaged from solutions as a hydrate, HIO3> which loses
water at 170°, and gives an anhydride, I2O5. Both these substances
are crystalline (sp. gr. I2O5 5-037, HIO3 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 sul-
phurous anhydride, hydrogen sulphide, carbonic oxide, &c. — form
chloride of iodine and water with hydrochloric acid, and with bases
form salts, not only normal MIO3, but also acid ; for example,
KIO3HIO3, KIO32HI03.83 «• With hydriodic acid iodic acid imme-
diately reacts, disengaging iodine, HIO3 -f 5HI = 8H20 -f 3I2.
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, 5NaC103 + 8Br2 + 8H2O = SNaBr + 5HC1O3 + HBr03.
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 IHC12, or IC13, or HIO3. If there be a small amount
of water, then the iodic acid may separate out directly as crystals, but a complete con-
version (Bornemann) only occurs when not less than ten parts of water are taken to
one part of iodine— IC1 + 8H2O + 2C12=IHO3H- 5HC1.
88 Schonebein and Ogier proved this. Ogier found that at 45° ozone immediately
oxidises iodine vapour, forming first of all the oxide I<.j03, 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 Kiimmerer showed that a solution of sp. gr. 2'127 at 14°, containing 2HI03,9HjO,
solidified completely in the cold. On comparing solutions HI + mH2O with HIO3 + mH2O,
•we find that the specific gravity increases but the volume decreases, whilst in the
•passage of solutions HCl+mH2O to HC103+wiH2O both the specific gravity and the
volume increase, which is also observed in certain other cases (for example, H3PO3 and
H5P04).
83 b» Ditte (1890) obtained many iodates of great variety. A neutral salt, 2(LiI03)HaO,
is obtained by saturating a solution of lithia with iodic acid. There is an analogous
ammonium salt, 2(NH4IO3)HjO. He also obtained hydrates of a more complex com-
position, such as 6(NHJO3)H2O and 6(NH4lO3)2H2O. Salts of the alkaline earths,
Ba(I03)2H2O and Sr(IO3)2H2O, may be obtained by a reaction of double decompositioa
from the normal salts of the type 2(MeIO3)H2O. 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(IOx)*2H?O. An anhydrous salt is thrown out if nitric
510 PRINCIPLES OF CHEMISTRY
As with chlorine, so with iodine, a periodic acid, HIO4) 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
2H2O (corresponding with HC1O4,2H20), but as it forms salts con-
taining up to 5 atoms of metals, this water must be counted as water
of constitution. Therefore IO(OH)a = HIO4,2H2O corresponds with
the highest form of halogen compounds, IX7.85 In decomposing (at
acid be added to the solutions. Analogous salts of cadmium, silver, and copper give
compounds of the type 2MelO64NH3 and Me"(IO3)24NH3, 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(lO3),(NH4),O, Cd(IO3)2(NH4)2O. Copper gives Cn(IO3)24(NH4)jO and
Cu(IO5)j(XH4)2O. These salts may be regarded as compounds of I2O5, and MeO and
(NH4)3O ; for example, Zn(IO3)..j(NH4).;>O may be regarded as ZnO(NH4)2OI2O5, or, as
derived from the hydrate, IsO52HoO = 2(HIOj)H2O.
64 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 Na4I2O9,3HaO.
6NaHO + 2NaIO3 + 401 = 4NaCl + Na4I2O9 -t- 3H2O.
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, Ag4I2C>9,3H2O. If this sparingly
soluble silver compound be dissolved in hot nitric acid, orange crystals of a salt having
the composition AgIO4 separate on evaporation. This salt is formed from the preceding
by the nitric acid taking up silver oxide— Ag4I.2O9 + >2HNO3 = 2AgNO3 + 2AgIO4-fH.2O.
The silver salt is decomposed by water, with the re-formation of the preceding salt,
whilst iodic acid remains in solution —
4AgIO4 + H2O = Ag4I2O9 + 2HIO4.
The structure of the first of these salts, Na4I2O9,3H2O, 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 IOX5, or IX7l like AgIO4?which is IO3(OAg). The composition of all the
salts of periodic acids are expressed by this type EX7. Kimmins (1889) refers all
the salts of periodic acid to four types — the meta-salts of HIO4 (salts of Ag, Cu, Pb),
the meso-salts of H3IO5 (PbH, Ag2H, CdH),the para-salts of H5IO6 (Na2H3, Na3H2), and
the di-salts of H4I2O9 (K4, Ag4, Ni2). The three first are direct compounds of the type
IX7, namely, IO3(OH), IO2(OH)3, and IO(OH)5, and the last are types of diperiodic
Baits, which correspond with the type of the meso-salts, as pyrophosphoric salts corre-
spond with orthophosphoric salts — i.e. aHjiOj — H2O = H4I2O9.
85 Periodic acid, discovered by Magnus and Ammermiiller, and whose salts were
afterwards studied by Langlois, Kammelsberg, 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 HC1O,2H2O
the water, 2H20, is not displaced by bases, and must be regarded as water of crystallisa-
tion, whilst in HI04,2H2O it must be regarded as water of hydration. We shall after-
wards see that the system ol the elements obliges us to consider the halogens as
substances giving a highest saline type, GXi, where G signifies a halogen, and X oxygen
(O = X2), 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.
THE HALOGENS 611
200°) or acting as an oxidiser, periodic acid first gives iodic acid, but it
may also be ultimately decomposed.
Compounds formed between chlorine and iodine must be classed
among the most interesting halogen bodies.86 These elements com-
bine together directly with evolution of heat, and form iodine
monochloride; 1C1, or iodine trichloride, IC13.87 As water reacts on
these substances, forming iodic acid and iodine, they have to be pre-
pared 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
88 With respect to hydrogen, oxygen, chlorine, and other elements, bromine occupies
on 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 respect-
ing iodine monochloride, IC1, 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 (1898) showed that of the two modifications of IC1, one is stable, and melts at
27° ; while the other, which easily passes into the first, and is formed in the absence of IClj,
melts at 14°. Schiitzenberger 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 deliques-
cent octahedra, having the composition IC14, which are therefore formed from the mono-
chloride 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 : IC15 + 6KHO
= DKCl + KIOg + 8H20 The existence -of a pentachloride IC15 is, however, denied,
because this substance has not been obtained in a free state.
Stortenbeker (1888) investigated the equilibrium of the system containing the mole-
cules I2, IC1, IC1S, and C12, in the same way that Roozeboom (Chapter X. Note 88) examined
the equilibrium of the molecules HC1, HC1,2H2O, and H2O. He found that iodine
mouochloride 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 18°'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 IHC19
is obtained, which does not seem to be altered by water. Besides this compound, iodine
and iodic acid are always formed, 10IC1 + 3H2O = HIOS + 6IHC12 + 2I2 ; and in this respect
iodine trichloride may be regarded as a mixture, IC1 + IC14 = 2IC13, but IClj + SH.jO
= IHO3 + 6HC1; hence iodic acid, iodine, the compound IHClj, and hydrochloric acid
are also formed by the action of water.
*10
612 PRINCIPLES OF CHEMISTRY
crystals of the monochloride, it is easy to obtain iodine trichloride ia
orange crystals, which melt at 34° and volatilise at 47°, but in so doing
decompose (into C12 and C1I). 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, C2H4), which
are capable of directly combining with chlorine aad iodine, also directly
combine with iodine monochloridfe.
518
THE neutral salt, sodium sulphate, Na2SO4, 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
put 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 un-
changed during this reaction, PbS04 + 2NaCl =PbCl2 + NaaS04, the
vapours will contain lead chloride, and the residue will contain the mix-
ture 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 PbCl2 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.
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 uugeneralised properties and relations.
614 PRINCIPLES OF CHEMISTRY
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, Ag2S04, in solution with common
salt, gives silver .chloride, because the latter is insoluble- in water,
Ag.2SO4 + 2NaCl = Na2S04 + 2AgCl. Sodium carbonate, mixed in
solution with the sulphates of iron, copper, manganese, magnesium, <fec.,
gives in solution sodium sulphate, and in the precipitate a carbonate
of the corresponding metal, because these salts of carbonic acid
ore insoluble in water; for instance, MgSO4 -f Na2C03 = Na2S04
+ MgCO3. 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,
CuSO4 + 2NaHO = Cu(HO)2 + Na2S04. Sulphate of magnesium,
MgS04, on being mixed in solution with common salt, forms, although
not completely, chloride of magnesium, and sodium sulphate. On cool-
ing the mixture of such (concentrated) solutions sodium sulphate is
deposited, as was shown in Chapter X. This is made use of for prepar-
ing it on the large scale in works where sea-water is treated. In this
case, on cooling, the reaction 2NaCl + MgSO4 = MgCl2 + Na2S04
takes place.
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,
Na2SO4,10H20, 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.
The methods of obtaining salts by means of double decomposition
2 Anhydrous (ignited) sodium sulphate, Na4S04, is known in trade as ' sulphate '
or salt-cake, in mineralogy thenardite. Crystalline decahydrated salt is termed in
mineralogy miralilite, and in trade Glauber's salt. On fusing it, the monohydrate
is obtained, together with a supersaturated solution.
SODIUM 515
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 possi-
bility of this occurrence ought to be foreseen according to Bertho) let'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 sepa-
rating them from other salts have not been generalised.4 These pro-
perties as yet remain subjects for investigation, and are rarely to
be foreseen. The crystallo- hydrate of the normal sodium sulphate,
Na2S04,10H2O, very easily parts with water, and may be obtained
in an anhydrous state if it be carefully heated until the weight re-
mains 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 SO3. 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 Na2SO4)10H2O.5 At 34° the latter fuses, and the
solubility decreases at higher temperatures.9 A concentrated solution
at 34° has a composition nearly approacnmg to Na2SO4 + 14H20,
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, NagS, or sodium sulphite,
Na^SOs, &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.2SO4, is obtained from NaCl without the
preliminary manufacture of ILjSO^. 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 SOj. Under
these conditions, HC1, 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 decomposi-
tion 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
84 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 deter-
minations of Gay-Lussac, Lovell, and Mulder.
« In Chapter L, Note 94, we have already seen that with many other sulphates the
616 PRINCIPLES OF CHEMISTRY
and the decahydrated salt contains 78 -9 of the anhydrous salt com-
bined 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, C12,8H2O (Chapter XI., Note 10). Not only the fused deca-
hydrated salt, but also the concentrated solution at 34° (not all at once,
but gradually), yields the monohydrated salt, Na2SO4,H2O. The hepta-
hydrated salt, Na2SO4,7H2O, 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 solu-
tion, and at 20° 52'8 parts. Hence, with respect to the anhydrous and
monohydrated salts, the solubility is identical, and falls with increas-
ing temperature, whilst with respect to decahydrated salt, the solubility
rises with increasing temperature. So that if in contact with a solution
solubility also decreases after a certain temperature is passed. Gypsum, CaS04,2H2O, lime,
and mariy other compounds present such a phenomenon. An observation of Tilden's (1884)
io 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 48 parts of anhydrous salt, at 140° 42 parts, at 160° 43 parts, at
180° 44 parts, at 230° 46 parts. According to Etard (1892) the solubility of 30 parts of
NajSO4 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, deter-
mined 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 com-
pounds, many useful statements can hardly be given.
' Already referred to in Chapter I., Note 56.
The example of sodium sulphate is historically very important for the theory of solu-
tions. 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 thesp. gr. of solutions at 15°/4°= 9,992 + 90'2p + 0'35ps
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.
= ri986 ; 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 88'8 to 87'6) ; but in the same way 83'S c.c. of 20 p.c. solution
we formed from (45-4/l'46 = ) 81'1 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.
SODIUM 517
of sodium sulphate there are only crystals of that heptahydrated salt
(Chapter I., Note 54), Na2SO4,7H2O, which is formed from saturated
solutions, then saturation sets in when the solution has the follow-
ing 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 raonohyd rated salt and a saturated solution. Thus sodium
sulphate has three curves of solubility : one for Na2SO4,7H2O (from 0'
to 26°), one for Na2S04,10H2O (from 0° to 34°), and one for
Na2SO4,H2O (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
Thus solutions of sodium sulphate may give crystallo-hydrates of
three kinds on cooling the saturated solution : the unstable hepta-
hydrated 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 hepta-
hydrated salt decomposes into them, probably according to the equa-
tion 3Na2SO4,7H2O = 2Na2S04,10H2O + Na2SO4,H2O. 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
Sodium sulphate, Na2SO4, only enters into, a few reactions of com-
bination with other salts, and chiefly with salts of the same acid
forming double sulphates. Thus, for example, if a solution of sodium
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 CaCls).
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° -27 5, at 20°
gives out +25, at 25° + 800 calories. For the decahydrated salt, Na2SO4,10H2O,
5°-4,225, 10°-4,000, 15°-8,570, 20°-3,160, 25°-2,775. Hence (just as in Chapter I.,
Note 56) the heat of the combination Na^SO^lOHaO at 5° =+3,125, 10° =+3,250,
20°= +8,200, and 25°= +3,050.
It is evident that the decahydrated salt dissolving in water gives a decrease of tempera-
ture. 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 NaoSO^lOHjO 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 form-id.
518 PRINCIPLES OF CHEMISTRY
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, Na2SO4 + H2SO4
= 2NaHSO4. This separates from hot solutions, whilst the crystallo-
hydrate, NaHS04,H2O,10 separates from cold solutions. The crystals
when exposed to damp air decompose into H2S04, which deliquesces,
and Na2SO4 (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, Na2S.207) at a low red
heat.11 bl» This anhydrous salt, at a bright red heat, parts with the
elements of sulphuric anhydride, the normal sodium sulphate remaining
behind — Na2S2O7 = Na2SO4 + SO3. 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.
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, Na2SO4 -» Ba(HO)a = BaSO4 + 2NaHO.
And by taking any salt of barium, BaX2, the corresponding salt of
sodium may be obtained, Na2S04 + BaX2 = BaSO4 + 2NaX. Barium
10 The yery large and well-formed crystals -of this salt resemble the -hydrate
H2SO4,HjO, 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 + 77'92y + 0'231p8 (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, HjSO4 + 25H20 with No.jSO4 + 25H3O increases from 488 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 NaHSO4,H2SO4 are separated (Schultz, 1868). This
compound fuses at about 100° ; the ordinary acid salt, NaHSO4, at 149°.
11 bis On decreasing the pressure, sodium hydrogen sulphate, NaH8O4, dissociates
much more easily than at the ordinary pressure ; it loses water and forms the pyrosul-
phate, NtujSaO7 ; this reaction is utilised in chemical works.
SODIUM 519
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 S03 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, Na2SO3 (for instance, in the formation
of glass) ; or else the decomposition proceeds further, and sodium
sulphide, Na2S, is formed, according to the equation Na2SO4 + 20
= 2CO2 + Na2S.
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, Na2CO3, 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 Na2SO4,
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
520 PRINCIPLES OF CHEMISTRY
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,SO4 -t- 2C = Na2S + 2CO2.
The sodium sulphide thus formed then enters into double decomposition
with the calcium carbonate taken, and gives calcium sulphide and
sodium carbonate, Na2S + CaCO3 = Na2CO3 + CaS.
Besides which, under the action of the heat, a portion of the excess
of calcium carbonate is decomposed into lime and carbonic anhydride,
CaCO3 = CaO + CO2, and the carbonic anhydride with the excess of
charcoal forms carbon monoxide, which towards the end of the opera-
tion 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 formulae, if it be con-
sidered that the product contains one equivalent of calcium oxide to
two equivalents of calcium sulphide.18 The sum of the reactions
may then be expressed thus : 2Na2SO4 + 3CaCO3 + 9C = 2Na2CO3
-+• CaO,2CaS -f 10CO Indeed, the quantities in which the substances
are mixed together at chemical works approaches to the proportion re-
quired 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
18 Calcium sulphide, CaS, like many metallic sulphides which are soluble in water, ia
decomposed by it (Chapter X.), CaS + H^O = CaO + HjS, 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 + 2Ca8
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 manu-
facture 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.
SODIUM
521
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
fio. 68.— Reverberatory furnace for the manufacture of sodium carbonate. P, grate. A, bridge.
M, hearth for the ultimate calcination of the mixture of sodium sulphate, coal, and calcium car-
bonate, wnich 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.
partially fused mass obtained at the end of the process is cooled, and
then subjected to methodical lixiviation l3 to extract the sodium car-
is 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 writer. This is easily effected if water be first poured
PlG. 69. — Apparatus for the methodical lixiriation of black ash, *c. Water flows into the tanks
from the pipes r, r, aud the saturated liquid is drawn off from c. c.
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
522 PRINCIPLES OF CHEMISTRY
bonate, the mixture of calcium oxide and sulphide forming the so-called
' soda waste ' or ' alkali waste.' u
which has already been subjected to a first lixiviatiou 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, conshting 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 con-
taining 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 ie
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°
Jn a current of air, in order to promote the separation of the water vapour. The large
transparent crystals (efflorescent in air) of NagCOjjlOHjO 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
iflecomposing 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. T21) of magnesium
chloride, which disengages sulphuretted hydrogen: CaS + MgCl2 + 2H2O = CaCl2
+ Mg(OH)2 + H2S. Sulphurous anhydride is passed through the residue in order to form
the insoluble calcium sulphite: CaCl.j + Mg(OH);j + SOs = CaSOs+MgCU + HtfO. The
SODIUM 623
The above-mentioned process for making soda was discovered in the
year 1808 by the French doctor Leblanc, and is known as the Leblano
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 XL, page 497). Even up to the present time this process is
carried on in Normandy. In France, where for a long time the manu-
facture 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.16
eolation 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 :
CaSO3 + 2H5S+2HC1 = CaCl, + 3H,0 + 89.
But most efforts have been directed towards avoiding the formation of soda
waste.
15 Among the drawbacks of the Leblauc 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 Leblauc process,
besides its simplicity and cheapness, are that almost the whole of the acids obtained
us 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 chloiide is decom-
posed 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 reverbe-
ratory furnaces. Here a. compound, Nagl^Sj, is formed, which is insoluble in waten
absorbs oxygen and carbonic anhydride, and then forms sodium carbonate and ferroa/l
sulphide ; this when roasted gives sulphurous anhydride, the indispensable materiaj
524 PRINCIPLES OF CHEMISTKY
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, NH4HCO3. 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,
NaHCO3, is precipitated and ammonium chloride, NH4C1, 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 NaHCO3 according to the equation NaCl + NH4HCO3
e= NH4C1 -f- NaHCO3. 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 beating. It is thus
obtained in a very pure state.17
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.2S from
NooSO4 (by heating it with carbon), dissolves it in water and subjects the solution to the
action of an excess of CO2 in coke towers, thus obtaining H2S (a gas which gives SO]
under perfect combustion, or sulphur when incompletely burnt, Chapter XX., Note 6) and
bicarbonate of sodium ; NajS + 2CO2 + 2H2O = HSS + 2HNaCOs. The latter gives soda and
CO2 when ignited. This process quite eliminates the formation of soda-waste (see Note 8)
and should in my opinion be suitable for the treatment of native Na2SO4, like that
which is found in the Caucasus, all the more since H2S 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 28 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 tho
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.
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 Bussia, 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 bu Mond (see Chapter XI., Note 8 bis) separates the NH^Cl from the residual solu-
tions by cooling (Chapter X., Note 44) ; ignites the sal-ammoniac and passes the vapour
over MgO, and so re-obtains the NH3 and forms MgCl : the former goes back for the
manufacture of soda, while the latter is employed either for making HC1 or CL2.
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
SODIUM 626
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 super-
saturated solution, and, according to the conditions, gives various com -
binations with water of crystallisation (mentioned on page 108), <fec.
A.t a red heat superheated steam liberates carbonic anhydride from
sodium carbonate and forms caustic soda, Na2CO3 + H2O = 2NaHO
+ CO2. 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
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 (Lb'wel, 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 82 parts (of anhydrous salt) in 100
water ; the other is more stable, and its solubility 20 parts (of anhydrous salt) per 1QO of
water. The solubility of the decahydrated salt in 100 water = at 0°, 7*0 ; at 20°, 91-7 ; at
80°, 87-2 parts (of anhydrous salt). At 80° the solubility is only 46-1, at 90° 45*7. at )00°,
46-4 parts (of anhydrous salt). That is, it falls as the temperature rises, like NajSO4.
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 + 104'5p
+ 0'166j>(. 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.o. 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= ffa •*• Ap + Bp*,
where S0 is the specific gravity of the water. Por 6 p.c. the sp. gr. 15°/4° = lg0520 ; for 10
p.c. 1*1057; for 15 p.c. 1*1608. The changes in the sp. gr. with the- temperature are
b$re almost £he same as with solutions of sodium chloride with an equal value of p
18 The resemblance is so great that, notwithstanding the difference in the molecular
composition of No^SOi an<l Na^COj, they ought to be classed under the type (NaO)aR,
where R = 8O2 or CO. Many other sodium salts also contain 10 mol. H2O.
19 According to the observations of Pickering. According to Rose, when solutions of
•odium carbonate are boiled a-oertain amount of carbonic anhydride is disengaged.
526 PBINCIPLES OF CHEMISTRY
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 in-
stance, with barium salts it precipitates an insoluble barium car-
bonate (Ba012 + Na2C03 = 2NaCl + BaCO3) and with the aluminium
salts it precipitates aluminium Tiydroxide, carbonic anhydride being
disengaged 3Na2CO3 + A12(SO4)3 + 3H2O = 3Na2S04 + 2A1(OH)3
-f 3CO2. 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), NaHCO3.20 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
*° The composition of this salt, however, may be also represented _as a combination
of carbonic acid, H2CO3, with the normal salt, NajCOs, just as the latter also com-
"bines with water. Such a combination is all the more likely because (1) there exists
another salt, Na2COs,2NaHCOj,2H20 (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 NaHCOj,Na2CO3,2H2O may be
easily obtained, as long as the temperature is above 86°. The natural urao (Boussinganlt)
lias, according to Laurent, the same composition. This salt it 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 astrachanite (Chapter XTV., 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 com-
pounds with water) the water of crystallisation of the normal salt separates and the water
id, 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 NagCOjHjO, this salt
•will, when powdered, absorb CO2 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
c4'8 ports of the anhydrous normal salt, but at 0° 100 parts of water dissolve 7 parts of
SODIUM 527
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 anhy-
dride.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, MgCO3. No such precipitate is formed by the acid salt,
because magnesium carbonate is soluble in the presence of an excess of
carbonic anhydride.
Sodium carbonate is used for the preparation of caustic soda23 —
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,*4
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 80° 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 80° of 27 parts. The solu-
bility therefore increases very rapidly with the temperature. And its saturated solution
is more stable than a solution of sodium bicarbonate. In fact, saturated solutions oi
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 ammo-
nium salt.
22 Crystalline sodium carbonate (broken into lumps) also absorbs carbonic anhydride,
but the water contained in the crystals is then disengaged: Na-jCOjjlOHjO + COj
= NajCOg, HjCOj + 9H2O, 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 Icfeenges like those
made at Essentuki, Vichy, &c.
e In chemistry, sodium oxide is termed ' soda,' which word must be carefully distin-
guished from the word sodium, meaning the metal.
** 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 in-
529 PRINCIPLES OF CHEMISTEY
and boiled in a cast-iron, wrought-irou, or silver boiler (sodium hydrox-
ide 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 Na2C03 + Ca(HO)a = CaCO3 + 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
fluence 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 con-
firmed by further investigation.
** As long as any undecomposed sodium carbonate remains in solution, excess
of acid added to the solution disengages carbonic anhydride, and the solution after dilu-
tion 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 hydrox-
ide 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 ha?
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
Ly evaporation (Note 14). If excess of lime and charcoal have been used, much sodium
hydroxide may-be obtained. After the removal as much as possible of the sodium carbon-
ate, a red liquid (from iron oxide) is left, containing sodium hydroxide mixed with com-
pounds 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 £nd from iron. The greater
part of the caustic soda, which forms so important an article of commerce, is manufactured
in this manner.
M Lb'wig 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 decom-
position with the ferric oxide, and the ferric carbonate produced is decomposed into
carbonic anhydride and ferric oxide, the action of which is renewed. Similar explana-
SODIUM 529
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,7H3O; they
fuse at + 6°.27 If the evaporation be conducted so long as water is dis-
- engaged, 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 disen-
gagement 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 1 30° ; at 55° water
dissolves an equal weight of it.32 Caustic soda is not only soluble in
tions expressing the reason for a reaction really add but little to that elementary con-
ception 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 COj move in a circle round the elements Na2O, but at the points of contact
with Fe^Oj the motion becomes elliptic with a long axis, and at some distance.f rom NogO
the elements of CO2 are parted, not having the faculty of attaching themselves to Fe2Os.
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,7HiO 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, 6, and 7 H2O.
K 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.
89 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.
80 As the molecular weight of NaHO = 40, the volume of its molecule =40/2-18 = 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.
81 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 + nHjO, on being mixed with water, evolve heat if they contain less than 6H2O,
but if more they absorb heat.
!* The specific gravity of solutions of sodium hydroxide at 16°/4° is given in the jghort
table below : —
NaHO, p.o. .6 10 15 SO 30 40
Sp..gr. . . . 1-057 1-118 1-169 1'224 1-881 1-486
530 PEINCIPLES OF CHEMISTRY
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.
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 poison-
ous 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
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).
S3 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, CjHs(OH)., (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
onion with the glycerin in the fat, as Chevrenl 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, C16HS2O > and
CigHagO-j, and the liquid oleic acid, CjsH^Oj. 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 Ota 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 akin 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 (tee Chapter IX., Note 15).
54 On this is founded the process of Henkoff and Engelhardt for treating bones
SODIUM 581
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, <fec. 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(NO3)2 + 2NaHO = Cu(HO)2 + 2NaNO3. 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 pheno-
menon 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 form-
ing 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 phos-
phide. When the hydrogen compound disengaged is capable of combin-
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.
55 As explained in Note 3&.
582 PBINCIPLES OF CHEMISTRY
ing 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, C12, takes the place
of the hydrogen, and forms the hypochlorite, NaCIO 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.
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, NaHSO4, NaHCO3), 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 irith poly basic
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,
SODIUM 583
H2CO3,' or sulphuric, H2S04 — 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,
NaHS04, and Na2S04, 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, CaCO3
and CaSO4.35M»
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 dis-
coveries in che.mistry, 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 con-
ductivity), 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 demon-
strated the compound nature of alkalis. On being decomposed by the
galvanic current, caustic soda disengages hydrogen and sodium at the
» Wi 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, H5PO4 — 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 atomio weight ; the feebler bases (like silver oxide), although corresponding with
univalent metals, do not form acid salts, while the feeblest bases {CuO, FbO, &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 ex-
pected 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
M the faculty of forming' compounds with hydrogen is characteristic of elements.
534 PRINCIPLES OF CHEMISTRY
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 The"nard, who observed that metallic iron at a high
temperature was capable of reducing caustic soda to sodium.'6 Brunner
latterly discovered that not only iron, but also charcoal, has this
property, although hydrogen has not.37 But still the methods of ex-
tracting 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. Further-
more, in practice the most important epoch in the history of the
production of sodium is comprised in the investigation of Sainte-CIaire
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.
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
56 Deville supposes that each 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 dis-
sociation 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 follow-
ing 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 fee 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
Na30 ; then, however, only one half is displaced, with the formation of NaHO
SODIUM 585
from the charcoal.38 The chalk on being heated loses carbonic anhy-
dride, leaiving infusible lime, which is permeated by the sodium
carbonate and forms a thick mass, in which the charcoal is intimately
mixed with the sodiunf 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 :
Na2CO3 + 20 = Na2 + 300
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.
In sodium works an iron tube, about a metre long and a decimeter
KC^WWVVVON'X V'>^NNVV^VVSNV^ VO>XVV-VC1CW^>^NNNV ^ N-^V>,-^X w^^xx\^-v^" ^> X-S.-.-WWW..-
FlO. 70.— Manufacture of sodium by Deville's process. A C, iron tube containing a mixture of
soda,, charcoal, and chalk. "B. condenser.
soda, charcoal, and chalk. B, condenser.
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
58 Since the close of the eighties in England, where the preparation of sodium is at
present carried out on a large commercial scale (from 18CO 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 carburettcd 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$ parts of 'metallic sodium. The reaction proceeds
more easily than with carbon or iron alone, and the decomposition o the NaHO proceeds
according to the equation: 8NaHO + C = Na2CO3 + 8H+Na. Subsequently, in 1891,
aluminium was prepared .by electrolysis (see Chapter XVII.), and metallic sodium found
•11
536 PRINCIPLES OF CHEMISTRY
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 in-
troduced into the aperture C, compelling the
vapours arid gases formed to pass through the
condenser B. This condenser consists of two
OTtr2SSB0S2; 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 de-
composes it, absorbing oxygen from it. Although the above described
methods of preparing sodium by chemical means have proved very con-
venient 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
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).
SODIUM 537
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 tempera-
ture 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, 1889V 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 bls It
forms alloys with most metals, combining with them, heat being some-
times 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.
It is easy to form an alloy of mercury and sodium having a crystal-
line structure, and a definite atomic composition, NaHg5. The alloy of
sodium with hydrogen or sodium hydride, Na2H, which has the external
J8 bU 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).
59 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 NaHg5, the amount of heat evolved increases,
after which the heat of formation falls, and the heat evolved decreases. In the formation
of NaHg5 about 18,500 calories are evolved ; when NaHgj is formed, about 14,000 ; and
for NaHg about 10,000 calories. Kraft regarded the definite crystalline amalgam as
having the composition of NaHgg, but at the present time, in accordance with Grimaldi's
results, it is thought to be NaHgs. A similar amalgam is very easily obtained if ft
8 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, KHg10. It may be mentioned here that the latent heat of fusion (of atomio
quantities) of Hg = 860(Personne), Na = 780 (Joannis), and K = 610 calories (Joannis).
40 Alloys are so similar to solutions (exhibiting such complete parallelism in proper-
ties) that they are included in the same class of so-called indefinite compounds. But la
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 witb
tin (Bailey 1892 found NajSn), lead (NaPb), bismuth (Na3Bi), &c. (Joannis 1892 and others)
have been investigated.
588 PRINCIPLES OF CHEMISTRY
appearance of a -metal,41 is a most instructive example of the character-
istics 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 ten>
peratures), 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 ia
as soft as this latter, when heated it becomes brittle, and decomposes
above 300°, evolving hydrogen. In this decomposition all the pheno-
mena of dissociation are very clearly shown — that is, the hydrogen
gas evolved has a definite tension 4a 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 4* we see other properties
than in such hydrogen compounds as HC1, H2O, H3N, H4C, or even
in the gaseous metallic hydrides AsH3, TeH2. 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.)*3bis
41 Potassium forms a similar compound, but lithium, under the same circumstances,
does not.
** The tension, of dissociation of hydrogen p, in millimetres of mercury, is: —
t = 830° 850° 400° 480°
for NajH p - 28 57 447 910
forKjH 46 72 548 1100
42 In general, during the formation of alloys the volumes change very slightly, and
therefore from the volume of NojH 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. Prom the fact that
the density of Na2H 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 con-
tained in the Na^ (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^H is equal-to 1'6 c.c., and conse-
quently 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.
« i>u We may remark that at low temperatures Na absorbs NHg and forma (NHsNa)z
(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 doea).
SODIUM 539
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.2Na, although it
does not displace hydrogen from the hydrocarbons.45 Sodium burns
44 H. A. Schmidt remarked that perfectly dry hydrogen chloride is decomposed
with great difficulty by sodium, although the decomposition proceeds easily with potas-
sium 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 bu Sodamide, NH2Na, (Chapter IV., Note 14), discovered by Gay-Lussac and
Thenard, 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 Thenard obtained an olive-
green, easily-fusible mass, soflamide, NH2Na, hydrogen being separated. This substance
with water forms sodium hydroxide and ammonia ; with carbonic oxide, CO, it forms
sodium cyanide, NaCN, and water, H20 ; 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, NH2Na,
only partially decomposes, with evolution of hydrogen, the principal part of it giving
Ammonia and sodium nitride, Na3N, according to the equation 8NH2Na=2NH3 + NNas.
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 800°-4000, 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 NH3. Pure NH2Na 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 NH2Na. It partially volatilises (sublimes) in vacuo at 200°, and breaks
up into 2Na + N2 + 2H2 at 600°. The same amide is formed when oxide of sodium is
heated in NH5: Na2O + 2NH3=2NaH2N + H2O. NaHO is also formed to some extent
by the resultant H2O. Potassium and lithium form similar amides. With water,
alcohol, and acids, NH2Na gives NH3 and NaHO, which react further. Anhydrous CaO
absorbs NH2Na when heated without decomposing it. When sodamide is heated with
SiO^, NH3 is disengaged, and silicon nitride formed. It acts still more readily upon
boric anhydride when heated with it: 2NH2Na + B.2O3 = 2BN + 2NaHO + H2O. When
slightly heated, NH2Na + NOC1 = NaCl + N2 + H2O (NHNa2 and NNa3 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 de-
struction of the organic groups and production of carbon.
4i 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
640 PRINCIPLES OF CHEMISTKY
"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
Wjth oxygen sodium unites in three degrees of combination, forming
a suboxide Na4O,46 bis an oxide, Na2O, and a peroxide, NaO. They are
thus termed because Na2O is a stable basic oxide (with water it forms
a basic hydroxide), whilst Na4O and NaO do not form corresponding
saline hydrates and salts. The suboxide is a grey intiammable 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 : Na40 + 3H2O = 4NaHO + H247
Oxide : Na2O + H2O = 2NaHO 48
Peroxide : Na2O2 + H2O = 2NaHO + O 49
of the substances contained jn 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 (CsHj)^ when treated with sodium, loses zinc and forms sodium
ethyl, C2H5Na, 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).
46 bis it ;s even doubtful whether the suboxide exists (see Note 47).
47 A compound, Na2Cl, which corresponds with the suboxide, in 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 NajCl; the
more so as the mass obtained gives- hydrogen when treated with water : Na^Cl + H2Q
= H+NaHO + NaCl, that is, it acts like suboxide of sodium. If Na«jCl realty exists as a
salt, then the corresponding base Na4O, according to the rule with other bases of the
composition M4O, 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
toot burn, the combustion only commencing with the formation of vapour — that is, when
SODIUM 541
All three oxides form sodium hydroxide with water, but only the
oxide Na2O is directly transformed into a hydrate. The other oxides
liberate either hydrogen or oxygen ; they also present a similar dis-
tinction with reference to many other agents. Thus carbonic anhydride
combines directly with the oxide Na2O, 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, Na2O, is the only salt-forming
considerably heated. Davy and Karsten obtained the oxides of potassium, K2O, and of
sodium, Na-jO, by heating the metals with their hydroxides, whence NaHO -f- Na=Na2O 4- 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 produeedj^NajO, "when "heated in an atmosphere of hydrogen,
gave a mixture of sodium and its hydroxide,: Na20 + 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'S. 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^O^ ; this
is obtained when sodium is burnt in an excess of oxygen. If NaN03 be melted, it gives
Na3O2 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^Oj + I^
=NazOl3+O. The compound NajOIj is akin to the compound Cu2OClij obtained by
oxidising CuCl. This reaction is one of the few in which iodine directly displaces
oxygen. The substance NajjOIj 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 : Na2C05=Na202-f-CO, whilst carbonic anhydride liberates oxygen
from it. With nitrous oxide it reacts thus: Na2O2 + 2N2O=2NaNO2+N2; with nitric
oxide it combines directly, forming sodium nitrite, NaO + NO = NaNO2. Sodium peroxide,
when treated with water, does not give hydrogen peroxide, because the latter in the
presence of the alkali formed (Na2Oj+2H2O=2NaHO + H202) decomposes into water
and oxygen. In the presence of dilute sulphuric acid it forms H2O2 (Na-jOj + H2SO4
= NajSOi + HaO2). Peroxide of sodium is now prepared on a large scale (by the action of air
upon Naat 800°) for bleaching wool, silk &c. (when it acts in virtue of the H202 formed).
The oxidising properties of Na2Og under the action of heat are seen, for instance, in the
fact that when heated with I it forma sodium iodate ; with PbO, Na-jPbOs > with pyrites,
sulphates, &o. When peroxide of sodium comes into contact with water, it evolves much
heat, forming H2O2, and decomposing with the disengagement of oxygen ; but, as a rule,
there is no explosion. But if Na2O2 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
^ajO2 on a large scale, Castner melts Na in an aluminium vessel, and at 800° passes
tost air deprived of a portion of its oxygen (having been already once used), and then
ordinary dry air over it.
542 PRINCIPLES OF CHEMISTRY
omde of £his metal, as water is in the case of hydrogen. Although the
peroxide H202 is derived from hydrogen, and Na2O2 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, CO2. Nitrogen and chlorine both give several
salt-forming oxides and types of salts. But of the oxides of nitrogen,
NO and NO2 do not form salts, as do N803, N204, and N2O5) although
N2O4 does not form special salts, and N2O5 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, Na2O,
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.
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 com-
pounds, 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.
543
CHAPTER XIH
POTASSIUM, RUBIDIUM, C/ESIUM, AND LITHIUM. SPECTRUM ANAEYSIS
JUST as the series of halogens, fluorine, bromine and iodine correspond
with the chlorine contained in common salt, so also there exists a cor-
responding series of elements : lithium, Li = 7, potassium, K = 39,
rubidium, Rb = 85, and caesium, 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, JBr = 80, and iodine,
1 = 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 K2SO4,Rb2S04
and €63804 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 K2S04
is 2-6638 of Rb2SO4, 8-6113, and of CsjSO^ 4'2484. 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 O'OOSl. The linear expansion (the maximum for the vertical axis) along the axis
of crystallisation is the same for all three salts, within the limits bl 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 caesium. 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 csesium salt, -and the differences are nearly in
the ratio 2 : 5. The lengths of the rhombic crystalline axes for £3804 are in the ratio
0-5727 : 1 : 0'7418, for Rb-jSC^ 0-6728 : 1 : 0-7485, and for Cs-jSO*, 0'67l2 : 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 crystallc*
graphic properties of the isomorphic rhombic sulphates RgSQt 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.
644 PRINCIPLES OF CHEMISTRY
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, SiO2,
potassium oxide, like sodium oxide, forms saline mineral substances
resembling glass. If other oxides, such as lime, CaO, and alumina,
A12O3, 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 com-
pounds as these which contain potash (potassium oxide), K2O, or soda
(sodium oxide), Na2O, and sometimes both together, silica, Si02, lime,
CaO, alumina, A1203, 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.,lbis are formed of such crystal-
line 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 (plagioclstse) 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 carnattite,
\ u« The origin of the primary rocks has been mentioned in Chapter X., Note 2,
POTASSIUM, RUBIDIUM, CAESIUM, AND LITHIUM
KMgCl3,6H20, occurs at Stassfurt. This carnallite2 Is now em-
ployed 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
2 Carnallite belongs to the number of doable 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 Stass-
furt 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.
5 The method of separating sodium chloride from potassium chloride has been de-
scribed 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°
82 85 40 46 . 57
When mixed with solutions of other salts the solubility of potassium chloride, naturally
varies, but not to any great extent.
j bin 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 + 63'29jJ + 0'226jp8, and therefore for 10 p.c. = l'0647, 20 p.c.
«. 1-1348, &c.
Potassium chloride combines with iodine trichloride to form a compound KC1 + IC15
= KIC14, 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 4 2C12 = KC1,IC13.
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
KC105 + I + 6HCl = KCl)ICl3 + 8Cl + 3H20. This is a kind of salt corresponding with
KIOj (unknown) in which the oxygen |s 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 com-
pounds with each other ; such as, KC1, C1I, KI, &c., whilst here they are grouped
together in the molecule KIC14. 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 Bb 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 : KI3, KBr2I, KClgl, and the corresponding
salts of rubidium and caesium, for instance, Cslj, CsBrlg, CsBr2I, CsClBrI CsCLI,
CsBrs, CsClBri,.CsCl4Br, and in general MX3 where X is a halogen. The colour of the
546 PRINCIPLES OF CHEMISTRY
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
The primary rocks contain an almost equal proportion of potassium
and sodium. But in sea water the compounds of the latter metal pre-
dominate. It may be asked, what became of the compounds of potas-
sium 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
crystals varies according to the halogen, thus CsI3 is black, CrBr5 yellowish red,CsBrI2
reddish brown, CsBr2I red, CsCl2Br yellow. The ctesium salts are the most stable, and
those of potassium least so, as also those which contain Br and I separately or together ;
for caesium no compounds containing Cl and I were obtained. The penta-haloid salts form
a smaller class ; among these salts potassium forms KC14I, rubidium KbCl^I, caesium
Csls, CsBr, CsCl4I, lithium LiClJ (with 4H20) and sodium NaCl4I (with 2H2O). The
most stable are those salts containing the metal with the greatest atomic weight — caesium
(see Chapter XI., Note 68).
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 in-
stance, 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 com-
pound 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 orthoclaae
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 combin-
ing weight of sodium (28) 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) other-
wise 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.
POTASSIUM, RUBIDIUM, CJESIUM, AND LITHIUM 547
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, C2HKO4, which is
employed for removing ink stains. Grape juice contains the so-called
cream of tartar, which is the acid tartrate of potassium, C4H5KO6.7
' 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.
? 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, K2O, CaO, MgO, and Fe2C>3, and of the four acid oxides, CO2, N2O5,
"PoO5, and S05, 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 com-
pounds 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 'reactions1
proceeding in organic substances in the presence of a small quantity of mineral substances
bas as yet been but little investigated, although there afe 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 maybe ex-
648 PBINjClPLES OF CHEMISTRY
This salt also separates as a sediment from wine. When the plants,
containing one or more of the salts of potassium, are burnt, thft
carbonaceous matter is oxidised, and in consequence the potassium is
obtained in the ash as carbonate, K2CO3, 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
pressed thus — two substances, A and B, do hot 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 combices with C, forming AC, and B reacts on this new
compound, which has a different store of chemical energy, forming-the compound AT* 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
Clements, 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 L, Note 19) not only
do not enter into the plants but kill them (poison them)
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 ia natural, since animals
consume plants. For example, milk, and especially human milk, contains a somewhat con-
siderable 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 con-
tains 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 VTIL,
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 com-
posed 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 gene-
rally taken and purified by crystallisation. Potassium carbonate crystallises with diffi-
culty, 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
POTASSIUM, RUBIDIUM, CESIUM, AND LITHIUM 549
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
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 BO 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
beating 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. KHCOj 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 K2CO38HgP in well-formed crystals at + 10°). There is no necessity to de-
scribe 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 'ley'— 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), K2CO3,6H2O,2Na2COs,6H2O. 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. H2O. This is evidently a combination by similarity, as in alloys,
eolations, &c.
8 bb About 25,000 tons of potash annually are now prepared from KC1 by this method
at Stassfurt.
9 Potassium sulphate, KjSOj, crystallises from its solutions in an anhydrous condi-
tion, in which respect it differs from the corresponding sodium salt, just as potassium car-
bonate 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, KHSO4, 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 tran»>
formations accomplished by means of sulphuric acid at a high temperature, because it ia
possible to take, in the form of this salt, a strictly definite quantity of sulphuric acid,
550 PBINCIPLES OF CHEMISTRY
bromide, and iodide 10 — by the action of the corresponding acid on KCI
and especially on the carbonate, whilst the hydroxide, caustic potasht
KHO, which is in many respects analogous to caustic soda, is easily ob-
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.
Bait 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 com-
pound containing eight equivalents of SO3 per equivalent of K3O. The salts of
rubidium, cajsium, 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, how-
ever, this method is rarely had recourse to, other more simple processes being em-
ployed 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 coloured 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 FeCO3) is precipitated (with an excess of iodine the pre-
cipitate is granular, and contains a compound of the suboxide and oxide of iron), while
potassium iodide remains in solution. Ferrous iodide, FeI2) 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).
POTASSIUM, fcUBIDIUM,. CESIUM, AND LITHIUM 551
tained 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.
Potassium cyanide, which presents in its chemical relations a certain
analogy with the halogen salts of potassium, is not only formed accord-
ing to the equation, KHO + HCN = H20 + KCN, but also when-
ever a nitrogenous carbon compound — for instance, animal matter — ia
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.
Tf the yellow prussiate be ground to a powder and dried, so that it loses
its water of crystallisation, it the"n melts at a red heat, and decomposes
into carbide of iron, nitrogen, and potassium cyanide, FeK4C6N6
= 4KCN -f FeC2 + N2. 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
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, KA1S208) and barium hydroxide,
BaHgOj. 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 9/04, but that of its solutions (see
Chapter XII., Note 18) at 15° S = 9,992 + 90'4p + 0'28p» (here p* is +, and for sodium
hydroxide it .is—). Strong solutions, when cooled, yield a*crystallo-hydrate, KHO,4HgO,
which dissolves in water, producing cold (like 2NaHO,7H20), whilst potassium hydroxide
in solution develops a considerable amount of heat.
u 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 end forms potassium cyanate, FeCOs + KQN = C0a + 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
552 PKINCIPLES OF CHEMISTKY
reaction, a swell-like that of bitter almonds, peculiar to prussic acid,
and acts as a most powerful poison. Although exceedingly stable
in a fused state, po'tassium 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 4- 2H2O = CHKO2 + NH3. 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 re-
ducing 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 + 2H20 + Zn =K2ZnC4N4 + 2KHO 4- H2 .
one part of charcoal powder to the mixture of 8 parts of anhydrous yellow prussiate and
8 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 some-
times 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 (Bossier and Gasslaker, 1892) of KCN and
NaCN, prepared by heating powdered and dried yellow prussiate with metallic sodium :
K4Fe (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 N2 ; 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):
POTASSIUM, RUBIDIUM, CAESIUM, AND LITHIUM 553
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 + HaO = ^AuKC^
+ 2KHO, which is taken advantage of for extracting gold from its
ores (Chapter XXIV.).13 bis Platinum, mercury, and tin. are not dis-
solved in a solution of potassium cyanide, even with access of air.
Potassium nitrate, or common nitre or saltpetre, KNO3, 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
15 *>>• A dilute eolation of KCN is taker*, not containing more than 1 per cent. ECN.
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.
14 Besides which Schloessing and Miintz, by employing e imilar methods to Pasteur,
showed that the formation of nitre in the decomposition of nitrogenous substances is
554 PRINCIPLES OF CHEMISTRY
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 col-'
lected, 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 con-
tact 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
brush-wood, 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 con-
taining various salts of nitric acid mixed with soluble orgaiiic 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.2CO3 + Ca(NO3)2 = 2KNO3 4- Ca003. 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
accomplished by the aid of peculiar micro-organisms (ferments), without which the
uimultaneous action of the other necessary conditions (alkalis, moisture, a temperature
of 37°, airland nitrogenous substances) cannot give nitre.
POTASSIUM. RUBIDIUM, CJESIUM, AND LITHIUM 555
repeated crystallisation. The greater part of the nitre used for
making gunpowder is now obtained from the sodium salt Chili salt-
petre 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 solu-
tions of potassium chloride and sodium nitrate is evaporated, sodium
chloride first separates, because this salt, which is formed by the
double decomposition KC1 + NaNO3 = KNO3 4- 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 solu-
tion of nitre, which cannot dissolve a further quantity of nitre but only
the impurities.
Nitre is a colourless salt having a peculiar cool taste. It crystal-
lises easily in long striated six-sided rhombic prisms terminating in
rhombic pyramids. Its crystals (sp. gr. T93) 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 bl3
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
M bi» 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- crystal-
lises 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 solidi-
fies -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 80 40 60 60 70 80 90
298° 283° 268° 242° 2S1U 281° 242° 284° 806°
which confirms Shaffgotsch's observation (1857) that the lowest fusing' point (about 231°)
is given by mixing molecular quantities (j> = 64'8) of the salts — that is, in the formation
of the alloy, KN05,NaNO3.
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 potas-
sium nitrate mixed with- 100— p parts by weight of sodium nitrate taken for solution,
556 PRINCIPLES OF CHEMISTRY
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 :
4KNO3 + 50 = 2K2C03 + 3CO2 + 2N2. This phenomenon depends on
the fact that oxygen in combining with carbon evolves more heat than
it do'es 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 • 2KN03-f 2S = K2SO4-f Nj
•f-SOo. 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, <fec. are in-
stances 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 temper-
and c be the quantity of the mixed salts which dissolves m 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 liqnid 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 potas-
sium 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 mini-
mum 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 nitro-
gen. The nitrite KNO2 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.
POTASSIUM, RUBIDIUM, CESIUM, AND LITHIUM 657
atures. Its application in gunpowder is based on this property ; gun-
powder consists of a mechanical mixture of finely-ground sulphur, nitre,
and charcoal. The relative proportion of these substances varies accord-
ing 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
maybe expressed by Ihe equation : 2KNO3 + 30 + S = K2S + 3CO2 + N2.
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 sul-
phur, 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 Gun-
powder explodes when heated to 800°, when struck, or by contact with a spark. A
compact or finely- divided mass of gunpowder barns 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, K2S, 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 87'4 parts of residue, K2S, 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 Bun-
sen, 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, 2KNOS + 8C + S = 2C + KSSO4 + COa + N2.
The residue in a blank cartridge often consists of a mixture of C, KZSO4, K3COS, and
KgSgOj. 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
558 PRINCIPLES OF CHEMISTRY
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, how-
ever, from that of sodium, because it easily combines with carbonic
oxide, forming an explosive and inflammable mass.17
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
will decrease. Under these circumstances the weight of the residue will be less — for
example, 4K2C03 + 4S = K2S04 + 3K2S + 4CO2. Besides which, carbonic oxide has been
found in the gases, and potassium bisulphide, K2S2, in the residue of gunpowder. The
amount of potassium sulphide, K2S, increases with the completeness of the combustion,
and is formed in the residue at the expense of the potassium sulphite. In recent tunes
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 con-
sequently gives (under proper- conditions of combustion) the possibility of communi-
cating 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
nitro-cellulose (Chapter VI., Note 87) 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 CO2, H2, &c.
17 The substances obtained in this case are mentioned in Chapter IX., Note 31.
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 nitro-
gen. But these, the first dalja 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 (11-6) KI ' 92 (84).
K 19 (19-5). BbCl 70 (60).
CsCl 89-5 (84-2J. Csl 188(180).
FeCl3 68 AgCl 80(71-7).
In brackets are given the densities corresponding with the formulae, according to
Avogadro-Gerhardt's law. This figure is not given for FeCl3, because in all probability
•under these conditions (the temperature at which it was determined) a portion of the
FeCl3 was decomposed. If it was not decomposed, then a density 81 would correspond
with the formula FeCls, and if the decomposition were Fe2Cl6 = 2FeCl2 + Cl2, 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 formulae that a better concord cannot be expected in such determinations.
V. Meyer (1887) gives S3 as the density of KL
POTASSIUM, KUBIDIUM, CESIUM, AND LITHIUM;. 059
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 coloured violet from
the presence of particles of potassium.21
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, K2H (between 200° and 411°), and with
oxygen it gives a suboxide K4O, oxide KaO, and peroxide, only more
oxygen enters into the composition of the latter than in sodium per-
oxide ; potassium peroxide contains K02, 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
19 The molecules of non-metals are more complex — for instance, Hj, Oj, Cls, &o. But
arsenic, whose superficial appearance recalls that of metals, but whose chemical proper-
ties approach more nearly to the non-metals, has a complex molecule containing ^84.
80 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 —
Na24 NaHOlS NaCl 28 NaNOj 87 NaaSOi 64
K 45 KHO 27 KC1 89 KNOS 48 K2SO4 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 carbouio
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 K2O + 0
= K2 + CO is reversible and the relation is the same in this case as between hydrogen
and zinc.
22 Potassium forms alloys with sodium in all proportions. The alloys containing 1 and
8 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.2K, NaK, NaK2, and NoK3 to be
44'5, 44-1, 48'8 and 44'4 thousand heat units respectively (for Na 42'6 and for K 45'4).
The formation of the alloy NaK2 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 NaK8 is now used for filling thermometers employed
for temperatures above 360°, when mercury boils.
*12
560 PRINCIPLES OF CHEMISTRY
sodium, 23, is as much greater than that of fluorine, 19, as that of
potassium, 39, is greater than that of chlorine, 35'5.
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, C4H5KO6
(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 KC1
and NaCl in solutions easily give double salts K2PtCl6 and Na2PtCl6,
with platinic chloride, PtCl4, 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 admix-
ture together, by means of their property of imparting different colours
to aflame. 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 undistin-
guishable, 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 coloured 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.
POTASSIUM, RUBIDIUM, CJESIUM, AND LITHIUM
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 dis-
persed in the prism, and the spectrum formed is observed through the
eye-piece of the other tele-
scope B. The third tube
D contains a horizontal
transparent scale (at the
outer end) which is
divided into equal divi-
sions-. 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 simul-
taneously with the spec- PIG. 72.— Spectroscope. The prism and table are covered with
A,,,, _ i , .!„ an opaque cover. The spectrum obtained from the flame
trum given by the rays coloured by a substance introduced on the wire is viewed
nacoi'nw thivninli fKo through B. A light Is placed before the scale D in order
passing tnrOUgn tne to uiuminate the image of the scale reflected through B
Slit of the tube C. In by the side of the prism.
this manner the image of the scale and the spectrum given by the
source of light under investigation are seen simultaneously. If the
a For accurate measurements and comparative researches more complicated spec-
troscopes are required which give a greater dispersion, and are furnished for this
purpose with several prisms — for example, in Browning's spectroscopa 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 reflec-
tion 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 oi the wave lengths it is particularly important that the spectroscope
should be furnished with diffraction gratings. The construction of spectroscopes des-
tined for special purposes (for example, for investigating the light of stars, or for deter-
mining 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 Boscoe,
Kayser, Vogel, and Lecoq de Boisbaudran.
562 PRINCIPLES OF CHEMISTRY
sun's rays be directed through the slit of the tube C, then the ob-
server looking through the eye-piece of B will see the solar spectrum,
and (if the aperture of the slit be narrow and the apparatus cor-
rectly 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 coloured 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 coloured
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 com-
pounds 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
84 The arrangement of all the parts of the apparatus so as to give the clearest possible
vision and accuracy of observation must evidently preoede every kind of spectroscopio
determination. Details concerning the practical use of the spectroscope must be
looked for in special works on the subject. In this trea ise 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 ths
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 thai, the Fraunhofer line A (the extreme prominent line in the red) cor-
responds 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 tha
127th division of the scale.
POTASSIUM, RUBIDIUM, CESIUM, AND LITHIUM 568
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 con-
tinuity 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, Angstrom, Kirchhoff, Corhu, 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 Fraun-
hofer 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 re-
fracting 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 ia 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 corre-
sponding with iron, are repeated, as Kirchhoff showed, in the solar
spectrum as dark Fraunhofer lines which occur in exactly the same situa-
tions 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
M 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 516-2,
498'8 and 498-2, &c., according to Liveing and Dewar.
•o bi" 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
$lit of the object tube.
564 PRINCIPLES OF CHEMISTRY
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 Ruhmkorffs coil is passed. Thus hydrogen gives a spectrum com-
posed of three lines — a red line corresponding with the Fraunhofer line
C, a green line corresponding with the line F, and a violet line corre-
sponding 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.
The correlation of the Fraunhofer lines with the spectra of metals
depends on the phenomenon of the so-called reversed of the spectrum. This
phenomenon consists in this, that instead of the bright spectrum cor-
responding with a metal, under certain circumstances a similar dark
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 < ABC D EbFGH
Wave-length 761'0 687'5 656'6 589'5-588'9 527'3 518'7 486'5 431-0 897'2
Colour red orange yellow green 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
POTASSIUM. RUBIDIUM, CAESIUM, AND LITHIUM
565
M I S I I I I I I $ | M 5 I 1 I I I
3 I I IgIS I I I I 15 1| I I I I I I I
e i i
i§§ i i i i i i i
Mill ISISI I I I I I IS I I I
I I I is-sss USE: I I I ISS I I
eg, _jp_>O_iQ "g^tf* ^ -^
,|
i i
spectrum in the form of Fraunhofer lines may be obtained, as will be
explained directly In order to clearly understand the phenomenon of
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-S 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
us 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). Conse-
quently, 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 dis-
posal for bright spectra of the in-
candescent 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
| I I Igl I I I Ijjgl I I I I I I
1 1 i2si
II
I I
i i i ss i es i i i i s i i i i i
i i
191
i i i i i i lisji i -i i i i i i i i
19 I8S I I i I IS I I I I I I I
11111 ills 1 1 1 1 1 1 1 1 iji
1 1
1 1
•* eo c* »-( a> 0
>o to ua to ^
I IIS I I I I
SgSSSJl I I I
666
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 per-
manganate, 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 solu-
tions of salts of didymium absorb rays of a certain refrangibility,
and therefore an impression of black lines is received,-1* as shown in
fig. 73. Many vapours (iodine) and gases (nitric peroxide) give similar
750710600 060 610690570 SfO *30 510 490 470
FIG. 73. — Absorption spectrum (Lecoq ile Boisbaudran) of salts of didymium in concentrated
and dilute solutions.
fepectra. 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
bbserved 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
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.
23 The method of observing absorption spectra consists iu taking a continuous
epectrunf of white light (one which doe's 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 ortho>
POTASSIUM, RUBIDIUM, CAESIUM, AND LITHIUM 667
linea 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 remark-
able progress made in all spectroscopic research dates from the in-
vestigations 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, Angstrom) that
the bright spectrum of the sodium flame gives two bright lines which
are in exactly the sa^ne 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 Same before it, a perfect super-
position 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 pro-
-duced — then there was no doubt that its appearance in the solar spectrvm
was due to the light passing somewhere through incandescent vapours
of sodium. Hence a new theory of reversed spectra 29 arose — that is,
chromatic 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.
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
FIG. 74.— Absorption spectra of nitrogen dioxide and iodine
for the genera] 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
Bnnsen's method sodium chloride is put into an apparatus for evolving hydrogen (the
•pray of the salt is then carried off by the hydrogen aud colours the flame with tbe
568 PRINCIPLES OF CHEMISTRY
of the relation between the waves of light emitted and absorbed by a
substance under given conditions of temperature ; this is expressed
by Kirchhoffs 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 coloured 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.
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
vaoours of the given metal. A similar phenomenon to that thus arti-
ficially 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 pre-
supposed that the sun itself, like all known sources of artificial light,
gives a continuous spectrum without Fraunhofer lines.31 We must
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
coloured 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.
50 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 coloured 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 ex-
planation and deduction of which must be looked for in text-books of physics, states that
the fraction AIE is a constant quantity depending on the nature of a substance (as A
depends on it) and determined by the temperature and wave-length.
31 Heated metals begin to emit light (only visible in the dark) at about 420° (vary-
ing 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.
POTASSIUM, RUBIDIUM, CJESIUM, AND LITHIUM 669
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 3a
contains very little, or no, metallic vapours, and as they cannot be sup-
posed 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 ^tself. 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 sepa-
rated 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 KirchhofFs
theory was formulated. By observations on the spectra of many
heavenly bodies, changes have been discovered going on in them,34 and
3* Brewster, as is mentioned above, first distinguished the atmospheric, cosmical
Fraunhofer lines from the solar lines. Janssen showed that the spectrum of the atmo-
sphere 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 : 680-622,
681-568, 535, 480-475. The same lines were found for liquid oxygen.
s* w» 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.
84 The great interest and vastness of astro-physical observations concerning the sun,
comets, stars, nebulae, &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
(he luminous note or wave-length of the light varies if the luminous (or absorbent) vapour
uid the earth front which we observe it approach or recede from each other; (his
570 PKINCIPLES OF CHEMISTRY
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 temper-
ature of the sun. This may indeed be concluded from the fact that
in ordinary speetroscopic 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
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 A, has an apparent wave-length A — — v, where n is the
n
velocity of light. Thus Tolon, Huyghens, and others proved that the star Aldebaran
approaches the solar system with a velocity of 80 kilometres per second, while Arcturus
is receding with a velocity of 45 kilometres. The majority of stars give a distinct
hydrogen spectrum, besides which nebulae 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 (pro-
tuberances) 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 897 and 393 appear at a higher temperature, and at a still higher temperature the
line 423 becomes quite invisible.
POTASSIUM. KUBLDIUM, CESIUM. AND LITHIUM .571
common salt be introduced into the flame of a gas-burner, a portion of
it is decomposed, first forming, in all probability, with water, hydro-
chloric acid and sodium hydroxide, and the latter then becoming partially
decomposed by the hydrocarbons, giving metallic sodium, whose incan-
descent 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 coloured 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
" r I • 3 I
Ea*> 3 F 6
: 1 IflHIHK '
i,iimi|iiM|iiii|iMi|lm|i|.i|HH|mi|
t\> 3 f 4 S 6
! Illl hf
1 ~El\> 3 T * i e
FIG. 75.— Bright spectra of copper compounds.
flalt be replaced by sodium, then either the bright line or the absorp-
tion 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, <fec., 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
672
PRINCIPLES OF CHEMISTRY
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
M 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.
Pliicker, Wiillner, Schuster, and others showed that at low temperatures and pres-
sures 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 N2 molecules are
obtained containing only one atom of nitrogen), or else it may be because at low tempera-
ture certain rays have a greater relative intensity than those which appear at higher tem-
peratures. 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,
AO f |B 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 absorp-
tion 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.
The spectra of solutions are best observed by means of Lecoq
de Boisbaudran's arrangement, shown in fig. 76. A bent capillary
tube, D P, 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.
FIG. 76.— Method of
showing the spec-
trum of substances
in solution.
POTASSIUM, RUBIDIUM, CESIUM, AND LITHIUM 578
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 ovon 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.
Spectrum analysis has not only endowed science with a knowledge
of the composition of distant heavenly bodies (of the sun, stars,
nebulae, comets, <fcc.), but has also given a new method for study-
ing 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
their light may be examined by placing the apparatus in front of the slit of a spectro-
scope. 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.
57 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 investi-
gators, 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 (Griinwald) 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*/(m'— 4),
if m varies as a series of whole numbers from 8 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 (896'8), and when
m is greater than 9, the ultra-violet lines of the hydrogen spectrum.
574 PRINCIPLES OF CHEMISTRY
salts and such very characteristic spectra that the least traces of
them 3* are discovered with great ease by means of the spectroscope.
For instance, lithium gives a very brilliant red coloration to a Same
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.
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,
epodurnene and lithia mica. Many compounds of lithium are in all
respects closely analogous to the corresponding compounds of sodium
88 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 8 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.
ss 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
nric 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 chlor-
ides of the alkaline metals KC1, 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^Oj = 2NaCl + LioCOj. 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 carbo-
nates are only sparingly soluble. Oxide of lithium, Li2O, 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 Li2 with O evolves 140,000 calories — that is,
more than Na^O (100,000 calories) and K2O (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 Li3N, like Na5N (see Chapter XII. Note 50).
LiCl, LiBr, and Lil form crystallo-hydrates with H2O, 2HjO, and 3H2O. As a rule,
LiBr2H2O crystallises out, but Bogorodsky (1894) showed that a solution containing
LiBr + 3'7H2O, cooled to — 62°, separates out crystals LiBr3H2O, which decompose at + 4°
with the separation of H2O. LiF is but slightly soluble (in 800 parts) in water (and still
less so in a solution of NH4F).
POTASSIUM, RUBIDIUM, CESIUM, AND LITHIUM 575
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, Li2CO3, 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,H20. 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 (sur-
rounded 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 consider-
able quantity of the metal has accumulated
,, . ... -iij .1 i • FlO. 77. — Preparation of lithium by the
On the Wire it IS Withdrawn, the metal IS action of a galvanic current on fused
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
Bunsen in 1860 tried to determine by means of the spectroscope
S9 bb Gu'ntz (1898) recommends adding KC1 to the LiCl in preparing Ll by this
method, and to act with a current of 10 amperes at 20 -volts, and not to heat above 460°,
so as to avoid the formation of Li2Cl.
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 hydro-
fluoric acid must be taken), and to treat the residue with sulphuric acid, evaporate to
drynesB, 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.
676 PRINCIPLES OF CHEMISTRY
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 ccesium, because it
colours a pale flame sky blue, which depends on its containing bright
blue rays, which appear in the spectrum of caesium 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 caesium. The amount of the oxides of caesium 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 KC1) forms
an abundant source for rubidium and also partly for caesium.
Rubidium also occurs, although in very small quantities, in the majority
of mineral waters. In a very few cases caesium is not accompanied by
rubidium ; thus, in a certain granite on the Isle of Elba, caesium has
been discovered, but not rubidium. This granite contains a very rare
mineral called pollux, which contains as much as, C4 per cent, of caesium
oxide. Guided by the spectroscope, and aided by the fact that the
double salts of platinic chloride and rubidium and csesium chlorides
are still less soluble in water than the corresponding potassium salt,
K2PtCl6,41 Bunsen succeeded in separating both metals from each
other and from potassium, and demonstrated the great resemblance
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 be 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
caesium 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 0'1S4 part, and the ceesium salt,
0'024 part; at 100° 6'1S parts of potassium platinochloride, KoPtClg, are dissolved,
0'634 part of rubidium platinochloride, and 0'177 part of caesium platinochloride.
From this it is clear how the salts of rubidium and caesium may be isolated. The
separation of caesium 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; caesium carbonate, CsjCOs, 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, caesium, and rubidium in the pre-
sence of hydrochloric acid, stannic chloride precipitates a double salt of caesium, which
is very slightly soluble. The salts of Rb and Cs ore closely analogous to those of
poUssium.
POTASSIUM. RUBIDIUM, CESIUM, AND LITHIUM 577
they bear to each other. The isolated metals,42 rubidium and caesium,
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
caesium by heating CsAlO2 with Mg (42 bls).
Judging by the properties of the free metals, and of their corre-
sponding and even very complex compounds, lithium, sodium, potas-
sium, rubidium, and csesium present an indubitable chemical resem-
blance. The fact that the metals easily decompose water, and that their
a Bunsen obtained rubidium by distilling a mixture of the tartrate with soot, and
Beketofi (1888) by heating the hydroxide with aluminium, 2EbHO + Al = Rb A1O.4 + H2 + Rb.
.By the action of 85 grams of rubidium on water, 94,000 heat units are evolved. Setterberg
obtained caesium (1882) by the electrolysis of a fused mixture of cyanide of caesium 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).
a 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
alnminates (Chapter XVII.), RA1O2— for example, 2KHO + A1 = KA1O2 + K + H2. 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 caesium upon this principle by heating its aluminate CsAlO., with
magnesium powder. In this case aluminate of magnesium is formed, and the whole of
the caesium is obtained as metal: 2CsAlO.j + Mg = MgOAl2O5 + 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? was
prepared by the precipitation of caesium -alums by caustic baryta, and evaporating the
resultant solution. We may add that N. N. Beketoff (1887) pfepared oxide of potassium,
K2O, 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 evolve?
24,920 cal. ; so that K2O + H2O gives 42,480 cal.), whence (knowing that K2 + O + H2O in
an excess of water evolves 164,500) it follows that K2 + 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 K2O evolves more heat
than Na<jO in combining with water (sen Chapter II. Note 9). Just as hydrogen displaces
half the Na from Na-jO forming NaHO, so also N. N. Beketoff found from experiment
and thermo-chemical reasonings that hydrogen-displaces half the potassium from K2O,
forming KHO and evolving 7,190 calories. Oxide of lithium, Ia2O, which is easily
formed by igniting Li2COs with carbon (when Li2O + 2CO is formed), disengages
26,000 cals. with an excess of water, while the reaction Li2 + O gives 114,000 cals. and
the reaction Llj + HgO gives only 18,000 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 RgO gives most heat with Li and least with K, while the formation of
RC1 evolves most heat with K (105,000 cats.) and least of all with Li (93,500 cals.).
Rubidium, in forming Rb2O, gives 94,000 cals. (Beketoff). Caesium, in acting upon an
excess of water, evolves 51,500 cals., and the reaction CB.J + O evolves about 100,000 cals.
i.e. more than K and Rb, and almost as much as Na — and oxide of caesium reacts
with hydrogen (according to the equation C^O + H = CsHO + Cs) more easily than any
of the oxides of the alkali metals, and this reaction takes place at the ordinary tem-
perature (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.
578 PRINCIPLES OF CHEMISTRY
hydroxides HHO and carbonates R2CO3 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 com-
pounds, 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 IC1 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.
Tt is thus correct to say that both the above-named orders of ele-
POTASSIUM, RUBIDIUM, CJESIUM,' AND LITHIUM 679
ments 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 re-
place 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 com-
pounds 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, K2O, HC10, and C12O, with water. It
must not be forgotten, however, that the halogens give, with oxygen,
besides compounds of -the type R2O> 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
The atomic weights of the alkali metals, lithium 7, sodium 23,
potassium 39, rubidium 85, and caesium 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
4S 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).
Schiitzenberger obtained a compound C2HSO(OC1), which he called chlorine acetate, by
acting on acetic anhydride, (C^HjO^O, with chlorine monoxide, C12O. With iodine this
compound gives off chlorine and forms iodine acetate, C2H3O(OI), which also is formed
by the action of iodine chloride on sodium acetate, CiHjO(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 C12O
itself, or IC1 or KNa. By the action of chlorine monoxide on a mixture of iodine and
acetic anhydride, Schiitzenberger also obtained the compound I(C2H304)3, which is
analogous to IC1S, because the group C2H5O2 is, like Cl, a halogen, forming salts with
the metals. Similar properties are found in iodosobenzene (Chapter XI., Note 79).
580 PRINCIPLES 'OF CHEMISTRY
of potassium, rubidium, and caesium 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 them-
selves ; thus lithium volatilises with difficulty, whilst sodium is obtained
by distillation, potassium volatilises more easily than sodium, and
rubidium and caesium as we have seen, are still more volatile.
581
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 com.
pounds, 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 35g5 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
1 Under favourable circumstances (by taking all the requisite precautions), the weight
of the equivalent may be accurately determined by this method. Thus Keynolds and
Ramsay (1887) determined the equivalent of zinc to be 82-7 by this method (from the
average of 29 experiments), whilst by other methods it has been fixed (by different
observers) between 82-55 and 88'95.
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).
* The most accurate determinations of this kind were carried on by btas, and will be
described in Chapter XXIV.
s bis The amount of electricity in one coulomb according to the present nomen-
clature of electrical units (see Works on Physics and Electro-technology) disengages
0'00001036 gram of hydrogen, 0-00112 gram of silver, 0-0008268 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
582 PRINCIPLES OF CHEMISTRY
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.
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 ?i-valent
metals it is equal to the weight of n equivalents. Thus aluminium,
41 = 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 M2O, MHO,
MCI, MNO3, M2SO4, &c., and in general MX, then for bivalent
metals like magnesium or calcium the corresponding compounds
will be MgO, Mg(HO)2, MgCl2, Mg(NO3)2, MgSO4, &c., or in general
MX2. *
By what are we to be guided in ascribing to some metals uni-
valency and to others bi-, ter-, quadri-, -u-valency 1 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 ?
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 furthef develop-
ment of our science.
There is only one side of the subject respecting the direct correlation between thermo-
chemical data and electro-motive force, which I thirik 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 experi-
ment and theory, because in it there takes place the decomposition of CuSO4 into Cu + O
together with the formation of Zn + O and ZnO + SO3Aq, 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^S04 + PbO-j, as Cheltzoff showed) E = 0'0436Q.
THE VALENCY AND SPECIFIC HEAT OF THE METALS 583
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 com-
position 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 com-
pounds correspond, like the compounds of nitrogen, with the forms
AsX3, and AsX5 ; for example, AsH3, AsCl3, AsFl5, As2O5, &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 caesium) ; so,
although very few volatile compounds of bismuth are known, it was
necessary to ascribe to them formulae 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, ZnC4H10, which boils at
118°, vapour density = 61 '3), and in the molecules of all these com-
pounds there is never less than 65 parts of zinc, which is equivalent to
H2, 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 H2 is replaced by O = 16), is a repre-
sentative 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 t£>
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 Cu2O and the-
*13
684 PRINCIPLES OF CHEMISTRY
oxide CuO— that is, the compounds CuX corresponding with the sub-
oxide are analogous (in the quantitative relations, by their composition)
to NaX or AgX, and the compounds of the oxide CuX2, to MgX2,
ZnX2, and in general to the bivalent metals. It is clear that in such
examples we must make a distinction between atomic Veights 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.
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
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 Avogadrc-
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, Marignoc, and others took advan-
tage of this method for determining the composition'of the compounds of many elements.
(8) The specific heat, according to Dnlong and Petit's law. Begnault, 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 atoni
(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 be taken as trivalent — that is, if the composition of
its chloride 'be taken as BeCls (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 compoSi-
Sion of its chloride as BeCl2, 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 BeCl2 + 200H2O = 1-0138— that is, greater than the corresponding solution
KC1 + 200H2O ( = 1-0121), and less than the solution MgCl2+200HjO ( = 1-0208), as would
follow from the magnitude of the molecular weight Bed, = 80, ginoe KC1=74'6 and
MgCl2 = 96.
THE VALENCY AND SPECIFIC HEAT OF THE METALS 585
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 conclu-
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 deter-
mination. 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 800° = 0-1267, and at 1,400°
= 0*4031. Between these last limits of temperature a change takes place in iron (a spon-
taneous heating, recalescence), as we shall see in Chapter XXII. For quartz SiO3
Pionchon gives Q=OT787 + 894<10-G-27<210-8 up to 400°, for metallic aluminium
(Eichards, 1892) at 0' 0'222, at 20° 0'224, at 100° 0'282 ; 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 0-28 0'29 0'44 0-46
Diamond O'lO 019 0'22 0'44 0'45
Boron ' 0'22 0'29 0'85
These determinations, which have been verified by Dewar, Le Chatelier (Chapter VIII.,
Note 18), Moissan, and Ganthier, 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 Ax Q at 0°=l-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 multipled by 12
686
PRINCIPLES OF CHEMISTRY
sions arrived at by Dulbng and Petit with respect to solid elementary
bodies.
Na
23
0-2934
Q =
AQ =
In
7
0-9408
6-59
Fe
56
0-112
6-27
A=±
Q =
AQ =
Pd
A= 106
Q = 0-0592
AQ = 6-28
Pt
A= 196
Q = 0-0325
AQ = 6-37
6-75
Cu
63
0-093
5-86
Ag
108
0-056
6-05
Au
198
0-0324
6-41
Mg
24
31
0-245
0-202
5-88
6-26
Zn
65
Br
80
0-093
0-0843
6-04
6-74
Sn
118
I
127
0-Q55
0-541,
6-49
6-87
Hg
200
Pb
206
0-0333
0-0315
6-66
6-49
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
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 = ll C = 12
Q= 0-94 0-42 0-24 0'20
AQ= 6'6 8-8 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 VTII. In the case of sulphur the
molecule contains at least Sa and its atomic heat = 32 x 6'163 = 5'22, which is distinctly
below the normal. If a large number of atoms of carbon are contained in the molecule
cf 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 approxi-
mate 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.
THE VALENCY AND SPECIFIC HEAT OF THE METALS 587
multiplying the specific heats found by experiment by the correspond-
ing 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 (Mendele'etf). If the same
composition be ascribed to the compounds of magnesium as to the
corresponding compounds of potassium, then the equivalent of mag-
nesium 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 x 0-245 — 5-9 j for calcium,
giving its compounds a composition CaX2 — for example CaCl2, CaSO4,
CaO (Ca = 40) — we obtain an atomic heat = 40 x 0-17 = 6-8, and for
barium it is equal to 137 x 0-05 = 6-8 ; that is, they must be counted
as bivalent, or that their atom replaces H2, Na2, or K2. This con-
clusion may be confirmed by a method of analogy, as we shall afterwards
see. The application of the principle of specific heats to the determi-
nation 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 pro-
fessor Cannizzaro.
Exactly the same conclusions respecting the bi valence 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,
MgCl2 and CaCl2, are 0'194 and 0-164, and of sodium and potassium
chlorides, NaCl and KC1, 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
5 It must be remarked that in the case of oxygen (and also hydrogen and carbon)
compounds the quotient of MQ/w, 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-l, MnO2 = 4'6, ice
(Q=0'604) = 8, SiOa = 8-5, &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
588 PRINCIPLES OF CHEMISTRY
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
extent be seen in the specific heat of the oxides. Thus for alumina, Al2Oj (Q=0'217),
MQ = 22'3, and therefore the quotient MQ/w=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
8'7, 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 = H-8 ;
MgSO4 = 27'0 and ZnSO4 = 28'0, &c.
6 If W be the amount of heat contained in a mass m of a substance at a temperature
',, and <fW the amount expended in heating it from t to t + dt, then the specific heat
Q = dWf(m x 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 Wiedernann 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 in-
considerable. According to Mallard and Le Chatelier it is = — — — per 1°, where M is
the molecular weight (for instance, for O2, M = 82). Therefore the specific heat of those
permanent gases which contain two atoms in the molecule (H2, O2, N2) 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 —278°); and therefore
MQ = 6'8 + a(273 + £), 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)=8'4, as it is in fact (Chapter IX., Note 17 bls). As regards
liquids (as we as the vapours formed by them), the specific heat always rises with the
temperature. Thus for benzene it equals 0'38 + 0'0014<. 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 EL, 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=Bt.
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 0-28 ; steam 0'475, liquid water TO. But the complexity of the relations exist-
ing 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 80°), 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 0-81, 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
THE VALENCY AND SPECIFIC HEAT OF THE METALS 589
complex quantity — including not only the increase of the energy of a
substance with its rise in temperature, but also the external work of
expansion7 and the internal work accomplished in the molecules
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 tem-
perature 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 ex-
pended in external work. In the case of gases the last quantity may be easily deter-
mined, 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 10338 x 0-00368,
or S8'02 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 l.'0"0448M cubic metres, and hence the external
work D in the heating of 1 kilo of the given gas through 10 = 0'0896/0-0448M, or D = 2/M.
Taking the magnitude of the internal work B for gases as negligeable if permanent gases
lire 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 Kegnault's experiments, for
oxygen it = 0'2175, for hydrogen 8'405, 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= 8-4050-1-000 = 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-
thai is, their product by the atomic weight is a constant quantity. In fact, for oxygen
this product =0-155x16 = 2-48, for hydrogen 2'40, for nitrogen 0-7724x14 = 2-414, and
therefore if A stand for the atomic weight we obtain the expression K x 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 2-45n + 2 to 2'45n. 'When w = l this ratio k = 1-8; when
'n = 2, A; = 1*4, when n = 3, k = 1-3, and with an exceedingly large number n, of atoms in the
"molecule, k = l. That is, the ratio between the specific heats decreases from 1-8 to TO
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, <fec.) and is 'found in reality to
be nearly 1-4, and for such gases as carbonic anhydride, nitric dioxide, and others it is
nearly rs. Kundt and Warburg (1875), by means of the approximate method mentioned
in Note 29, Chapter VII., determined k for mercury vapour when n=l, 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 = 2'48, only under the condition
590 PRINCIPLES OF CHEMISTRY
causing them to decompose according to the rise of temperature* —
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.
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 89 — that is, the latter
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, Kegnault ; 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 assxvmed, 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/w, where^ 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 tlie
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,
N2O4, which, when heated, gradually passes into NOg — thai is, chemical work of decom-
position pi'oceeds, 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$
to tL it cannot but suffer a chemical change (that is, the state of the atoms in the mole-
cules changes more or less in one way or another) if dissociation sets in at a temper-
ature ti. 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.
THE VALENCY AND SPECIFIC HEAT OF THE METALS 591
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 rescm
blance 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 + 2H2O = CaH2O2 + H2 — that is, hydrogen is liberated
and a hydroxide of the metal formed. These hydroxides are bases
which neutralise nearly all acids. However, the hydroxides RH2O2 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.
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, HO, 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
9 As if NaH = Mg and KH = Ca, which is in accordance with their valency. KH
includes two monov&lent elements, and is a bivalent group like Ca.
592 PRINCIPLES OF CHEMISTRY
and fall as a precipitate, whilst, the metals of the alkalis remain in
solution : RX2 + Na2CO3 = RCO3 + 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 CaCO8 and
MgC63, frequently met with in nature, are soluble in an excess of
water saturated with carbonic anhydride,™ 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 in-
soluble precipitate of calcium carbonate separates out. It may confi-
dently 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 pre-
cipitation of the carbonates from the sea water, because the plants
absorb C02, and many of the organisms CaCO3, 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 pre-
dominates, because it is present in rooks and running water in greater
quantity than magnesia, and in this case these sedimentary rocks are
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, NaHCO3.
THE VALENCY AND SPECIFIC HEAT OF THE METALS 593
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 car-
bonic anhydride so easily as the limestones under the action of acids.
Dolomites ' ' 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 car-
bonate itself is exceedingly common in tb\s form in nature, and is -then
known as cole 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 con-
siderable 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, CaSO4,2H2O. 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.
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 liine 12 itself and with
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 XEL), when we shall obtain a mixture of the carbonates of calcium
and magnesium. Haitinger heated a mixture of calcium carbonate, CaCOs, with a solu-
tion of an equivalent quantity of magnesium sulphate, MgSO^ in a closed tube at 200°.
and then a portion of the magnesia actually passed into the state of magnesium car-
bonate, MgCO3, and a portion of the lime was converted into gypsum, CaS04. Lubavin
(1892) showed that MgCO5 ia more soluble than CaCOsin salt water, which is of somft
significance in explaining the composition of sea water.
11 The undoubted action of lime in increasing the fertility of soils — if not ia every
594 PRINCIPLES OF CHEMISTRY
marl — that is, with clay mixed with a certain quantity of calcium car-
bonate, strata of which are found nearly everywhere
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 frotn 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
case, at all events, with ordinary soils which have long been under corn — is based not
BO 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.
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 deve-
lops more heat (about 140 thousand calories) than Kj+ O or Na.2 + 0 (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 + Cl2 gives 150 and Naxj + Clj
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).
THE VALENCY AND SPECIFIC HEAT OF THE METALS 595
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 duoride,
is fused in a close crucible. The latter substances only serve to facili-
tate 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, MgCl2 -I- Naa
= Mg + 2NaCl. In working on a large scale, the powdery metallic
magnesium is then subjected to distillation at a white heat. The dis-
tillation of the magnesium is necessary, because the undistilled metal is
not homogeneous u 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
ut about 1000° It is malleable and ductile, like the generality of
metals, so tha_t 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, MgH2O2, 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
14 Commercial magnesium generally contains a certain amount of magnesium nitride
(Deville and Caron), Mg5N2 — 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 MgsN2
is easily formed and is the sole product when Mg is heated to redness in a current of NH5.
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.
596 PRINCIPLES OF- CHEMISTRY
evolves about 140 thousand heat units, and that the product of com-
bustion, 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 pro-
duction 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
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
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 tempera-
ture, 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
u; 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 Ruhrakorff's coil (D Mendeleeff, 1889).
1? According to the observations of Maack, Comaille, Bcittger, and others. The re-
duction 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.
19 Davy, on heating magnesia in chlorine, concluded that there was a complete sub-
stitution, 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 Dot complete and is limited by a reverse reaction.
THE VALENCY AND SPECIFIC HEAT OF THE METALS 597
much more easily than that of calcium salts. For example, mag-
nesium 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, MgH202). 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 grea£ 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, MgCl2 + 2KHO = Mg(HQ)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 per-
ceptible quantity of magnesium hydroxide — namely, one part is dis-
solved by 55,000 parts of water. Such a solution, however, has
an alkaline reaction, and gives, with a salt of phosphoric acid, *
precipitate of magnesium phosphate, which is still more insoluble.
Magnesia is not only dissolved by acids, forming salts, but it also dis-
places certain other bases — for example, ammonia from ammonium
'salts when boiled; and the hydroxide also absorbs carbonic anhy-
dride from the air. The magnesium salts, like those of calcium, potas-
sium, 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.
The power of magnesium salts to form double and basic salts is
very frequently shown in reactions, and is specially marked as re-
gards ammonium salts. If saturated solutions of magnesium and
ammonium sulphates are mixed together, a crystalline double salt
Mg(NH4)2(SO4)2,6H2O,20 is immediately precipitated. A strong
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° 17'8 parts of the anhydrous
salt. At about 180° it loses all its water.
598 PRINCIPLES OF CHEMISTRY
solution of ordinary ammonium carbonate dissolves magnesium
oxide or carbonate, and precipitates crystals of a double salt,
Mg(NH4)2(CO3)2,4H2O, from which water extracts the ammonium
carbonate. With an excess of an ammonium salt the double salt passes
into solution/1 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(NH4)Cl3,6H2O.2a
The salts of potassium, like those of ammonium, are able to enter into
combination with the magnesium salts.23 For instance, the double
salt, MgKCl3,6H2O, which is known as carnattite,1** 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 satu-
rated solution of magnesium sulphate dissolves potassium sulphate, and
solid magnesium sulphate is soluble in a saturated solution of potassium
sulphate. A double salt,- K2Mg(SO4)2,6H2O, which closely resembles
the above-mentioned ammonium salt, crystallises from these solutions.25
J1 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, 2MgCl2 •*• 2NH4.OH — Mg(OH)j
+ Mg.NH4Cls + NH4C1. A solution of ammonium chloride reacts with magnesia, evolving
ammonia and forming a solution of the same salt, MgO + 8NH4Cl = MgNH4Cls+H2O
+ 2NH3.
Among.the double' salts of ammonium and magnesium, the phosphate, MgNH4PO4,6H2O,
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.
25 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 pre-
dominate 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(S04),8H2O (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 Riidorff, the double salts, like camallite, MgK2(S04)2,6HQO, and the alums,
all belong to this order (1888). But such salts as tartar emetic, the double oxalates, and
THE VALENCY AND SPECIFIC HEAT OF THE METALS 599
The nearest analogues of magnesium are able to give exactly similar
double salts, both in crystalline form (moaoclinic system) and coru-
donble 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
K2Mg(S04)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, NaxjMg(SO4).,, which crystallises
with 4 and 6 mol. OH2. The first crystallo-hydrate, MgNa2(SO4)2,4H2O, 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 oi water
(according to the equation MgNas(S04)2,4H2O + 13H,jO = Na2SO4,10H2O + MgSO4,7H2O),
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 Koozeboom, 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 astrakhauite 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/jSO^lOK^O in grams =322/l'46=220'5 cubic centimetres, and by
.MgSO4,7H20 = 246/l'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
ISHjO = 234 c.c., hence their sum = 880'5 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 aetrakhanite, 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. Koozeboom showed that there are two
Emits 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 2MgSO4,7HoO + 2NaCl
passes into MgNa^(SO4)o,4H2O + MgCl2,CH2O + 4H.A whilst below 31° there is not this
600 PRINCIPLES OF CHEMISTRY
position ; they, like this salt (see Chapter X\7'.), 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, MgSO4.26 It occurs at Stassfurt as kieserite, MgSO4,H20,
and generally separates from solutions as a heptahydrated salt,
MgS04,7H2O, and from supersaturated solutions as a hexahydrated
salt, MgSO4,6H2O ; at temperatures below 0° it crystallises out as a
dodecahyd rated salt, MgSO4,12H2O, and a solution of the com-
position MgSO,,2H2O solidifies completely at — 5°.27 Thus between
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,
K2SO4MgSO4,4H.2O, 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 r6le 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.jSO^lOILjO
and MgSO4,7H2O. Water is disengaged in both cases, and hence the volumes are altered.
86 This salt, and especially its crystallo-hydrate with 7H2O, 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 hydro-
chloric acid until about half 'of the salt remains, the calcium carbonate is mostly dis-
solved and magnesium carbonate is left, which by treatment with sulphuric acid gives a
solution of magnesium sulphate.
27 The anhydrous salt, MgSO4 (sp. gr. 2'61), attracts moisture (7 mol. H2O) 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 2MgSO4 + C = 2SO2 + COj
+ 2MgO. The monohydrated salt (kieserite), MgSO4,H2O (sp. gr. 2'56), dissolves in
water with difficulty ; it is formed by heating the other crystallo-hydrates to 185°.
The hexahydrated salt is dimorphous. If a solution, saturated at the boiling-point,
be prepared, • arid cooled -without access of crystals of the heptahydrated salt, then
MgSO4,6H2O crystallises out in monoclinic prisms (Loewel, Marignac), which are quite
as unstable as the salt, Na.jSO4,7H2O ;-but if prismatic crystals of the cubic system
of the copper-nickel salts of the composition MSO4,6H20 be added, then crystals of
MgSO4,6H2O are deposited on them as prisms of the cubic system (Lecoq de Boisbau-
dran). The common crystallo-hydrate, MgSO^H-jO, Epsom salts, belongs to the rhombiQ
system, and is obtained by crystallisation below 80°. Its specific gravity is 1'69. In a
vacuum, or at 100°, it loses 5H2O, at 132° GH2O, and at 210° all the 7H2O (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 trans-
formation into the more stable common form. Fritzsche, by cooling saturated solutions
below 0°, obtained ft mixture of crystals of ice and of a dodecahydratad salt, which easily
THE VALENCY AND SPECIFIC HEAT OF THE METALS 601
'water and magnesium sulphate there may exist several definite and
more or less stable degrees of equilibrium ; the double salt
MgSO4K2SO4)6H2O may be regarded as one of these equilibrated
systems, the more so since it contains 6H.20, whilst MgSO4 forms its
most stable system with 7H2O, and the double salt may be considered
as this crystallo-hydrate in which one molecule of water is replaced
by the molecule K2SO4.28
The power of forming basic salts is a very remarkable peculiarity
of magnesia and other feeble bases, and especially of those corre-
sponding with polyvalent metals. The very powerful bases corre-
sponding 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,
Bplit up at temperatures above 0°. Guthrie showed that dilute solutions of magnesium
sulphate, when refrigerated, separate ice until the solution attains a composition
MgSO4,24H2O, which will completely freeze into a crystallo-hydrate at — 5'3°. According
to Coppet and Riidorff, 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) t = 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 MgSO4 in the presence of the hexa-
hydrated salt, 84'67 MgSO4 in the presence of the hexagonal heptahydrated salt, and
only 26 parts of MgSO4 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 phenol-
phthalein.
The specific gravity of solutions of certain salts of magnesium and calcium reduced
to 16°/4° (see my work cited, Chapter I., Note 19), are, if water at 4° =10,000,
MgS04 :
MgCl2 : s = 9,992
CaCl2 : s = 9,992
K 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,(SO4)2,6H2O lose all their water at 135°, whilst MgSO4,7H2O only parts
with CH,O.
602 PRINCIPLES OF CHEMISTRY
Mg(HO)2,4MgC03,9H2O ; 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, 4MgSO4
-f- 4Na,CO3 + 4H20 = 4Na2S04 + CO2 + Mg(OH)2,3MgCO3,3H2O.
This basic salt forms the ordinary drug magnesia (magnesia atba), 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, MgCO3, •tfhich 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 pro-
ducts 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, MgCl2. 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, MgCl2,6H20, separates ; 32 but if it be
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 evolu-
tion of hydrochloric acid is observed just as in the action of sulphuric acid on common
salt, MgSO4 + 2NaCl + H2O=Na2SO4-l-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, MCI and MgX2, it follows, according to Ber-
thollet's teaching, that MgCl2 is also present.
32 As the crystallo-hydrates of the salts of sodium often, contain 10H4O, so many of
the salts of magnesium contain 6H2O.
THE VALENCY AND SPECIFIC HEAT OF THE METALS 603
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 evolu-
tion 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, MgCl2,2NH4Cl)
and then ignited (460°), when the sal-ammoniac is converted into
vapour and a fused mass o£ 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 heatiny
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 pre-
paration 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
53 This decomposition is most simply defined as the result of the two reverse reactions,
MgCl2 + H2O = MgO + 2HC1 and MgO + 2HC1 = MgCl2 + H2O, or as a distribution
between O and C12 on the one hand and H2 and Mg on the other. (With O, MgCl2 gives
chlorine, see Chapter X., Note 33, and Chapter II., Note 3bis and others, where the
reactions and applications of MgCl2 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, MgCl2,6H2O, forms the limit of the reversibility. But an inter-
mediate 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
ft solid mass, called cylolite, used for flooring, &c.
We may remark that MgBrj crystallises not only with 6H2O (temperature of fusion
162°), but also with 10H2O (temperature of fusion + 12b, formed at - 18°. Panfiloff, 1894).
w According to Thomson, the combination of MgCl2 with 6H2O evolves 33,000 calories,
and its solution in an excess of water 36,000.
604 PRINCIPLES OF CHEMISTRY
is formed with even greater facility under the action of heat in the
presence of steam, when HC1 is formed, according to the equation
2MgCl2 + H20 = MgOMgCl, + 2HC1.3' bi*
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 decom-
poses calcium chloride36 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 caleiom 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
34 bls Hence MgClj may be employed for the preparation of chlorine and hydrochloric
acid (Chapters X. and XI.). In general magnesium chloride, which i& obtained in large
quantities from sea water and Stassfurt carnallite, may find numerous practical uses.
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 48° 0'85 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 (Lie's-Bodart). Dumas advises carrying on this
reaction in a closed space under pressure.
THE VALENCY AND SPECIFIC HEAT OF THE METALS 605
metal is volatile. Calcium decomposes water at the ordinary tem-
perature, 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 : CaCO3=CaO + C02.
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
Calcium oxide — that is, quicklime — is a substance (sp. gr. 3- 15)
3S 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 tacool somewhat, the lime
raked out, and the same process repeated. In the continuously acting furnace's, con-
structed 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 charac-
terised 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, CaN2Og, 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 : CaN2O0
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.
606
PRINCIPLES OF CHEMISTRY
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, CaH2O2 (specific
gravity 2'07) is a most common alkaline substance, employed largely
fio. 78.— Continually-acting kilo for burning lime. Tlie lime is charged from above and calcined by
four lateral grates, R, M. D, fire-bars. B, spnce for withdrawing the burnt lime. K, stoke-
house. M, fire grate. Q, R, tinder-grate.
in building for making mortars or cements, in which case its bind-
ing property is mainly due to the absorption of carbonic anhydride.40
59 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
THE VALENCY AND SPECIFIC HEAT OF THE METALS 607
Lime, like other alkalis, acts on many animal and vegetable sub-
stances, and for this reason has many practical uses— for example,
for removing fats, and in agriculture for accelerating the decom-
position of organic substances in the so-called composts or accumu-
lations 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 dis-
tinguishable 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 4l If
lime-water be mixed with hydrogen peroxide minute crystals of
calcium peroxide, CaO2,8H20, separate ; this compound is very un-
stable 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
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 experi-
ment) 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 build-
ings 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.
*14
608 PRINCIPLES OF CHEMISTRY
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, CaX2, which
are distinguished by their stability in comparison with the salts
MgX2 ; neither does lime so easily form basic and double salts as
magnesia.
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
43 The act of heating brings the substance into that state of internal motion which ia
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 Kiihlmann is particularly instructive in this sense. Anhydrous (perfectly
dry) barium oxide does not react with monohydrated sulphuric acid, H2S04 (containing
neither free water nor anhydride, SOs). 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 th«
following experiment. Lime, or barium oxide, is placed in a flask or retort having au
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 tem-
perature 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.
THE VALENCY AND SPECIFIC HEAT OF THE METALS 609
temperature. Just as every volatile substance Las 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 be 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 carbon-
ate 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 (combi-
nation) cannot proceed because the second (decomposition, dissociation)
45 Experience has shown that by moistening partially-burnt lime with water and re-
heating 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.
41 Before the introduction of Deville's theory of dissociation, the modus operandi of
decompositions like that under consideration was understood in the sense that decompo-
sition starts at a certain temperature, and that it is accelerated by a rise of temperatun,
bat 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 complete-
ness 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.
610 PRINCIPLES OF CHEMISTRY
cannot take place, and thus all the most important phenomena with re-
spect to the behaviour of lime towards carbonic anhydride are explained
by starting from one common basis.45
Calcium carbonate, CaCCX, 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 cleav-
age 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 aragon-
ite ; 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
43 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, CaCO3,5H2O (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 hand, it is probable that an acid salt, CaHjfCOj)^, 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 atmo-
sphere) dissolves T086 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 8 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.
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 distribu-
tion (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
THE VALENCY AND SPECIFIC HEAT OF THE METALS 611
Calcium sulphate in combination with two equivalents of water,
CaS04,2H2O, 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, CaSO4,
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-trans-
parent gypsum, or alabaster, is often carved into small statues. Besides
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 laevo 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 crys-
tallises with the progress of time. The colloidal state of calcium carbonate is par-
ticularly 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).
«7 According to Le Chatelier (1888), 1JH20 is lost at 120°— that is, H20,2CaSQ4 is
formed, but at 194° all the water is expelled. According to Shenstoue and Cundall
(1888) gypsum begins to lose water at 70° in dry air. The semi-hydrated compound
H,2O,2CaSO4 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, <tc.)
612 PRINCIPLES OF CHEMISTRY
•which an anhydrous calcium sulphate, CaSO4, 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 ignit-
ing gypsum. If gypsum be very strongly heated it shrinks and loses
its power of combining with water.48 bis One part of calcium sulphate
requiresat 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
As lime is a more energetic base than magnesia, so calcium chloride,
CaCl2, 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, CaCl2,6H2O, which melts
at 30° 50
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 desicated at 120°, easily gives super-
saturated solutions with respect to CaSO4,2H2O, which contain as much as 1 part of
CaSC>4 to 110 parts of water. Boiling dilute hydrochloric acid dissolves gypsum,
forming calcium chloride. The behaviour of gypsum towards the alkaline carbonates
lias been described in Chapter X. Alcohol precipitates gypsum from its aqueoua
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 sulphate, 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, Struve,
and others obtained the salt CaKr^SO^sjHjO ; 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, NajCatCOj^BHsO,
"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(NO3)jCa(OH)2,24H2O by leaving powdered caustic lime in a saturated solution of
Ca(NC>3)j until it solidified. This salt is decomposed by water.
50 Calcium chloride has a specific gravity 2'20, or, when fused, 2'12, and the sp. gr. of
the crystallised salt CaClj.CHjO 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. CaCl2 boils
at 130°, 70 p.c. at 158°. Superheated steam decomposes calcium chloride with more diffi-
culty than magnesium chloride and with greater ease than barium chloride (Kuhnheim).
Sodium does not decompose fused calcium chloride even on prolonged heating (Liea-
Bodart), but an alloy of sodium with zinc, lead, and bismuth decomposes it, forming an
THE VALENCY AND SPECIFIC HEAT OF THE METALS 613
Just as for potassium, K = 39 (and sodium, Na — 23), there are
Jhe near analogues, Rb = 85 and Cs = 133, and also another, Li = 7, so
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
absoibs 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
v(1889), and found that CaCls,6H2O melts at 300-2, 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 CaCl2,4H2O£ are formed, which at temperatures above 88-4° are converted
into the crystallo-hydrates CaCl2,2H2O, whilst at temperatures below 18° the £ variety
passed into the more stable CaCl2)4H3Oa, 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 |8, which is easily produced but is unstable, and
the 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° 80° 40° 60°
CaCl2,6H2O 60 75 100 (102'8)
CaCl2,4H2Oa 90 101 117.
CaCl2,4H2O0 104 114
CaCl2,2H2O (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
80° for the first two salts and about 45° for the salts with 4H2O and 2H2O. The crystals
CaCl2,2H2O may, however, be obtained (Ditte) at the ordinary temperature from solu-
tions 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 CaCl2,H2O 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.CaCl2,4H2O/3, we will give the temperatures t at
which the passage of one hydrate into another takes place and at which the solution
CaCl2 + nH2O, 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
-65° 14-5 ice CaCl2,6H2O 0
+ 29-8° 6'1 CaCl2,6H2O CaCl2,4H20 6'8
45-3° 4-7 CaCl2,4H2O CaCl2)2HaO 11'8
175-5° 2-1 CaCl2,2H2O CaCl2,H2O 842
260° 1-8 CaCl2,H2O CaCl2 Several atmospheres
Solutions of calcium chloride may serve as a convenient example for the study of th6
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 CaCl2,6H2O.
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
614 PRINCIPLES OF CHEMISTRY
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 caesium 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 acquain-
tance 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, BaSO4) 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, SrSO4, 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 (BaCO3 forms
the mineral witherite ; SrC03, 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 BaSO4
4-4C=BaS + 4CO. The residue is then treated with water, in which
the barium sulphide is soluble.52 When boiled with hydrochloric acid,
crystallo-hydrate CaClj^H^O separates, and that a solution of the above composition
(CaCl2,14H2O requires 44'0 parts calcium chloride per 100 of water) solidifies as a cryo-
hydrate at about —55°.
61 The action of barium sulphate on sodium and potassium carbonates is given on
p. 437.
M Barium sulphide is decomposed by water, BaS + 2H20 = H2S + Ba(OH)2 (the reac-
tion is reversible), but both substances are soluble in water, and their separation is com-
plicated 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 anhy-
THE VALENCY AND SPECIFIC HEAT OF THE METALS 615
fcariutn chloride, BaCl2, is obtained in solution, and the sulphur is dis-
engaged as gaseous sulphuretted hydrogen, BaS + 2HCl=BaCl2-f H2S.
In this manner barium sulphate is converted into barium chloride,93
and the latter by double decomposition with strong nitric acid or
nitre gives the less soluble barium nitrate, Ba(N03)2,64 or with sodium
dride, 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.2O2 and SrH202 are prepared on a large scale, being applied to many reactions ; for/
example, strontium hydroxide is prepared for sugar works for extracting crystallisablel
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 Grouveu, by heating a mixture of charcoal and
strontium sulphate with magnesium and potassium sulphates, showed the easy decom-
posability depending on the formation of double salts, such as SrS,K2S, which are easily
soluble in water, and give a precipitate af strontium carbonate with .carbonic anhydride.
In such examples as these we see that the force which binds double salts may play apart
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 BaSO4 after being diluted with water.
55 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 decom-
positions 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, MnCl2, 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 man-
ganese. 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 fixe ' or ' permanent
white.
The solution of one part of calcium chloride at 20° requires T86 part of water, the
solution of one part of strontium chloride requires 1*88 part of water at the same tem-
perature, 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 stron-
tium crystallise out from solution with great ease in combination with water ; they form
BaCl2,2H2O and SrCl2,6H2O. The latter (which separates out at 40°) resembles the
salts of Ca and Mg in composition, and £tard (1892) obtained SrCl3,2HyO from solutions
at 90-180'. We may also observe that the crystallo-hydrates BaBr^HgO and Bal^H^O
are known.
6! The nitrates 8r(NOs)2 (in the cold its solutions give a crystallo-hydrate containing
4H2O) and Ba(N03)2 are so very sparingly soluble in water that they separate in consider-
able quantities when a solution of sodium nitrate is added to a strong solution of either
616 PRINCIPLES OF CHEMISTRY
carbonate a precipitate of barium carbonate, BaCO3. 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, BaH2O2,8H20,
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, BaS04. 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 poisouous.
Barium nitrate is directly decomposed by the action of heat, barium
oxide being left behind. The same takes place with barium car-
bonate, 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 reten-
tion of the water, although it parts with it when strongly ignited.55 bu
With oxygen the anhydrous oxide gives, as already mentioned iu
barium or strontium chloride.- They are obtained by the action of nitric acid on the carbon-
ates 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 3'6 7'4 18'8 90'8
SrO O'S 0'7 1-4 8 9
Supersaturated solutions are easily formed.
The anhydrous oxide . BaO fuses in the oxyhydrogen flame. When ignited in th»
vapour of potassium, the latter takes up the oxygen ; whilst in chlorine, oxygen is sepa-
rated and barium chloride formed.
Mbii Brugellmann, by heating BaH202 in a graphite or clay crucible, obtained BaO
to needles, sp. gr. 5'82, and by heating in a platinum crucible-Tin crystals belonging to
THE VALENCY AND SPECIFIC HEAT OF THE METALS 617
Chapters III. and IV., a peroxide, BaO2.56 Neither calcium nor
.strontium oxides are able to give such a peroxide directly, but they
form peroxides under the action of hydrogen peroxide.
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, caesium) increase with the atomic
•weight, and this succession clearly shows itself in many of their corre-
sponding compounds. Thus, for instance, the solubility of the
hydroxides RH2O2 and the specific gravity57 rise in passing from
calcium to strontium and barium, while the solubility of the sulphates
the cubical system, sp. gr. 5'74. SrO is obtained in the latter form from the nitrate*
JFhe following are the specific gravities of the oxides from different sources :—
MgO CaO SrO
from RN,Ofi 3'38 8'25 4'76
„ RCO3 3-48 3'26 4'45
„ RH2OZ 8-41 8-25 4'57
86 The property of barium oxide of absorbing oxygen when heated, and giving the
peroxide, BaO2, 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 per-
oxide 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,,O4 (obtained by Professor Scheme), which in course of time evolves oxygen (Chap-
ter IV., Note 21).
47 Even in solutions a gradual progression in the increase of the specific gravity shows
itself, not only for equivalent solutions (for instance, RC12 + 200HjO), but even with ao
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 determina-
tions) :
BeCla : S = 9,992 + 6T21p + O'lllp*
CaCLj : S = 9,992 + 80'24p + 0'476p*
SrCla : S = 9,992 +85'57p + 0-783pJ
BaClj : S = 9,992 + 86'56p + 0'813|>a
618 PRINCIPLES OF CHEMISTRY
decreases,58 and therefore in the case of magnesium and beryllium, as
metals whose atomic weights are still less, we should expect the solu-
bility of the sulphates to be greater, and this is in reality the case.
Just as in the series of the alkali metals we saw the metals potas-
sium, rubidium, and cjesium 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), «fec.59 Beryllium was so named because it occurs
in the mineral beryl. The metal is also called glucinum (from the
Greek word yAv/cv's, ' 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 consider-
able 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 : Al203,3BeO,6Si02. 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
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.
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 = 18'5,
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, Avdc'eff
(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 wan 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
Avde'eff's view corresponded with the truth — that is, that beryllium is bivalent,
which therefore necessitated the denial -of its tri valency. 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 — deter-
mining the vapour density of BeCl2 ( = 40, Chapter VII., Note 21), which gave an
Undoubted proof of its bivalency 'tee also Note 8).
THE VALENCY AND SPECIFIC HEAT OF THE METALS 619
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 "Wohler and Bussy. Wohler
obtained metallic beryllium (like magnesium) by acting on beryllium
chloride, BeCl2, with potassium (it is best prepared by fusing
K2BeF4 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 diffi-
cultly 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
80 Beryllium oxide, like aluminium oxide, is precipitated from solutions of its salts
by alkalis as a gelatinous hydroxide, BeH3O2, 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 accord-
ing 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 jO2.
Another characteristic of the salts of beryllium is that they give with aqueous am-
monia 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 BeSO4,4H20. 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 hydro-
•chloric acid.
620 PRINCIPLES OF CHEMISTRY
at a red heat, but gaseous hydrochloric acid is decomposed by it when
slightly heated, with evolution of hydrogen and development of a con-
siderable 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.
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 . R3N, so the nitrides
of magnesium (Note 14), calcium, strontium, and barium, have the
composition R3N2, for example, Ba3N2, as might be expected from the
diatomicityof 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, Ba3N2 + 2CO =* Ba(CN)2 + 2BaO.61
The metals of the alkaline earths, just like Na and K, absorb
hydrogen under certain conditions, and form pulverulent easily oxidis-
able metallic hydrides, whose composition corresponds exactly to that
of Na2H and K2H, with the substitution of K2 and Na2 by the atoms
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 VALENCY AND SPECIFIC HEAT OF THE METALS 621
Be, Mg, Ca, Sr, and Ba. The hydrides of the metals of the alkaline
earths were discovered by 0. 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 hydrogen62 when heated, so that these hydrides are distin-
guished for their considerable stability under heat, but they oxidise
very easily.63
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.
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, 1804.)
It is probable that in the free metals of the alkaline earths hitherto obtained a portion
was frequently in combination with nitrogen and hydrogen.
c3 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 J 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 83 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 C2M (Chapter VIII.,
Note 12 bis), for instance, C2Ba, 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 car-
bon. The resultant compounds, C2M, evolve acetylene, C2H2, with HC1, just as N2Mj
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
•xtending our knowledge of their reactions prove to be of great interest.
.
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