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time being is not functioning in its highest recognised valency,
often exhibits a readiness to unite with additional atoms to form

64 Introductory Outlines

new compounds : thus ammonia combines eagerly with hydro-
chloric acid, forming ammonium chloride

NH 3 + HC1 = NH 4 C1.

Carbon monoxide unites directly with chlorine to form carbonyl


Carbon monoxide also combines with an additional atom of
oxygen, and gives carbon dioxide, thus

2CO + 2 = 2C0 2 .

In this last action it will be seen that the molecule of carbon
monoxide, in being converted into the dioxide, takes up one atom
of oxygen ; but as the molecule of oxygen is the smallest isolated
particle, it follows that the two atoms contained in such a molecule
must first separate, and each one then furnishes the requisite
additional oxygen for one molecule of carbon monoxide. In the
union of carbon monoxide with chlorine, and of ammonia with
hydrochloric acid, are we to suppose that the same action takes
place ? That is to say, do the two atoms in the molecule of
chlorine separate from each other and unite with carbon, thereby
satisfying its tetrad valency, in the manner here expressed ?


Cl C1 + CO= >C = 0.

And in the case of ammonia and hydrochloric acid, do the
hydrogen and chlorine atoms part, and each unite with the
nitrogen atom, thereby raising it from the trivalent to the penta-
valent condition ? thus

Cl H

H Cl + H N H = H N H.

! I


Valency of the Elements 65

Or are we to suppose that the two molecules, without losing their
integrity, become held together as independent molecules, by
virtue of the unsatisfied affinities of the carbon, or the nitrogen,
as the case may be, in which case the compounds might be repre-
sented thus

Cl H Cl

|C = H N H.

Cl |


This question would be settled by determining the vapour-
density of the compound. If, for instance, we were to find the
vapour-density of ammonium chloride to be 26.75, then the com-
pound having the composition NH^Cl would have the normal
molecular volume, that is, its molecule would occupy two unit
volumes,* and the conclusion would be that the vapour consisted
of single molecules of the composition represented by the formula
NH 4 C1. But ammonium chloride at ordinary temperatures is a
solid, and when heated to the temperature necessary to convert it
into vapour its molecules break up into separated molecules of the
two original gases ammonia, NH 3 , and hydrochloric acid, HCl.t
So that we are unable to gain any information in this direction
as to the mode in which the atoms are disposed in the compound.
When the two gases are brought together under ordinary con-
ditions, they combine with the evolution of considerable heat,
owing to loss of energy ; this is taken as evidence that true
chemical action, in the sense of atomic rearrangement, has re-
sulted, hence it is believed that in this compound the nitrogen
is united with the five monovalent atoms, and consequently is

In the case of carbonyl chloride, COC1 2 , the vapour-density can
be ascertained, this compound existing in the gaseous condition
at the ordinary temperature. Its vapour-density, determined by
experiment, is found to be 50.6. This number, divided into the
molecular weight of the compound having the composition COC1 2 ,
gives practically the number 2 as the molecular volume of the
compound. Hence we conclude that these four atoms constitute
a single molecule.

There are a number of combinations, however, in which mole-

* See p. 43. f See Dissociation, p. 89.


66 Introductory Outlines

cules of different compounds unite, that do not so readily admit of
explanation, because in neither of the molecules is there any
atom functioning in a lower state of valency than that which
it is known to be capable of. For example, the monovalent
elements fluorine and hydrogen form the compound hydrofluoric
acid, HF; fluorine also combines with the monovalent element
potassium, forming potassium fluoride, KF. Both of these com-
pounds come under the head of saturated compounds, in the sense
that neither of them contains an atom which is known to be
capable of exercising a higher valency than it exhibits in these
compounds. Nevertheless these two molecules unite together and
form a definite chemical compound, known as hydrogen-potassium

Again, the divalent element zinc combines with two atoms of
the monad element chlorine, forming zinc chloride, ZnCl 2 ; the
two monovalent elements sodium and chlorine also combine,
giving the compound sodium chloride, NaCl. Both of these
substances must be regarded as saturated compounds, and yet
they unite with each other, forming a distinct chemical compound,
known as sodium zinc chloride. Such compounds as these are
known as double salts, and examples might be multiplied almost
indefinitely. A similar union of molecules, where the recognised
valency of the atoms is all satisfied, is seen in a large number
of compounds containing water of crystallisation ; * for example,
the divalent element copper, in combination with two atoms of
chlorine, forms cupric chloride, CuCl 2 . The divalent element
oxygen, in combination with two hydrogen atoms, forms water,
H 2 O. When cupric chloride crystallises from aqueous solution,
each molecule of the chloride unites to itself two molecules of
water, which is therefore termed water of crystallisation.

In chemical notation, it is usual to represent compounds of this
order by placing the formulae of the different molecules that have
entered into union in juxtaposition, with a comma between ;
accordingly, the examples here quoted would be indicated thus

Hydrogen potassium fluoride . . HF,KF.
Sodium zinc chloride .... ZnCl 2 ,NaCl.
Crystallised cupric chloride . . CuCl 2 ,2H 2 O.

Combinations of this order are by no means confined to the
* See page 216.

Valency of the Elements 67

union of two kinds of molecules, as the following examples will
serve to show :

Platinum sodium chloride . . PtCl 4 ,2NaCl,6H 2 O.
Mercuric potassium chloride . 2HgCl 2 ,KCl,2H 2 O.

At the present time our knowledge of the nature of the union
between these various molecules is too imperfect to admit of any
precise explanation ; such compounds are frequently distinguished
as molecular combinations.

It is quite possible that the unit which has been adopted for estimating
valency, namely, i monovalent atom, is after all only an extremely rough and
crude measure, which is incapable of appreciating smaller differences of com-
bining capacity that may, and most probably do, exist. Its use may be com-
pared to the adoption of a single unit, say I gramme, for the estimation of
mass or weight ; when, if a given quantity of matter has a weight equal to i
gramme, but less than 2 grammes, its weight would be i; if greater than 2
grammes, but less than 3, then its weight would be 2 a method of estimating
which tacitly assumes that no intermediate weights of matter between the
various multiples of the selected unit are possible. There is no evidence to
show that the combining capacity of an element is exactly expressed by simple
multiples of a monovalent atom.

For example, i hydrogen atom unites with i chlorine atom, that is to say,
with a mass of chlorine weighing 35.5 times its own weight ; and we say that
the mutual affinities of these atoms are satisfied. But for anything we know
to the contrary, an atom of hydrogen may have an affinity for chlorine which
would enable it to unite with a mass of chlorine weighing 40 or 45 or 50 times
its own weight, but not a mass weighing 71 (35.5 x 2) times its own. But since
a mass of chlorine 35.5 times the weight of a hydrogen atom is the smallest
quantity that is ever known to take part in a chemical change, is the chemically
indivisible mass we call an atom, it follows that as the hydrogen atom has not
sufficient combining capacity to unite with 2 atoms, it is compelled to be
satisfied with i. It might still, however, retain a residual combining capacity.
Or the residual combining capacity may be lodged in the chlorine atom,
which may be conceived as being able to unite with a greater weight of
hydrogen than is represented by i atom, but not so much as that of 2

Each of the elements fluorine, chlorine, bromine, and iodine unites with
i atom of hydrogen, and we represent their compounds in a similar manner,

H - F ; H - Cl ; H - Br , H - I ;

but we make an enormous assumption if we suppose that in each of these
compounds the mutual affinities of the atoms is equally satisfied.

The trend of modern thought, however, lies in the direction of an electrical
interpretation of valency. The fact that atoms are always associated with
fixed and definite charges of electricity, that valency, indeed, could be measured
in terms of electric units (the outcome of Faraday's Law, chap, xi.) seemed

68 Introductory Outlines

at one time only to emphasise the difficulty of explaining such cases as those
above mentioned ; but the mo:~e recent developments in this region of physics
have led to modified views as to the nature of the bond which unites atoms
together. Stated in briefest outline, this chemical "bond" or unit of affinity,
which forme: ly has been regarded in the light of a single line of force a
fi action of a bond being considered as altogether inadmissible io now rega ded
as a. bundle of lines of force (a Faraday bundle). Under appropriate condi-
tions, such as the proximity of suitable molecules or ions, it is conceived that
some strands of the bundle may become loosened from one of the attached
atoms and thus become available for attraction by similar wandering strands
from other molecules. Obviously, therefore, this view admits of practically an
unbroken gradation in degrees of chemical affinity. Instead, therefore, of
residual affinity, we have varying fractions of the total bundle of lines of
force which in its entirety constitutes the chemical " bond."

A modification of this view, recently advanced by Sir W. Ramsay,* substitutes
electrons f for this ' ' bundle of lines of force. " Atoms are regarded as carrying
with them a " reserve of electrons," electrons which may be inactive, or latent.
1'hus taking ch.orine as an example, he says: "It appears likely that each
atom of chlorine carries with it no fewer than seven electrons, . . . latent as it
wrre, not revealing themselves in such a compound as common salt. . . .
These valencies are manifested in such compounds as perchloric acid."

* Presidential address, Chem. Soc., 1909. f See page 104.


UNDER the head of the general properties of gases it will be con-
venient to consider the following subjects : *

1. The relation of gases to heat.

2. The relation of gases to pressure.

3. The liquefaction of gases.

4. Diffusion of gases.

5. The kinetic theory of gases.

The Relation of Gases to Heat. The fact that substances
expand when heated, and again contract upon being cooled, was
observed in very early times. The fact also that all substances do
not undergo the same alterations in volume when subjected to the
same changes of temperature has been Jong known ; but it was not
until the beginning of the nineteenth century that it was proved by
Charles and Gay-Lussac that all gases expanded and contracted
equally when exposed to the same alterations of temperature.
This law is generally known as the Law of Charles, and may be
thus stated : When a gas is heated, the pressure being constant, it
increases in volume to the same extent whatever the gas may be.

The increase in bulk suffered by I volume of a gas in being
heated from o to i is termed the coefficient of expansion, and if
the law of Charles is true all gases will have the same coefficient.

Modern research has shown that the law of Charles is not abso-
lutely true, and the extent to which gases deviate from the strict
expression will be seen from the coefficients of expansion given in
the following table :

* The study of these subjects belongs more especially to the science of
physics or chemico-physics. For fuller information on these points than can
be included within the scope of this book students are referred to specicJ
treatises on physics.


7O Introductory Outlines

Air 003665^

Hydrogen 003667 I

Carbon monoxide .... .003667 j

Nitrogen 003668^

Nitrous oxide . .003676

Carbon dioxide 003688

Cyanogen ...... .003829

Sulphur dioxide 003845

It will be noticed that the first four gases have almost the same
coefficient of expansion : these gases are all very difficult of lique-
faction, and it will be seen that the coefficient rapidly rises in the
case of the other gases, which are easily liquefied.

For purposes of ordinary calculation it is usual to adopt the
coefficient of expansion of air as applicable to all gases. It will
be obvious that since the volume of a gas is affected by alterations
of temperature, it becomes necessary, when measuring the volume
of a gas, to have regard to the particular temperature at which the
measurement is made, and in order to compare volumetric measures
they must be all referred to some standard temperature. This
standard temperature is by general consent o C.

Taking the fraction .003665, therefore, for the coefficient

i volume of a gas at o becomes I + .003665 volumes at i

I o i + .003665 x 2 2

or i ., o i -1- .003665 / /

Therefore the volume at f equals the volume at o multiplied by
i + .003665 /. Let v be the volume at /, and v the volume at o,

V V (l + .003665/),

and conversely the volume at o equals the volume at t divided by
i + .003665 t

I + .003665 1

The vulgar fraction equivalent to .003665 is %\^. 273 volumes
at o become 273 -f / at /.

What is known as the absolute temperature of a substance is the
number of degrees above 273 C. Taking this point as the zero,
the absolute temperature of melting ice, for example, will be 273.
Charles' law, therefore, may be thus stated : The "volume of an}

Relation of Gases to Pressure 71

gas, under constant pressure, is proportional to the absolute tem-

The Relation of Gases to Pressure. The effect of increase
of pressure upon a gas is to diminish its volume. The law which
connects the volume occupied by a gas, with the pressure to which it
is subjected, was discovered by Robert Boyle (1661), and is known
as Boyle's Law. It may be thus stated : The vohtme occupied by
a given weight of any gas is inversely as the pressure. The
general truth of this law may readily be illustrated by subjecting a
gas to varying pressures, and it will be seen that when the pressure
is doubled the volume of gas is reduced to one-half, and so on.

Just as in the case of the law of Charles, modern investigations
have shown that the law of Boyle is not a mathematical truth. It
is found not to be absolutely true of any gas, for, with the exception
of hydrogen, all gases are more compressible than is demanded by
the law. Hydrogen deviates from the law in an opposite sense, in
that it requires a higher pressure than the law would indicate, in
order to reduce a volume of it to a given point. These deviations
from Boyle's law are explained by the operation of two causes ;
first, the attraction exerted by gaseous particles upon each other ;
second, the fact that increased pressure diminishes the space
between the molecules, and not the actual space occupied by the
molecules of a gas. When the former cause predominates, the
gas deviates from the law by being more compressible ; in the case
of hydrogen the second cause operates more powerfully. (See
Kinetic Theory of Gases.) For ordinary purposes of calculation
the law of Boyle may be regarded as true.

As the volume of a given weight of gas is so intimately related
to the pressure, and as the atmospheric pressure is variable, it
becomes necessary, in all quantitative manipulation with gases, to
know the actual pressure under which the gas is at the time of
measurement, and to refer the volume to a standard pressure.
The pressure that has been adopted as the standard is that of a
column of mercury 760 mm. in height. (See Atmosphere.)

If 77 equals the volume of gas measured at p pressure, and v t
the volume at the standard pressure, then

V = -.

In practice it is most usual to make both correction for tempe-

7 2 Introductory Outlines

rature and pressure together ; then i> being the volume at the
standard temperature and pressure, we get

V i +.003665 t ' 760 '

The Liquefaction of Gases. Under certain conditions of tem-
perature and pressure, the law of Charles and the law of Boyle both
completely break down. According to
the law of Charles, 100 c.c. of a gas at
o C. should occupy 96.4 c.c. if the tem-
perature were lowered to 10. If 100
c.c. of the gas sulphur dioxide at o C.
be confined in a glass tube standing in
mercury, and the gas be cooled to 10
by surrounding the tube with a freezing
mixture, it will be found that the volume
of gas, instead of occupying 96.4 c.c.,
has been reduced to a few cubic centi-
metres only, and that the surface of the
mercury in the tube is wet owing to the
presence of a minute layer of a colourless
liquid upon it. In this case the law of
Charles has broken down, and the sul-
phur dioxide has passed from the gaseous
\^V to the liquid state.

Similarly, according to the law of
Boyle, loo c.c. of a gas measured at the
standard pressure should occupy 25 c.c.
when exposed to a pressure of four additional atmospheres. If
100 c.c. of the gas sulphur dioxide be enclosed in one limb of a long
(j-tube, as shown in Fig. i, the other limb being filled with air,
and the two gases be simultaneously exposed to increased pressure
by raising the mercury reservoir, it will be seen that at first the
gases in both tubes are compressed equally. As the pressure
approaches three atmospheres, however, the mercury will be seen

* The student should familiarise himself with the method of calculating the
changes of volume suffered by gases, by changes of temperature and pressure,
by working out a number of examples such as the following :

1. If 30 litres of gas are cooled from 25 to o, what is the diminution in
volume, the pressure being constant? Ans. 2.51 litres.

2. If a litre of air at o weighs 1.293 grammes when the barometer is at

FIG. i.

Liquefaction of Gases 73

to rise much more rapidly in the tube containing the sulphur
dioxide, and when the mercury reservoir has been raised to such a
height that the gases are subjected to four atmospheres, the sulphur
dioxide will have completely broken down, and will be entirely con-
verted into a few drops of liquid, which appear upon the surface of
the mercury. The air meantime, in the other limb, will be found to
occupy 25 c.c., as that gas at that pressure obeys Boyle's law almost
absolutely. We see, therefore, that at a certain temperature and at
a certain pressure the gas sulphur dioxide begins rapidly to depart
from the laws of Charles and Boyle, and ultimately passes into the
liquid condition.

All gases, when exposed to certain conditions of temperature
and pressure, conditions which are special for each different
gas, will pass from the gaseous to the liquid state ; and
as the point at which liquefaction takes place is approached,
the departures from Boyle's law become more and more pro-

The first substance, recognised as being under ordinary condi-
tions a true gas, that was transformed into the liquid condition
was chlorine, which was liquefied in the year 1806 by Northmore.
The true nature of this liquid was
not understood until Faraday inves-
tigated the subject.

In his earlier experiments Fara-
day's method consisted in sealing
into a bent glass tube (Fig. 2) sub-
stances which, when heated, would
yield the gas ; the substances being
contained in one limb of the tube,
and the empty limb being immersed FIG. 2.

in ice. The pressure exerted by the gas thus generated in a con-
fined space was sufficient to cause a portion of it to condense to

760 mm., what will be the weight of a litre of air at 27, the barometer
standing at the same height? Ans. 1.177 grammes.

3. What will be the weight of a litre of air at 42 when the barometer stands
at 735 mm. ? Ans. 1.084 grammes.

4. Air at a temperature of 15 is enclosed in a vessel and heated to 93.
Compare the pressure of the enclosed air with that of the atmosphere. Ans.
As 61 : 48.

5. What will be the volume, at the standard temperature and pressure, of
500 c.c. of hydrogen, measured at 20, and under a pressure of 800 mm.?
Ans. 490 c.c.

74 Introductory Outlines

the liquid state, and the liquid collected in the cooled limb. In
this way Faraday liquefied such gases as chlorine, sulphur dioxide,
ammonia, cyanogen. In his later experiments Faraday compressed
the gas by means of a small compression pump, and at the same
time applied a low degree of cold, and by so doing he succeeded
in liquefying carbon dioxide, hydrochloric acid, nitrous oxide, and
other gases. There were a number of gases, however, which Fara-
day found it impossible to liquefy, such as hydrogen, oxygen, nitro-
gen, marsh gas, nitric oxide, carbon monoxide, &c. It became the
custom to call these permanent gases, and this term was applied to
them until the year 1877.

In that year it was proved by Pictet, and independently by Cail-
letet, that under sufficiently strong pressure, and a sufficiently low
degree of cold, the so-called permanent gases could in the same
way be reduced to the liquid condition. Pictet's method was in
principle the same as that employed by Faraday, the difference
being that with the machinery at his disposal he was able to
employ enormously increased pressure and a greater degree of
cold. For the liquefaction of oxygen, a quantity of potassium
chlorate was heated in a strong wrought-iron retort, to which was
connected a long horizontal copper tube of great strength and small
bore. At the extreme end of this tube there was a pressure gauge
capable of indicating pressures up to 800 atmospheres, and a stop-
cock. The tube was cooled by being contained in a wider tube,
through which a constant stream of liquid carbon dioxide, at a tem-
perature of 120 to 140, was caused to flow.

The machinery employed to maintain this flow of liquefied car-
bon dioxide was somewhat elaborate, consisting of condensing and
exhaust pumps for liquefying and rapidly evaporating sulphur
dioxide, and similar condensing and exhaust pumps for liquefying
and rapidly evaporating carbon dioxide : the sulphur dioxide being
merely the refrigerating agent used to assist the liquefaction of
the carbon dioxide. This machinery was driven by two eight-
horse-power engines. As the potassium chlorate was heated
and oxygen evolved, the internal pressure in the retort and
copper tube rapidly rose, and its amount was indicated by the

When the stop-cock upon the end of the tube was opened, liquid
oxygen was forcibly driven out in the form of a jet.

In the method employed by Cailletet, the pressure to which the
gas is subjected is obtained by purely mechanical means. The

Liquefaction of Gases


gas to be liquefied is introduced into a glass tube (Fig. 3), the
narrow end of which consists of a strong capillary tube. The tube
carries a metal collar, which enables it to be secured in position
in the strong steel bottle (Fig. 4), by means of a nut, E' (Fig. 5),
which screws into the mouth. The bottle, which is partially filled
with mercury, is connected by means of a flexible copper tube of
fine bore with a small hydraulic pump, by means of which water
is forced into the steel bottle. The water so driven in forces the

FIG. 3.

FIG. 4.

FIG. 5.

mercury up into the glass tube T, and thereby compresses the
contained gas. In this way a pressure of several hundred atmos-
pheres may be applied to the gas. In his earlier experiments
Cailletet depended almost entirely for the refrigeration he required
upon the fact, that when a gas is allowed suddenly to expand it
undergoes a great reduction in temperature. This method of
cooling may be termed internal refrigeration. In the case of
oxygen, the gas was first subjected to a pressure of 300 to 400

Introductory Outlines

atmospheres, and was then allowed suddenly to expand by a rapid
release of the pressure. The result of the sudden expansion was
to momentarily lower the temperature of the gas to such a point
that the tube was filled with a fog, or mist, consisting of liquid
particles of oxygen.

Online LibraryGeorge S NewthA text-book of inorganic chemistry → online text (page 7 of 67)