R. M. (Robert Martin) Caven.

The foundations of chemical theory; online

. (page 17 of 23)
Online LibraryR. M. (Robert Martin) CavenThe foundations of chemical theory; → online text (page 17 of 23)
Font size
QR-code for this ebook

solution, and the consequent hydrolysis of the precipitated mercurous

The series of the metals must be traversed far in the upward
direction before those metals are reached which form normal but
not basic carbonates. Such are the metals of the alkaline earths
and the alkalis. Calcium carbonate, for example, is normal when
precipitated, and never becomes basic. This salt is soluble to a
minute extent in pure water, to which it imparts a faintly alkaline
reaction. It is true that hydrolysis of the dissolved salt takes place,
but this results in the formation of hydroxide and bicarbonate, thus:
2CaCO 3 + H 2 O ^= 2Ca(HC0 3 ) 2 + Ca(OH) 2 ,

OH' ions showing alkalinity being due to the ionization of Ca(OH) 2 .
The same phenomenon appears more markedly in the case of alkali
carbonates, which show a strongly alkaline reaction. Thus, sodium
carbonate reacts with water in this way:

Na 2 C0 3 + H 2 ^= NaHC0 3 + NaOH,

and so its strongly alkaline reaction is explained; for, according to
the ionization theory, a salt producing basic and acidic ions only, e.g.
2Na- and CO 3 ", would be neutral in reaction.


It has been seen above that only the most powerful metals form
hydrogen or acid sulphates; the same is true in regard to hydrogen
or acid carbonates. Thus it is only the alkali metals that form solid
hydrogen carbonates, and these increase in stability from sodium to
caesium in the series:

NaHC0 3) KHC0 3 , RbHCO 3 , CsHCO 3 .

The alkaline earth metals, with ferrous iron and magnesium, form
unstable hydrogen carbonates in solution however, e.g. Ca(HCO 3 ) 2 ;
and there is this difference between these hydrogen carbonates in
solution, and the solid hydrogen carbonates of the alkalis, viz. that
the former are more soluble in water than the corresponding normal
carbonates, whilst the latter are less soluble, and are precipitated, as,
for example, in the case of NaHCO 3 , by passing carbon dioxide gas
into cold saturated solutions of the normal carbonates.

The superior solubility of calcium carbonate in water containing
carbon dioxide in solution over its solubility in pure water is of
profound importance in nature; for it is the cause, not only of the
temporary hardness of water, but of the disintegration of calcareous
rocks, as well as of their original formation through the agency of
marine organisms, which form their shells from calcium carbonate
held in aqueous solution by carbonic acid. These facts are repre-
sented by the following reversible reaction:

CaCO 3 + H 2 O + CO 2 ^= Ca(HCO 3 ) 2 .
Hydrated Salts.

Water of crystallization is of common occurrence in crystallized
salts, and since its presence has a great influence on physical pro-
perties, the student must on no account ignore it in formulat-
ing a salt. The influence of temperature and atmospheric con-
ditions on hydrated salts will be dealt with in another place; it
may here be remarked, however, that when such salts are coloured
the corresponding anhydrous compounds are invariably of a different
colour. Thus, for example:

CuSO 4 -5H 2 O is blue; CuSO 4 is white.

FeSO 4 -7H 2 O green; FeSO 4 white.

NiSO 4 -7H 2 O deep green; NiSO 4 yellow.

CuCl 2 -2H 2 O bluish green; CuCl 2 brown.

FeCl 3 -6H 2 O yellow; FeCl 3 iron-black.

CoCl 2 -6H 2 O crimson; CoCl 2 blue.

CoBr 2 '6H 2 O dark red; CoBr 2 green.

CoI 2 '6H 2 O dark red; CoI 2 violet.


The proportion of water varies much in different hydrated salts;
ammonium oxalate, for example, has 1 molecule of water to 1 mole-
cule of salt, ordinary sodium phosphate has 12, and the alums 24;
whilst between these extremes there are salts containing 2, 5, 6, 7,
and 10, and less frequently 3, 4, and 8 molecules of water. Occa-
sionally, too, the same salt will crystallize with varying proportions
of water according to the temperature of its formation. Thus, for
example, manganous sulphate, MnSO 4 , forms crystallo-hydrates with
1, 4, 5, and 7 molecules of water at different temperatures.

The way in which water is combined chemically in crystallo-
hydrates constitutes a problem the discussion of which is beyond
the scope of the present work; nevertheless it will be well to
tabulate here the commonest hydrated salts according to the mole-
cular proportions of water they contain:

iH 2 2CaS0 4 -H 2 0.

H 2 O Na 2 CO 3 -H 2 O; (NH 4 ) 2 C 2 O 4 -H 2 O.

2H 2 O BaCl 2 -2H 2 O; CuCl 2 -2H 2 O; CaSO 4 -2H 2 O.

3H 2 O K 4 Fe(CN) 6 -3H 2 O.

4H 2 O NaNH 4 HPO 4 -4H 2 O.

5H 2 O Na 2 S 2 O 3 -5H 2 O; CuSO 4 -5H 2 O; Bi(NO 3 ) 3 .5H 2 O.

6H 2 O CaCl 2 -6H 2 O; MgCl 2 -6H 2 O; CoCl 2 -6H 2 O; FeCl 3 -

6H 2 O; CrCl 3 -6H 2 O; FeSO 4 -(NH 4 ) 2 SO 4 -6H 2 O,

and similar double sulphates.
7H 2 , MgS0 4 -7H 2 0; ZnSO 4 -7H 2 0; FeSO 4 .?H 2 O;

NiSO 4 -7H 2 O; CoSO 4 -7H 2 O.
8H 2 O .. . Ba(OH) 2 -8H 2 O; BaO 2 -8H 2 O.

10H 2 O
12H 2 O
18H 2 O
24H 2

Na 2 CO 3 -10H 2 O; Na 2 SO 4 -10H 2 O; Na 2 B 4 O 7 -10H 2 O.
Na 2 HPO 4 -12H 2 O; Na 2 HAsO 4 .12H 2 O.
A1 2 (SO 4 ) 3 -18H 2 O.
K 2 SO 4 -A1 2 (SO 4 ) 3 .24H 2 O, and other alums.

Double and Complex Salts.

DOUBLE SALTS are those which have a definite chemical indi-
viduality in the solid state, but break up more or less completely
in aqueous solution into their constituent single salts. Crystallized
potassium alum, K 2 SO 4 A1 2 (SO 4 ) 3 24 H 2 0, for example, is un-
doubtedly a chemical compound, and not a mixture of its two
constituent salts; but when dissolved in water it gives the separate
reactions of aluminium and potassium sulphates, so that its solution
is a mixture of these two salts, the process of solution having been
accompanied evidently by disintegration of the double salt.

The formula for alum is sometimes halved, thus:

KA1(SO 4 ) 2 -12H 2 O.

(D60) 14


Now it is always wise to accept the simplest available formula
in default of evidence to the contrary; and there is no direct
evidence regarding the molecular magnitude of a solid alum. It
may be objected, however, that the above formula suggests a
complex rather than a double salt, since it does not show complete
molecules of the two constituent sulphates. This objection would
perhaps have little weight were it not for a peculiar change which
solid chromic alum undergoes when heated to 90 C. The violet
crystals then turn green, with loss of water, changing into a salt
which contains no free sulphate, since its solution gives no pre-
cipitate with barium chloride. The change is thus formulated,
ignoring water of crystallization:

K 2 S0 4 .Cr 2 (S0 4 ) 3 K 2 [Cr 2 (S0 4 )J or 2K[Cr(SO 4 )J.

So a double salt becomes a complex salt; potassium chromic
sulphate becomes potassium chromisulphate, the potassium salt of
chromisulphuric acid, HCr(SO 4 ) 2 , a compound which is actually
formed when chromic sulphate is warmed with sulphuric acid.

That K[Cr(SO 4 ) 2 ] is so different from K 2 SO 4 -Cr(S0 4 ) 3 .24H 2 O
is a good reason for not writing the formula for any alum in a way
to suggest relationship to the former of these compounds.

Double salts are very numerous. Besides the double sulphates
and isomorphous selenates there are double chlorides, bromides, and
iodides, and less frequently double carbonates and nitrates.

DOUBLE SULPHATES. The alums, and salts of which ferrous
ammonium sulphate is a well-known example, may be mentioned.

Alums are isomorphous salts of the type

M 2 'S0 4 .X 2 -(S0 4 ) 3 '24H 2 0,
where M = Na, K, NH 4 , Kb, Cs, Tl, Ag,
and X = Al, Fe, Cr, Mn, Ga, Ti, Rh.

They are formed by mixing the constituent salts in aqueous
solution, in proportions approximating to theoretical requirements,
and crystallizing. The alums are less soluble than their con-
stituent salts, and this is particularly the case with those of the
extremely electropositive metals rubidium and caesium.

Double sulphates of the ferrous ammonium sulphate type are
the salt FeSO 4 -(NH 4 ) 2 SO 4 -6H 2 and others in which Mg", Zn",
Cu", Mn", Co", Ni" may take the place of Fe", and other alkali
metals that of NH 4 . The relation between these double sulphates
and the heptahydrated sulphates, e.g. FeSO 4 7H 2 O, is interesting.


It was found by Graham that one of the seven molecules of
water in this salt required a higher temperature for its expulsion
than the other six. This seventh molecule Graham called con-
stitutional water, because it appeared to enter into the constitution
of the salt more intimately than the other six molecules. It is this
molecule which seems to be displaced by ammonium or other
alkali sulphate in the formation of the double salt. The relation-
ship may be thus shown:

FeSO 4 .H 2 O-6H 2 O : FeSO 4 .(NH 4 ) 2 SO 4 -6H 2 O.

It should be^remarked that the ammonium sulphate in the
double salt exerts a protective influence over the ferrous sulphate,
for ferrous ammonium sulphate is less oxidizable by the air than
ferrous sulphate, and for this reason is preferred for the purpose
of volumetric analysis.

DOUBLE CHLORIDES. The mineral carnallite is KC1 MgCl 2 6 H 2 0,
to which there corresponds the ammonium salt NH 4 ClMgCl 2 6H 2 O.
The solubility of magnesium and manganous hydroxides in am-
monium chloride solution, with the corresponding fact that the
hydroxides of these metals are not precipitated by ammonia in
presence of ammonium chloride, is sometimes attributed to the
formation in solution of complex ions, such as (MgCl 3 )', derived
from NH 4 ClMgCl 2 . The salts themselves, however, are usually
regarded as double rather than complex salts. The double chloride
NaCl'AlCl 8 is a volatile compound, the formation of which was a
part of an early process for the preparation of metallic aluminium.

Examples of double salts of another type are sodium potassium
tartrate /Rochelle salt), NaKC 4 H 4 6 4H 2 O, microcosmic salt,
NaNH 4 HP0 4 4H 2 0, and magnesium ammonium phosphate,
MgNH 4 P0 4 -6H 2 0.

These are formulated differently from the alums and other
double salts, as containing two or more metallic radicles within
the same molecule. Since, however, these salts show no complex
ions in solution, and their molecular magnitudes are unknown, it
may be that they should be put in the same category as other
double salts.

COMPLEX SALTS are those which, derived originally from single
salts, are so stable as to maintain their individuality in solution,
one of the metals appearing as a basic ion, whilst the other has
become part of a complex acidic ion, so that its metallic nature is


masked. Potassium ferrocyanide, K 4 Fe(CN) 6 , is a familiar example
of a complex salt. It appears to be composed of 4KCN + Fe(CN) 2 ,
and is indeed formed by adding potassium cyanide to ferrous sul-
phate solution until the precipitated cyanide has been redissolved,
and then boiling the solution. Thus a remarkable change takes
place; the iron ceases to behave as a basic radicle and becomes
part of an acidic complex, so that it gives no ferrous reactions in
solution. No ferrous salt is present, only a potassium salt potas-
sium ferrocyanide which ionizes in solution thus:

K 4 Fe(CN) 6 :^= 4K' + [Fe(CN) 6 ]"".

So profound is this change, and so stable the complex salt, that
from its concentrated solution sulphuric acid separates hydroferro-
cyanic acid, H 4 Fe(CN) 6 , as a white solid.

Alum and potassium ferrocyanide, as representatives of double
and complex salts respectively, present extremes, but there are
gradations between them. The behaviour of nickel and cobalt
salts with potassium cyanide furnishes a case in point. The fol-
lowing reactions take place:

NiS0 4 + 2KCN = Ni(CN) 2 + K 2 SO 4 ; Ni(CN) 2 + 2 KCN = K 2 Ni(CN) 4 .
CoS0 4 + 2KCN = Co(CN) 2 + K 2 S0 4 ; Co(CN) 2 +4KCN = K 4 Fe(CN) 6 .

Both K 2 Ni(CN) 4 and K 4 Co(CN) 6 are complex rather than double
salts, for they do not contain nickelous and cobaltous ions; more-
over, K 4 Co(CN) 6 is plainly analogous to K 4 Fe(CN) 6 . From each of
these solutions, however, the simple cyanide Ni(CN) 2 or Co(CN) 2
is reprecipitated by dilute acid. These are examples of complex
salts, therefore, which are less stable than ferrocyanide. When a
solution of potassium cobaltocyanide is boiled in presence of air it
undergoes oxidation to cobalticyanide thus:

2 K 4 Co(CN) 6 + H 2 O + O = 2 K 3 Co(CN) fl + 2 KOH,

and this latter salt is much more stable than cobaltocyanide, in this
respect resembling ferro- or ferricyanide. The fact that nickel
forms no such stable complex salt, nickelic salts being unknown,
underlies the well-known separation of these two metals.

The student meets with other examples of complex acids and
salts in the course of chemical analysis. Hydrofluosilicic acid,
HgSiFg, is evidently composed of 2HF + SiF 4 , but it contains the
complex ion [SiF 6 ]". Potassium platinichloride, or chloroplatinate,


K 2 PtCl 6 , and the corresponding acid H 2 PtCl 6 , formed when platinum
is dissolved in aqua regia, are of the same type, and so is the
corresponding stannichloride, K 2 SnCl 6 . Potassium cobaltinitrite,
K 3 Co(NO.J 6 , formed as a yellow crystalline precipitate when potas-
sium nitrite is added to a cobaltous solution acidified with acetic
acid, is of the same type as K :i Co(CN) 6 and K 3 Fe(CN) 6 .

Ammonium phospho-molybdate is a complex salt of a different
kind, in which from 10 to 14 molecules of MoO 3 are combined with
(NH 4 ) 3 PO 4 . It is formed in presence of nitric acid, and when
dissolved by ammonia suffers hydrolysis into simple phosphate
and molybdate. Potassium antimonyl tartrate, or tartar emetic,
[KSbOC 4 H 4 O 6 ] 2 H 2 O, is a complex rather than a double salt, for
it dissolves in water without hydrolysis, which antimonious salts
will not do. It is therefore best regarded as the potassium salt of
antimonyl-tartaric acid, [K(SbOC 4 H 4 6 )] 2 H 2 O.

Types of Chemical Compounds

HYDRIDES. Metallic and non-metallic.

Neutral oxides, including suboxides.

Basic oxides.

Acidic oxides, including mixed anhydrides.

Saline oxides.

Peroxides divided into poly- and superoxides.

HALIDES. Metallic and non-metallic.

SULPHIDES. Metallic, metalloidal, and non-metallic.

OXYSALTS. Sulphates, carbonates, &c.




A classical illustration of chemical change, at once simple and
valuable, is furnished by the work of Priestley and Lavoisier on
mercuric oxide. Priestley heated mercuric oxide by concentrating
the sun's rays upon it with a lens, so as to decompose it into
mercury and oxygen. The reaction is represented thus:

2HgO 2Hg + O 2 .

This mercuric oxide could previously be obtained, as was shown
by Geber, by gently heating mercury for a long time in the air,
when atmospheric oxygen united with the metal, thus:

2Hg + O 2 2HgO.

Lavoisier combined these two operations by first heating mercury
at a moderate temperature in a confined space, and noting the volume
of air absorbed, and then collecting the mercuric oxide formed and
heating it more strongly; this resulted in the evolution of a volume
of oxygen equal to that of the air which was previously absorbed.
So the possibility of reversing a chemical reaction was established,
a fact now represented thus:

2Hg + O 2 :^ 2HgO.

This simple illustration has been chosen because it gives rise to
various questionings, the consideration of which leads far into the
subject of chemical change in general.

Thus it is a surprising thing that the compound of mercury and
oxygen should be a red, crystalline powder, so different from its
constituent elements, and the question at once occurs whether the
whole of chemistry is full of surprises like that. The elementary
student is rather led to suppose that it is. At least there are many
such surprises which lend to chemical science a fantastic charm for
the youthful mind. For example, the vapour of sulphur is led over



red-hot charcoal, and, instead of yellow crystals and black lumps,
there appears a colourless liquid with a quite extraordinary smell;
or ammonia and hydrogen chloride gases are brought together, and
instead of a neutral gas resulting from the combination of an acid
and an alkaline gas, there is dense white smoke which settles down
as solid sal ammoniac.

" It is the unexpected that happens" might apparently be said of
chemical change. Indeed, the difference between physical mixture
and chemical combination often appears to be this the properties
of a mixture are what might be expected, they are the mean of
those of the constituent parts of the mixture, whilst the properties
of a chemical compound often could not be expected, for they are
unrelated to those of the constituent elements.

Yet, if the impression gained from the facts above considered
were true, and if the above epigram were a generalization of
chemistry then there would be no chemical science; chemistry
would be but a catalogue of curious material phenomena.

The scientific thinker is thus met with the fundamental question
of the relation between the properties of compounds and those of
their constituent elements; he is led, indeed, to the threshold of a
field of inquiry as broad as chemistry itself.

This particular inquiry may, however, be carried a little further.
The greatest differences in properties are seen between elements
and their simplest compounds. When a compound is converted
into another compound by the addition or substitution of other
elements, the physical differences brought about are not so great.
Consider, for example, the series of paraffin hydrocarbons
C n H 2n+2 (see p. 118). An increment of CH 2 causes no surprising
change in the properties of a hydrocarbon; on the contrary it
causes an almost constant alteration of boiling-point and other
physical properties. Or, consider the influence of the substi-
tution of chlorine for hydrogen in the CH 3 group of acetic acid,
CH 3 .COOH. The chloracetic acids CH 2 C1-COOH, CHC1 2 -COOH,
CC1 3 COOH, stand in order of increasing strength; thus the electro-
negative element chlorine has had a specific influence in increasing
the strength of the acid.

The colour, also, of a complex chemical compound is definitely
related to its constitution, and is modified by the substitution of
one element or group for another within the molecule. The art
of producing synthetic dyes depends among other things on a


knowledge of the influence of certain substituents on the colour
of the compound formed. The same is true regarding the thera-
peutic properties of synthetic drugs.

But, more generally, if, according to the periodic law, the
properties of the elements and their compounds are periodic
functions of the atomic weights, then this law should at least
relate the properties of a particular compound to those of a similar
compound of an analogous element. That it does this is shown
by a systematic study of oxides, chlorides, and other simple com-
pounds. Our present knowledge of the law, however, fails to
account completely for the properties of a particular compound.
It is known, for example, that the solid iodides of imperfect metals
are brightly coloured, although the constituent ions are colourless;
such iodides are: PbI 2 , Hg 2 I 2 , HgI 2 , SnI 4 , SbI 8 ; but why PbI 2 is
yellow, for instance, and SnI 4 scarlet, is quite unknown. The same
is true of the colours of some ions; thus there is no theory of the
inherent property of manganese which causes the colour of the
green manganate or crimson permanganate ion.

These are examples of the questionings to which a consideration
of the superficial properties of a simple chemical compound gives
rise. The human mind desires an explanation of the unexpected.
Why is mercuric oxide red; why is permanganate solution crimson;
what hidden properties do the constituent elements of these com-
pounds possess which are revealed in so striking a manner on
combination? And these questions cannot at present be answered.

Eeversible Reactions.

A second question for consideration suggested by the reaction
between mercury and oxygen is that of the reversibility of a
chemical change. Can all chemical reactions be reversed; if not,
why not?

Speaking generally, the possibility of reversing a chemical
reaction depends on the realization of suitable conditions. Some
chemical changes brought about by heat are so profound that
their reversal in the narrower sense is not possible. Sugar, for
example, is destroyed if heated strongly; it is commonly said to
be burnt, and the final result of the burning of sugar in air is the
conversion of its elements into carbon dioxide and water. Can
such a change be reversed?

Growing plants can reconvert the carbon of carbon dioxide



and water into sugar, and the chemist can laboriously synthesize
a kind of sugar; but that is not a true reversal of the chemical
reaction of decomposition, because the synthetic changes do not
follow the same route as the changes of decomposition.

Now, reverting to the reaction between mercury and oxygen,
it would appear from Lavoisier's experiment that there is a certain
minimum temperature at which visible combination between the
elements takes place, and a somewhat higher temperature at which
visible decomposition of the compound formed sets in. These
temperatures cannot be stated because they are conditioned, but it
may be judged that below or above a limited range of temperature
oxygen and mercury do not combine or remain in combination

Such a statement, however, is not very satisfactory, because,
whilst it recognizes the increased activity of mercury and oxygen
molecules, due to rise of temperature, which first promotes com-
bination between the two elements, and subsequently causes
disruption of the compound formed, it takes no account of the
physical state or concentration of the combining elements; or
otherwise, since oxygen is a gas, that it may escape from the
mercury altogether when evolved and so render a reversal of the
reaction impossible.

It is worth while to attempt to gain a clear mental picture
of this reversible reaction, since it illustrates fundamental principles
which underlie chemical reactions in general.

Chemical .Equilibrium.

Suppose the flask A in the figure con-
tains mercury and mercuric oxide in con-
tact with oxygen gas, the pressure of the
latter being indicated by the manometer B,
consisting of a U-tube containing mercury;
and suppose that the flask is heated in a
chamber, the outline of which is shown by
the dotted line, to a temperature t C., within
the limits between which a reaction be-
tween mercury and oxygen is known to

take place. Then the manometer will show in which direction the
reaction is proceeding. If combination is taking place, the mercury
will rise in the limb nearer the flask, owing to diminution of gas

Fig. 45


pressure; if decomposition, the mercury in the nearer limb will
be depressed because of increase of pressure. But in either case
equilibrium will eventually result, and the mercury in the mano-
meter will become and remain stationary, registering a certain gas
pressure; this is not because nothing is taking place, but because
the two opposite reactions are proceeding at equal rates, and a
state of dynamic equilibrium has been attained, which is suitably
represented by the equation:

2Hg + O 2 ^^ 2HgO.

If the temperature is altered the pressure will likewise alter. If,
for instance, the temperature is raised, further decomposition will
begin, but the accumulating oxygen will soon bring it to a stand-
still and a higher constant pressure will be registered corresponding
to the higher temperature.

This reversible chemical reaction, depending upon temperature,
is an example of thermal dissociation^ and the pressure at which
equilibrium is reached at a given temperature is called the dis-
sociation pressure tor that temperature.

The following dissociation pressures have been measured for
the reaction under discussion. 1

Temperature. Millimetres Hg Pressure.

360 C. 90

380 141

400 231

420 387

440 642

460 1017

480 1581

If the temperatures are represented graphically by abscissae, and
the corresponding pressures by ordinates, the relation between
them is given by a curve which takes the form shown in fig. 46.

The curve may be thus interpreted: at the lower temperatures
there is little increase of pressure as the temperature rises, tendency

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 17 19 20 21 22 23

Online LibraryR. M. (Robert Martin) CavenThe foundations of chemical theory; → online text (page 17 of 23)