various agents, yield carbonic oxide; amongst these are many organic
or carboxylic acids. The simplest are formic and oxalic acids. Formic
acid, CH_{2}O_{2}, on being heated to 200°, easily decomposes into
carbonic oxide and water, CH_{2}O_{2} = CO + H_{2}O.[27 bis] Usually,
however, carbonic oxide is prepared in laboratories, not from formic
but from oxalic acid, C_{2}H_{2}O_{4}, the more so as formic acid is
itself prepared from oxalic acid. The latter acid is easily obtained by
the action of nitric acid on starch, sugar, &c.; it is also found in
nature. Oxalic acid is easily decomposed by heat; its crystals first
lose water, then partly volatilise, but the greater part is decomposed.
The decomposition is of the following nature: it splits up into water,
carbonic oxide, and carbonic anhydride,[28] C_{2}H_{2}O_{4} = H_{2}O
+ CO_{2} + CO. This decomposition is generally practically effected
by mixing oxalic acid with strong sulphuric acid, because the latter
assists the decomposition by taking up the water. On heating a mixture
of oxalic and sulphuric acids a mixture of carbonic oxide and carbonic
anhydride is evolved. This mixture is passed through a solution of an
alkali in order to absorb the carbonic anhydride, whilst the carbonic
oxide passes on.[28 bis]
[27] The so-called yellow prussiate, K_{4}FeC_{6}N_{6}, on being heated
with ten parts of strong sulphuric acid forms a considerable
quantity of very pure carbonic oxide quite free from carbonic
anhydride.
[27 bis] To perform this reaction, the formic acid is mixed with
glycerine, because when heated alone it volatilises much below its
temperature of decomposition. When heated with sulphuric acid the
salts of formic acid yield carbonic oxide.
[28] The decomposition of formic and oxalic acids, with the formation
of carbonic oxide, considering these acids as carboxyl
derivatives, may be explained as follows:--The first is H(COOH)
and the second (COOH)_{2}, or H_{2} in which one or both halves of
the hydrogen are exchanged for carboxyl; therefore they are equal
to H_{2} + CO_{2} and H_{2} + 2CO_{2}; but H_{2} reacts with
CO_{2}, as has been stated above, forming CO and H_{2}O. From this
it is also evident that oxalic acid on losing CO_{2} forms formic
acid, and also that the latter may proceed from CO + H_{2}O, as we
shall see further on.
[28 bis] Greshoff (1888) showed that with a solution of nitrate of
silver, iodoform, CHI_{3}, forms CO according to the equation
CHI_{3} + 3AgNO_{3} + H_{2}O = 3AgI + 3HNO_{3} + CO. The reaction
is immediate and is complete.
In its physical _properties_ carbonic oxide resembles nitrogen; this is
explained by the equality of their molecular weights. The absence of
colour and smell, the low temperature of the absolute boiling point,
-140° (nitrogen, -146°), the property of solidifying at -200°
(nitrogen, -202°), the boiling point of -190° (nitrogen, -203°), and the
slight solubility (Chapter I., Note 30), of carbonic oxide are almost the
same as in those of nitrogen. The chemical properties of both gases are,
however, very different, and in these carbonic oxide resembles hydrogen.
Carbonic oxide burns with a blue flame, giving 2 volumes of carbonic
anhydride from 2 volumes of carbonic oxide, just as 2 volumes of hydrogen
give 2 volumes of aqueous vapour. It explodes with oxygen, in the
eudiometer, like hydrogen.[29] When breathed it acts as a strong poison,
being absorbed by the blood;[30] this explains the action of charcoal
fumes, the products of the incomplete combustion of charcoal and other
carbonaceous fuels. Owing to its faculty of combining with oxygen,
carbonic oxide acts as a powerful reducing agent, taking up the oxygen
from many compounds at a red heat, and being itself transformed into
carbonic anhydride. The reducing action of carbonic oxide, however, is
(like that of hydrogen, Chapter II.) naturally confined to those oxides
which easily part with their oxygen--as, for instance, copper
oxide--whilst the oxides of magnesium or potassium are not reduced.
Metallic iron itself is capable of reducing carbonic anhydride to
carbonic oxide, just as it liberates the hydrogen from water. Copper,
which does not decompose water, does not decompose carbonic oxide. If a
platinum wire heated to 300°, or spongy platinum at the ordinary
temperature, be plunged into a mixture of carbonic oxide and oxygen, or
of hydrogen and oxygen, the mixture explodes. These reactions are very
similar to those peculiar to hydrogen. The following important
distinction, however, exists between them--namely: the molecule of
hydrogen is composed of H_{2}, a group of elements divisible into two
like parts, whilst, as the molecule of carbonic oxide, CO, contains
unlike atoms of carbon and oxygen, in none of its reactions of
combination can it give two molecules of matter containing its elements.
This is particularly noticeable in the action of chlorine on hydrogen and
on carbonic oxide respectively; with the former chlorine forms hydrogen
chloride, and with the latter it produces the so-called carbonyl
chloride, COCl_{2}: that is to say, the molecule of hydrogen, H_{2},
under the action of chlorine divides, forming two molecules of
hydrochloric acid, whilst the molecule of carbonic oxide enters in its
entirety into the molecule of carbonyl chloride. This characterises the
so-called _diatomic_ or _bivalent_ reactions of radicles or _residues_.
H is a monatomic residue or radicle, like K, Cl, and others, whilst
carbonic oxide, CO, is an indivisible (undecomposable) bivalent radicle,
equivalent to H_{2} and not to H, and therefore combining with X_{2} and
interchangeable with H_{2}. This distinction is evident from the annexed
comparison:
HH, hydrogen. CO, carbonic oxide.
HCl, hydrochloric acid. COCl_{2}, carbonyl chloride.
HKO, potash. CO(KO)_{2}, potassium carbonate.
HNH_{2}, ammonia. CO(NH_{2})_{2}, urea.
HCH_{3}, methane. CO(CH_{3})_{2}, acetone.
HHO, water. CO(HO)_{2}, carbonic acid.
[29] It is remarkable that, according to the investigations of Dixon,
perfectly dry carbonic oxide does not explode with oxygen when a
spark of low intensity is used, but an explosion takes place if
there is the slightest admixture of moisture. L. Meyer, however,
showed that sparks of an electric discharge of considerable
intensity produce an explosion. N. N. Beketoff demonstrated that
combustion proceeds and spreads slowly unless there be perfect
dryness. I think that this may he explained by the fact that water
with carbonic oxide gives carbonic anhydride and hydrogen, but
hydrogen with oxygen gives hydrogen peroxide (Chapter VII.), which
with carbonic oxide forms carbonic anhydride and water. The water,
therefore, is renewed, and again serves the same purpose. But it
may be that here it is necessary to acknowledge a simple contact
influence. After Dixon had shown the influence of traces of
moisture upon the reaction CO + O, many researches were made of a
similar nature. The fullest investigation into the influence of
moisture upon the course of many chemical reactions was made by
Baker in 1894. He showed that with perfect dryness, many chemical
transformations (for example, the formation of ozone from oxygen,
the decomposition of AgO, KClO_{3} under the action of heat, &c.)
proceeds in exactly the same manner as in the presence of
moisture; but that in many cases traces of moisture have an
evident influence. We may mention the following instances: (1) Dry
SO_{3} does not act upon dry CaO or CuO; (2) perfectly dry
sal-ammoniac does not give NH_{3} with dry CaO, but simply
volatilises; (3) dry NO and O do not react; (4) perfectly dry
NH_{3} and HCl do not combine; (5) perfectly dry sal-ammoniac does
not dissociate at 350° (Chapter VII., Note 15 bis); and (6)
perfectly dry chlorine does not act upon metals, &c.
[30] Carbonic oxide is very rapid in its action, because it is absorbed
by the blood in the same way as oxygen. In addition to this, the
absorption spectrum of the blood changes so that by the help of
blood it is easy to detect the slightest traces of carbonic oxide
in the air. M. A. Kapoustin found that linseed oil and therefore
oil paints, are capable of giving off carbonic oxide while drying
(absorbing oxygen).
Such monatmic (univalent) residues, X, as H, Cl, Na, NO_{2}, NH_{4},
CH_{3}, CO_{2}H (carboxyl), OH, and others, in accordance with
the law of substitution, combine together, forming compounds, XX';
and with oxygen, or in general with diatomic (bivalent) residues,
Y--for instance, O, CO, CH_{2}, S, Ca, &c. forming compounds XX´Y; but
diatomic residues, Y, sometimes capable of existing separately may
combine together, forming YY´ and with X_{2} or XX´, as we see from
the transition of CO into CO_{2} and COCl_{2}. This combining power of
carbonic oxide appears in many of its reactions. Thus it is very easily
absorbed by cuprous chloride, CuCl, dissolved in fuming hydrochloric
acid, forming a crystalline compound, COCu_{2}Cl_{2},2H_{2}O,
decomposable by water; it combines directly with potassium (at 90°),
forming (KCO)_{_n_}[31] with platinum dichloride, PtCl_{2}, with
chlorine, Cl_{2}, &c.
[31] The molecule of metallic potassium (Scott, 1887), like that of
mercury, contains only one atom, and it is probably in virtue of
this that the molecules CO and K combine together. But as in the
majority of cases potassium acts as a univalent radicle, the
polymeride K_{2}C_{2}O_{2} is formed, and probably
K_{10}C_{10}O_{10}, because products containing C_{10} are formed
by the action of hydrochloric acid. The black mass formed by the
combination of carbonic oxide with potassium explodes with great
ease, and oxidises in the air. Although Brodie, Lerch, and Joannis
(who obtained it in 1873 in a colourless form by means of NH_{3}K,
described in Chapter VI., Note 14) have greatly extended our
knowledge of this compound, much still remains unexplained. It
probably exists in various polymeric and isomeric forms, having
the composition (KCO)_{_n_} and (NaCO)_{_n_}.
But the most remarkable compounds are (1) the compound of CO with
metallic nickel, a colourless volatile liquid, Ni(CO)_{4}, obtained by
L. Mond (described in Chapter XXII.) and (2) the compounds of carbonic
oxide with the alkalis, for instance with potassium or barium hydroxide,
&c.--although it is not directly absorbed by them, as it has no acid
properties. Berthelot (1861) showed that potash in the presence of water
is capable of absorbing carbonic oxide, but the absorption takes place
slowly, little by little, and it is only after being heated for many
hours that the whole of the carbonic oxide is absorbed by the potash. The
salt CHKO_{2} is obtained by this absorption; it corresponds with an acid
found in nature--namely, the simplest organic (carboxylic) acid, _formic
acid_, CH_{2}O_{2}. It can be extracted from the potassium salt by means
of distillation with dilute sulphuric acid, just as nitric acid is
prepared from sodium nitrate. The same acid is found in ants and in
nettles (when the stings of the nettles puncture the skin they break, and
the corrosive formic acid enters into the body); it is also obtained
during the action of oxidising agents on many organic substances; it is
formed from oxalic acid, and under many conditions splits up into
carbonic oxide and water. In the formation of formic acid from carbonic
oxide we observe an example of the synthesis of organic compounds, such
as are now very numerous, and are treated of in detail in works on
organic chemistry.
Formic acid, H(CHO_{2}), carbonic acid, HO(CHO_{2}), and oxalic acid,
(CHO_{2})_{2}, are the simple organic or carboxylic acids, R(CHO_{2})
corresponding with HH and HOH. Commencing with carbonic oxide, CO, the
formation of carboxylic acids is clearly seen from the fact that CO is
capable of combining with X_{2}, that is of forming COX_{2}. If, for
instance, one X is an aqueous residue, OH (hydroxyl), and the other X
is hydrogen, then the simplest organic acid--formic acid, H(COOH)--is
obtained. As all hydrocarbons (Chapter VIII.) correspond with the
simplest, CH_{4}, so all organic acids may be considered to proceed from
formic acid.
In a similar way it is easy to explain the relation to other compounds
of carbon of those compounds which contain nitrogen. By way of an
example, we will take one of the carboxyl acids, R(CO_{2}H), where R is
a hydrocarbon radicle (residue). Such an acid, like all others, will
give by combination with NH_{3} an ammoniacal salt, R(CO_{2}NH_{4}).
This salt contains the elements for the formation of two molecules of
water, and under suitable conditions by the action of bodies capable
of taking it up, water may in fact be separated from R(CO_{2}NH_{4}),
forming by the loss of one molecule of water, _amides_, RCONH_{2}, and
by the loss of two molecules of water, _nitriles_, RCN, otherwise known
as _cyanogen compounds_ or _cyanides_.[32] If all the carboxyl acids are
united not only by many common reactions but also by a mutual conversion
into each other (an instance of which we saw above in the conversion
of oxalic acid into formic and carbonic acids) one would expect the
same for all the cyanogen compounds also. The common character of
their reactions, and the reciprocity of their transformation, were
long ago observed by Gay-Lussac, who recognised a common group or
radicle (residue) cyanogen, CN, in all of them. The simplest compounds
are _hydrocyanic_ or _prussic acid_, HCN, cyanic acid, OHCN, and free
cyanogen, (CN)_{2}, which correspond to the three simplest carboxyl
acids: formic, HCO_{2}H, carbonic, OHCO_{2}H, and oxalic, (CO_{2}H)_{2}.
Cyanogen, like carboxyl, is evidently a monatomic residue and acid,
similar to chlorine. As regards the amides RCONH_{2}, corresponding
to the carboxyl acids, they contain the ammoniacal residue NH_{2},
and form a numerous class of organic compounds met with in nature and
obtained in many ways,[33] but not distinguished by such characteristic
peculiarities as the cyanogen compounds.
[32] The connection of the cyanogen compounds with the rest of the
hydrocarbons by means of carboxyl was enunciated by me, about the
year 1860, at the first Annual Meeting of the Russian Naturalists.
[33] Thus, for instance, _oxamide_, or the amide of oxalic acid,
(CNH_{2}O)_{2}, is obtained in the form of an insoluble
precipitate on adding a solution of ammonia to an alcoholic
solution of ethyl oxalate, (CO_{2}C_{2}H_{5})_{2}, which is formed
by the action of oxalic acid on alcohol: (CHO_{2})_{2} +
2(C_{2}H_{5})OH = 2HOH + (CO_{2}C_{2}H_{5})_{2}. As the nearest
derivatives of ammonia, the amides treated with alkalis yield
ammonia and form the salt of the acid. The nitriles do not,
however, give similar reactions so readily. The majority of amides
corresponding to acids have a composition RNH_{2}, and therefore
recombine with water with great ease even when simply boiled with
it, and with still greater facility in presence of acids or
alkalis. Under the action of alkalis the amides naturally give off
ammonia, through the combination of water with the amide, when a
salt of the acid from which the amide was derived is formed:
RNH_{2} + KHO = RKO + NH_{3}.
The same reaction takes place with acids, only an ammoniacal salt
of the acid is of course formed whilst the acid held in the amide
is liberated: RNH_{2} + HCl + H_{2}O = RHO + NH_{4}Cl.
Thus in the majority of cases amides easily pass into ammoniacal
salts, but they differ essentially from them. No ammoniacal salt
sublimes or volatilises unchanged, and generally when heated it
gives off water and yields an amide, whilst many amides volatilise
without alteration and frequently are volatile crystalline
substances which may be easily sublimed. Such, for instance, are
the amides of benzoic, formic, and many other organic acids.
The reactions and properties of the amides and nitriles of the organic
acids are described in detail in books on organic chemistry; we will here
only touch upon the simplest of them, and to clearly explain the
derivative compounds will first consider the ammoniacal salts and amides
of carbonic acid.
As carbonic acid is bibasic, its ammonium salts ought to have the
following composition: _acid carbonate of ammonium_, H(NH_{4})CO_{3},
and _normal carbonate_, (NH_{4})_{2}CO_{3}; they represent compounds
of one or two molecules of ammonia with carbonic acid. The acid salt
appears in the form of a non-odoriferous and (when tested with litmus)
neutral substance, soluble at the ordinary temperature in six parts
of water, insoluble in alcohol, and obtainable in a crystalline form
either without water of crystallisation or with various proportions
of it. If an aqueous solution of ammonia be saturated with an excess
of carbonic anhydride, and then evaporated over sulphuric acid in
the bell jar of an air-pump, crystals of this salt are separated.
Solutions of all other ammonium carbonates, when evaporated under
the air-pump, yield crystals of this salt. A solution of this salt,
even at the ordinary temperature, gives off carbonic anhydride, as do
all the acid salts of carbonic acid (for instance, NaHCO_{3}), and
at 38° the separation of carbonic anhydride takes place with great
rapidity. _On losing carbonic anhydride_ and water, the acid salt is
converted into the normal salt, 2(NH_{4})HCO_{3} = H_{2}O + CO_{2} +
(NH_{4})2CO_{3}; the latter, however, decomposes in solution, and can
therefore only be obtained in crystals, (NH_{4})_{2}CO_{3},H_{2}O, at
low temperatures, and from solutions containing _an excess of ammonia_
as the product of dissociation of this salt: (NH_{4})_{2}CO_{3} = NH_{3}
+ (NH_{4})HCO_{3}. But the normal salt,[34] according to the general
type, is capable of decomposing _with separation of water_, and forming
_ammonium carbamate_, NH_{4}O(CONH_{2}) = (NH_{4})_{2}CO_{3}-H_{2}O;
this still further complicates the chemical transformations of the
carbonates of ammonium. It is in fact evident that, by changing the
ratios of water, ammonia, and carbonic acid, various intermediate salts
will be formed containing mixtures or combinations of those mentioned
above. Thus the ordinary commercial _carbonate of ammonia_ is obtained
by heating a mixture of chalk and sulphate of ammonia (Chapter VI.),
or sal-ammoniac, 2NH_{4}Cl + CaCO_{3} = CaCl_{2} + (NH_{4})_{2}CO_{3}.
The normal salt, however, through loss of part of the ammonia, partly
forms the acid salt, and, partly through loss of water, forms carbamate,
and most frequently presents the composition NH_{4}O(CONH_{2}) +
2OH(CO_{2}NH_{4}) = 4NH_{3} + 3CO_{2} + 2H_{2}O. This salt, in parting
under various conditions with ammonia, carbonic anhydride, and water,
does not present a constant composition, and ought rather to be regarded
as a mixture of acid salt and amide salt. The latter must be recognised
as entering into the composition of the ordinary carbonate of ammonia,
because it contains less water than is required for the normal or acid
salt;[35] but on being dissolved in water this salt gives a mixture of
acid and normal salts.
[34] The acid salt, (NH_{4})HCO_{3}, on losing water ought to form the
_carbamic acid_, OH(CNH_{2}O); but it is not formed, which is
accounted for by the instability of the acid salt itself. Carbonic
anhydride is given off and ammonia is produced, which gives
ammonium carbamate.
[35] In the normal salt, 2NH_{3} + CO_{2} + H_{2}O, in the acid salt,
NH_{3} + CO_{2} + H_{2}O, but in the commercial salt only 2H_{2}O
to 3CO_{2}.
Each of the two ammoniacal salts of carbonic acid has its corresponding
amide. That of the acid salt should be acid, if the water given off takes
up the hydrogen of the ammonia, as it should according to the common type
of formation of the amides, so that OHCONH_{2}, or _carbamic acid_, is
formed from OHCO_{3}NH_{4}. This acid is not known in a free state, but
its corresponding ammoniacal salt or _ammonium carbamate_ is known. The
latter is easily and immediately formed by mixing 2 volumes of _dry_
ammonia with 1 volume of dry carbonic anhydride, 2NH_{3} + CO_{2} =
NH_{4}O(CONH_{2}); it is a solid substance, smells strongly of ammonia,
attracts moisture from the air, and decomposes completely at 60°. The
fact of this decomposition may be proved[36] by the density of its
vapour, which = 13 (H = 1); this exactly corresponds with the density of
a mixture of 2 volumes of ammonia and 1 volume of carbonic anhydride. It
is easily understood that such a combination will take place with any
ammonium carbonate under the action of salts which take up the water--for
instance, sodium or potassium carbonate[37]--as in an anhydrous state
ammonia and carbonic anhydride only form one compound, CO_{2}2NH_{3}.[38]
As the normal ammonium carbonate contains two ammonias, and as the amides
are formed with the separation of water at the expense of the hydrogen of
the ammonias, so this salt has its symmetrical amide, CO(NH_{2})_{2}.
This must be termed carbamide. It is identical with urea, CN_{2}H_{4}O,
which, contained in the urine (about 2 per cent. in human urine), is for
the higher animals (especially the carnivorous) the ordinary product of
excretion[39] and oxidation of the nitrogenous substances found in the
organism. If ammonium carbamate be heated to 140° (in a sealed tube,
Bazaroff), or if carbonyl chloride, COCl_{2}, be treated with ammonia
(Natanson), urea will be obtained, which shows its direct connection with
carbonic acid--that is, the presence of carbonic acid and ammonia in it.
From this it will be understood how urea during the putrefaction of urine
is converted into ammonium carbonate, CN_{2}H_{4}O + H_{2}O = CO_{2} +
2NH_{3}.
[36] Naumann determined the following dissociation tensions of
the vapour of ammonium carbamate (in millimetres of mercury):--
-10° 0° +10° 20° 30° 40° 50° 60°
5 12 30 62 124 248 470 770
Horstmann and Isambert studied the tensions corresponding to
excess of NH_{3} or CO_{2}, and found, as might have been
expected, that with such excess the mass of the salt formed (in a
solid state) increases and the decomposition (transition into
vapour) decreases.
[37] Calcium chloride enters into double decomposition with ammonium
carbamate. Acids (for instance, sulphuric) take up ammonia, and
set free carbonic anhydride; whilst alkalis (such as potash) take
up carbonic anhydride and set free ammonia, and therefore, in this
case for removing water only sodium or potassium carbonate can be
taken. An aqueous solution of ammonium carbamate does not entirely
precipitate a solution of CaCl_{2}, probably because calcium
carbamate is soluble in water, and all the (NH_{3})_{2}CO_{2} is
not converted by dissolving into the normal salt,
(NH_{4}O)_{2}CO_{3}.
[38] It must be imagined that the reaction takes place at first between
equal volumes (Chapter VII.); but then carbamic acid,
HO(CNH_{2}O), is produced, which, as an acid, immediately combines
with the ammonia, forming NH_{4}O(CNH_{2}O).
[39] Urea is undoubtedly a product of the oxidation of complex
nitrogenous matters (albumin) of the animal body. It is found in
the blood. It is absorbed from the blood by the kidneys. A man
excretes about 30 grams of urea per day. As a derivative of
carbonic anhydride, into which it is readily converted, urea is in
a sense a product of oxidation.
Thus urea, both by its origin and decomposition, is an amide of carbonic
acid. Representing as it does ammonia (two molecules) in which hydrogen
(two atoms) is replaced by the bivalent radicle of carbonic acid, urea
retains the property of ammonia of entering into combination, with acids
(thus nitric acid forms CN_{2}H_{4}O,HNO_{3}), with bases (for instance,
with mercury oxide), and with salts (such as sodium chloride, ammonium
chloride), but containing an acid residue it has no alkaline properties.
It is soluble in water without change, but at a red heat loses ammonia
and forms _cyanic acid_, CNHO,[39 bis] which is a nitrile of carbonic
acid--that is to say, is a cyanogen compound, corresponding to the acid
ammonium carbonate, OH(CNH_{4}O_{2}), which on parting with 2H_{2}O ought
to form cyanic acid, CNOH. Liquid cyanic acid, exceedingly unstable at
the ordinary temperatures, gives its stable solid polymer cyanuric acid,
O_{3}H_{3}C_{3}N_{3}. Both have the same composition, and they pass one
into another at different temperatures. If crystals of cyanuric acid be
heated to a temperature, _t_°, then the vapour tension, _p_, in
millimetres of mercury (Troost and Hautefeuille) will be:
_t._ 160°, 170°, 200°, 250°, 300°, 350°
_p._ 56, 68, 130, 220, 430, 1,200
The vapour contains cyanic acid, and, if it be rapidly cooled, it
condenses into a mobile volatile liquid (specific gravity at 0° =
1·14). If the liquid cyanic acid be gradually heated, it passes into
a new amorphous polymeride (cyamelide), which, on being heated, like
cyanuric acid, forms vapours of cyanic acid. If these fumes are heated
above 150° they pass directly into cyanuric acid. Thus at a temperature
of 350°, the pressure does not rise above 1,200 mm. on the addition of
vapours of cyanic acid, because the whole excess is transformed into
cyanuric acid. Hence, the above-mentioned figures give the tension
of dissociation of cyanuric acid, or the greatest pressure which the
vapours of HOCN are able to attain at a given temperature, whilst at
a greater pressure, or by the introduction of a larger mass of the
substance into a given volume, the whole of the excess is converted into
cyanuric acid. The properties of cyanic acid which we have described
were principally observed by Wöhler, and clearly show the _faculty of
polymerisation of cyanogen compounds_. This is observed in many other
cyanogen derivatives, and is to be regarded as the consequence of the
above-mentioned explanation of their nature. All cyanogen compounds are
ammonium salts, R(CNH_{4}O_{2}), deprived of water, 2H_{2}O; therefore
the molecules, RCN, ought to possess the faculty of combining with two
molecules of water or with other molecules in exchange for it (for
instance, with H_{2}S, or HCl, or 2H_{2}, &c.), and are therefore
capable of combining together. The combination of molecules of the same
kind to form more complex ones is what is meant by polymerisation.[40]
[39 bis] Its polymer, C_{3}N_{3}H_{3}O_{3}, is formed together
with it. Cyanic acid is a very unstable, easily changeable liquid,
while cyanuric acid is a crystalline solid which is very stable at
the ordinary temperature.
[40] Just as the aldehydes (such as C_{2}H_{4}O) are alcohols (like
C_{2}H_{6}O) which have lost hydrogen and are also capable of
entering into combination with many substances, and of
polymerising, forming slightly volatile polymerides, which
depolymerise on heating. Although there are also many similar
phenomena (for instance, the transformation of yellow into red
phosphorus, the transition of cinnamene into metacinnamene, &c.)
of polymerisation, in no other case are they so clearly and simply
expressed as in cyanic acid. The details relating to this must be
sought for in treatises on organic and theoretical chemistry. If
we touch on certain sides of this question it is principally with
the view of showing the phenomenon of polymerisation by typical
examples, for it is of more frequent occurrence than was formerly
supposed among compounds of several elements.
Besidea being a substance very prone to form polymerides, cyanic acid
presents many other features of interest, expounded in greater detail in
organic chemistry. However we may mention here the production of the
cyanates by the oxidation of the metallic cyanides. Potassium cyanate,
KCNO, is most often obtained in this way. Solutions of cyanates by the
addition of sulphuric acid yield cyanic acid, which, however, immediately
decomposes: CNHO + H_{2}O = CO_{2} + NH_{3}. A solution of ammonium
cyanate, CN(NH_{4})O, behaves in the same manner, but only in the cold.
On being heated it completely changes because it is transformed into
urea. The composition of both substances is identical, CN_{2}H_{4}O, but
the structure, or disposition of, and connection between, the elements is
different: in the ammonium cyanate one atom of nitrogen exists in the
form of cyanogen, CN--that is, united with carbon--and the other as
ammonium, NH_{4}, but, as cyanic acid contains the hydroxyl radicle of
carbonic acid, OH(CN), the ammonium in this salt is united with oxygen.
The composition of this salt is best expressed by supposing one atom of
the hydrogen in water to be replaced by ammonium and the other by
cyanogen--_i.e._ that its composition is not symmetrical--whilst in urea
both the nitrogen atoms are symmetrically and uniformly disposed as
regards the radicle CO of carbonic acid: CO(NH_{2})_{2}. For this reason,
urea is much more stable than ammonium cyanate, and therefore the latter,
on being slightly heated in solution, is converted into urea. This
remarkable isomeric transformation was discovered by Wöhler in 1828.[41]
Formamide, HCONH_{2}, and _hydrocyanic acid_, HCN, as a nitrile,
correspond with formic acid, HCOOH, and therefore ammonium formate,
HCOONH_{4}, and formamide, when acted on by heat and by substances which
take up water (phosphoric anhydride) form hydrocyanic acid, HCN, whilst,
under many conditions (for instance, on combining with hydrochloric acid
in presence of water), this hydrocyanic acid forms formic acid and
ammonia. Although containing hydrogen in the presence of two acid-forming
elements--namely, carbon and nitrogen[42]--hydrocyanic acid does not
give an acid reaction with litmus (cyanic acid has very marked acid
properties); _but it forms salts_, _MCN_, thus presenting the properties
of a feeble acid, and for this reason is called an _acid_. The small
amount of energy which it has is shown by the fact that the cyanides of
the alkali metals--for instance, potassium cyanide (KHO + HCN = H_{2}0 +
KCN) in solution have a strongly alkaline reaction.[43] If ammonia be
passed over charcoal at a red heat, especially in the presence of an
alkali, or if gaseous nitrogen be passed through a mixture of charcoal
and an alkali (especially potash, KHO), and also if a mixture of
nitrogenous organic substances and alkali be heated to a red heat, in all
these cases the alkali metal combines with the carbon and nitrogen,
forming a metallic cyanide, MCN--for example, KCN.[43 bis] Potassium
cyanide is much used in the arts, and is obtained, as above stated, under
many circumstances--as, for instance, in iron smelting, especially with
the assistance of wood charcoal, the ash of which contains much potash.
The nitrogen of the air, the alkali of the ash, and the charcoal are
brought into contact at a high temperature during iron smelting, and
therefore, under these conditions, a considerable quantity of potassium
cyanide is formed. In practice it is not usual to prepare potassium
cyanide directly, but a peculiar compound of it containing potassium,
iron, and cyanogen. This compound is potassium ferrocyanide, and is also
known as _yellow prussiate of potash_. This saline substance (_see_
Chapter XXII) has the composition K_{4}FeC_{6}N_{6} + 2H_{2}O. The name
of cyanogen ([Greek: kuanos]) is derived from the property which this
yellow prussiate possesses of forming, with a solution of a ferric salt,
FeX_{3}, the familiar pigment Prussian blue. The yellow prussiate is
manufactured on a large scale, and is generally used as the source of
the other cyanogen compounds.
[41] It has an important historical interest, more especially as at
that time such an easy preparation of substances occurring in
organisms without the aid of organic life was quite unexpected,
for they were supposed to be formed under the influence of the
forces acting in organisms, and without the latter their formation
was considered impossible. And in addition to destroying this
illusion, the easy transition of NH_{4}OCN into CO(NH_{2})_{2} is
the best example of the passage of one system of equilibrium of
atoms into another more stable system.
[42] If ammonia and methane (marsh gas) do not show any acid
properties, that is in all probability due to the presence of a
large amount of hydrogen in both; but in hydrocyanic acid one atom
of hydrogen is under the influence of two acid-forming elements.
Acetylene, C_{2}H_{2}, which contains but little hydrogen,
presents acid properties in certain respects, for its hydrogen is
easily replaced by metals. Hydronitrous acid, HN_{3}, which
contains little hydrogen, also has the properties of an acid.
[43] Solutions of cyanides--for instance, those of potassium or
barium--are decomposed by carbonic acid. Even the carbonic
anhydride of the air acts in a similar way, and for this reason
these solutions do not keep, because, in the first place, free
hydrocyanic acid itself decomposes and polymerises, and, in the
second place, with alkaline liquids it forms ammonia and formic
acid. Hydrocyanic acid does not liberate carbonic anhydride from
solutions of sodium or potassium carbonates. But a mixture of
solutions of potassium carbonate and hydrocyanic acid yields
carbonic anhydride on the addition of oxides like zinc oxide,
mercuric oxide, &c. This is due to the great inclination which the
cyanides exhibit of forming double salts. For instance,
ZnK_{2}(CN)_{4} is formed, which is a soluble double salt.
[43 bis] The conversion of the atmospheric nitrogen into cyanogen
compounds, although possible, has not yet been carried out on a
large scale, and one of the problems for future research should be
the discovery of a practical and economical means of converting
the atmospheric nitrogen into metallic cyanides, not only because
potassium cyanide has found a vast and important use for the
extraction of gold from even the poorest ores, but more especially
because the cyanides furnish the means for effecting the synthesis
of many complex carbon compounds, and the nitrogen contained in
cyanogen easily passes into other forms of combination such as
ammonia, which is of great importance in agriculture.
If four parts of yellow prussiate be mixed with eight parts of water and
three parts of sulphuric acid, and the mixture be heated, it decomposes,
volatile hydrocyanic acid separating. This was obtained for the first
time by Scheele in 1782, but it was only known to him in solution. In
1809 Ittner prepared anhydrous prussic acid, and in 1815 Gay-Lussac
finally settled its properties and showed that it contains only hydrogen,
carbon, and nitrogen, CNH. If the distillate (a weak solution of HCN) be
redistilled, and the first part collected, the anhydrous acid may be
prepared from this stronger solution. In order to do this, pieces of
calcium chloride are added to the concentrated solution, when the
anhydrous acid floats as a separate layer, because it is not soluble in
an aqueous solution of calcium chloride. If this layer be then distilled
over a new portion of calcium chloride at the lowest temperature
possible, the prussic acid may be obtained completely free from water. It
is, however, necessary to use the greatest caution in work of this kind,
because prussic acid, besides being extremely poisonous, is exceedingly
volatile.[44]
[44] The mixture of the vapours of water and hydrocyanic acid, evolved
on heating yellow prussiate with sulphuric acid, may be passed
directly through vessels or tubes filled with calcium chloride.
These tubes must be cooled, because, in the first place,
hydrocyanic acid easily changes on being heated, and, in the
second place, the calcium chloride when warm would absorb less
water. The mixture of hydrocyanic acid and aqueous vapour on
passing over a long layer of calcium chloride gives up water, and
hydrocyanic acid alone remains in the vapour. It ought to be
cooled as carefully as possible in order to bring it into a liquid
condition. The method which Gay-Lussac employed for obtaining pure
hydrocyanic acid consisted in the action of hydrochloric acid gas
on mercuric cyanide. The latter may he obtained in a pure state if
a solution of yellow prussiate be boiled with a solution of
mercuric nitrate, filtered, and crystallised by cooling; the
mercuric cyanide is then obtained in the form of colourless
crystals, Hg(CN)_{2}.
If a strong solution of hydrochloric acid be poured upon these
crystals, and the mixture of vapours evolved, consisting of
aqueous vapour, hydrochloric acid, and hydrocyanic acid, be passed
through a tube containing, first, marble (for absorbing the
hydrochloric acid), and then lumps of calcium chloride, on cooling
the hydrocyanic acid will be condensed. In order to obtain the
latter in an anhydrous form, the decomposition of heated mercury
cyanide by hydrogen sulphide may be made use of. Here the sulphur
and cyanogen change places, and hydrocyanic acid and mercury
sulphide are formed: Hg(CN)_{2} + H_{2}S = 2HCN + HgS.
Anhydrous prussic acid is a very mobile and volatile liquid; its specific
gravity is 0·697 at 18°; at lower temperatures, especially when mixed
with a small quantity of water, it easily congeals; it boils at 26°, and
therefore very easily evaporates, and at ordinary temperatures may be
regarded as a gas. An insignificant amount, when inhaled or brought into
contact with the skin, causes death. It is soluble in all proportions in
water, alcohol, and ether: weak aqueous solutions are used in
medicine.[45]
[45] A weak (up to 2 p.c.) aqueous solution of hydrocyanic acid is
obtained by the distillation of certain vegetable substances. The
so-called laurel water in particular enjoys considerable notoriety
from its containing hydrocyanic acid. It is obtained by the
steeping and distillation of laurel leaves. A similar kind of
water is formed by the infusion and distillation of bitter
almonds. It is well known that bitter almonds are poisonous, and
have a peculiar characteristic taste. This bitter taste is due to
the presence of a certain substance called amygdalin, which can be
extracted by alcohol. This amygdalin decomposes in an infusion of
bruised almonds, forming the so-called bitter almond oil, glucose,
and hydrocyanic acid:
C_{10}H_{27}NO_{11} + H_{2}O = C_{7}H_{6}O + CNH + 2C_{6}H_{12}O_{6}
Amygdalin in Water Bitter Hydrocyanic Glucose
bitter almonds almond acid
oil
If after this the infusion of bitter almonds be distilled with
water, the hydrocyanic acid and the volatile bitter almond oil are
carried over with the aqueous vapour. The oil is insoluble in
water, or only sparingly soluble, while the hydrocyanic acid
remains as an aqueous solution. Bitter almond water is similar to
laurel water, and is used like the former in medicine, naturally
only in small quantities because any considerable amount has
poisonous effects. Perfectly pure anhydrous hydrocyanic acid keeps
without change, just like the weak solutions, but the strong
solutions only keep in the presence of other acids. In the
presence of many admixtures these solutions easily give a brown
polymeric substance, which is also formed in a solution of
potassium cyanide.
The salts MCN--for instance, potassium, sodium,
ammonium--as well as the salts M´´(CN)_{2}--for example, barium,
calcium, mercury--are soluble in water, but the cyanides of manganese,
zinc, lead, and many others are insoluble in water. They form double
salts with potassium cyanide and similar metallic cyanides, an example
of which we will consider in a further description of the yellow
prussiate. Not only are some of the double salts remarkable for their
constancy and comparative stability, but so also are the soluble salt
HgC_{2}N_{2}, the insoluble silver cyanide AgCN, and even potassium
cyanide in the absence of water. The last salt,[46] when fused, acts
as a reducing agent with its elements K and C, and oxidises when fused
with lead oxide, forming potassium cyanate, KOCN, which establishes the
connection between HCN and OHCN--that is, between the nitriles of formic
and carbonic acids--and this connection is the same as that between
the acids themselves, since formic acid, on oxidation, yields carbonic
acid. Free cyanogen, (CN)_{2} or CNCN, corresponds to hydrocyanic acid
in the same manner as free chlorine, Cl_{2} or ClCl, corresponds to
hydrochloric acid. This composition, judging from what has been already
stated, exactly expresses that of the nitrile of oxalic acid, and, as
a matter of fact, oxalate of ammonia and the amide corresponding with
it (oxamide, Note 33), on being heated with phosphoric anhydride, which
takes up the water, yield _cyanogen_, (CN)_{2}. This substance is also
produced by simply heating some of the metallic cyanides. Mercuric
cyanide is particularly adapted for this purpose, because it is easily
obtained in a pure state and is then very stable. If mercuric cyanide be
heated, it decomposes, in like manner to mercury oxide, into metallic
mercury and cyanogen: HgC_{2}N_{2} = Hg + C_{2}N_{2}.[47] When cyanogen
is formed, part of it always polymerises into a dark brown insoluble
substance called _paracyanogen_, capable of forming cyanogen when
heated to redness.[48] Cyanogen is a colourless, poisonous gas, with a
peculiar smell and easily condensed by cooling into a colourless liquid,
insoluble in water and having a specific gravity of 0·86. It boils at
about -21°, and therefore cyanogen may be easily condensed into a liquid
by a strong freezing mixture. At -35° liquid cyanogen solidifies. The gas
is soluble in water and in alcohol to a considerable extent--namely,
1 volume of water absorbs as much as 4-1/2 volumes, and alcohol 23
volumes. Cyanogen resists the action of a tolerably high temperature
without decomposing, but under the action of the electric spark the
carbon is separated, leaving a volume of nitrogen equal to the volume
of the gas taken. As it contains carbon it burns, and the colour of the
flame is reddish-violet, which is due to the presence of nitrogen, all
compounds of which impart more or less of this reddish-violet hue to
the flame. During the combustion of cyanogen, carbonic anhydride and
nitrogen are formed. The same products are obtained in the eudiometer
with oxygen or by the action of cyanogen on many oxides at a red heat.
[46] This salt will be described in Chapter XIII.
[47] For the preparation it is necessary to take completely dry mercuric
cyanide, because when heated in the presence of moisture it gives
ammonia, carbonic anhydride, and hydrocyanic acid. Instead of
mercuric cyanide, a mixture of perfectly dry yellow prussiate and
mercuric chloride may be used, then double decomposition and the
formation of mercuric cyanide take place in the retort. Silver
cyanide also disengages cyanogen, on being heated.
[47] _Paracyanogen_ is a brown substance (having the composition of
cyanogen) which is formed during the preparation of cyanogen by
all methods, and remains as a residue. Silver cyanide, on being
slightly heated, fuses, and on being further heated evolves a gas;
a considerable quantity of paracyanogen remains in the residue.
Here it is remarkable that exactly half the cyanogen becomes
gaseous, and the other half is transformed into paracyanogen.
Metallic silver will be found in the residue with the
paracyanogen; it may be extracted with mercury or nitric acid,
which does not act on paracyanogen. If paracyanogen be heated in a
vacuum it decomposes, forming cyanogen; but here the pressure _p_
for a given temperature _t_ cannot exceed a certain limit, so that
the phenomenon presents all the external appearance of a physical
transformation into vapour; but, nevertheless, it is a complete
change in the nature of the substance, though limited by the
_pressure of dissociation_, as we saw before in the transformation
of cyanuric into hydrocyanic acid, and as would be expected from
the fundamental principles of dissociation. Troost and
Hautefeuille (1868) found that for paracyanogen,
_t_ = 530° 581° 600° 635°
_p_ = 90 143 296 1,089 mm.
However, even at 550° part of the cyanogen decomposes into carbon
and nitrogen. The reverse transition of cyanogen into paracyanogen
commences at 350°, and at 600° proceeds rapidly. And if the
transition of the first kind is likened to evaporation, then the
reverse transition, or polymerisation, presents a likeness to the
transition of vapours into the solid state.
The relation of cyanogen to the metallic cyanides is seen not only in the
fact that it is formed from mercuric cyanide, but also by its forming
cyanide of sodium or potassium on being heated with either of those
metals, the sodium or potassium taking fire in the cyanogen. On heating a
mixture of hydrogen and cyanogen to 500° (Berthelot),[49] or under the
action of the silent discharge (Boilleau), hydrocyanic acid is formed, so
that the reciprocity of the transitions does not leave any doubt in the
matter that all the nitriles of the organic acids contain cyanogen, just
as all the organic acids contain carboxyl and in it the elements of
carbonic anhydride. Besides the amides,[50] the nitriles (or cyanogen
compounds, RCN), and nitro-compounds (containing the radicle of nitric
acid, RNO_{2}), there are a great number of other substances containing
at the same time carbon and nitrogen, particulars of which must be sought
for in special works on organic chemistry.
[49] Cyanogen (like chlorine) is absorbed by a solution of sodium
hydroxide, sodium cyanide and cyanate being produced: C_{2}N_{2} +
2NaHO = NaCN + CNNaO + H_{2}O. But the latter salt decomposes
relatively easily, and moreover part of the cyanogen liberated by
heat from its compounds undergoes a more complex transformation.
[50] If, in general, compounds containing the radicle NH_{2} are called
amides, some of the _amines_ ought to be ranked with them; namely,
the hydrocarbons C_{_n_}H_{2_m_}, in which part of the hydrogen is
replaced by NH_{2}; for instance, methylamine, CH_{3}NH_{2},
aniline, C_{6}H_{5}NH_{2}, &c. In general the amines may be
represented as ammonia in which part or all of the hydrogen is
replaced by hydrocarbon radicles--as, for example, trimethylamine,
N(CH_{3})_{3}. They, like ammonia, combine with acids and form
crystalline salts. Analogous substances are sometimes met with in
nature, and bear the general name of _alkaloids_; such are, for
instance, quinine in cinchona bark, nicotine in tobacco, &c.
CHAPTER X
SODIUM CHLORIDE--BERTHOLLET'S LAWS--HYDROCHLORIC ACID
In the preceding chapters we have become acquainted with the most
important properties of the four elements, hydrogen, oxygen, nitrogen,
and carbon. They are sometimes termed the _organogens_, because
they enter into the composition of organic substances. Their mutual
combinations may serve as types for all other chemical compounds--that
is, they present the same atomic relations (types, forms, or grades of
combinations) as those in which the other elements also combine together.
Hydrogen, HH, or, in general, HR.
Water, H_{2}O, " " H_{2}R.
Ammonia, H_{3}N, " " H_{3}R.
Marsh gas, H_{4}C, " " H_{4}R.
One, two, three, and four atoms of hydrogen enter into these molecules
for one atom of another element. No compounds of one atom of oxygen with
three or four atoms of hydrogen are known; hence the atom of oxygen does
not possess certain properties which are found in the atoms of carbon
and nitrogen.
The faculty of an element to form a compound of definite composition
with hydrogen (or an element analogous to it) gives the possibility
of foretelling the composition of many other of its compounds. Thus,
if we know that an element, M, combines with hydrogen, forming, by
preference, a gaseous substance such as HM, but not forming H_{2}M,
H_{3}M, H_{n}M_{m}, then we must conclude, on the basis of the law of
substitution, that this element will give compounds M_{2}O, M_{3}N,
MHO, MH_{3}C, &c. Chlorine is an example of this kind. If we know
that another element, R, like oxygen, gives with hydrogen a molecule
H_{2}R, then we may expect that it will form compounds similar to
hydrogen peroxide, the metallic oxides, carbonic anhydride, or carbonic
oxide, and others. Sulphur is an instance of this kind. Hence the
elements may be classified according to their resemblance to hydrogen,
oxygen, nitrogen, and carbon, and in conformity with this analogy it
is possible to foretell, if not the properties (for example, the
acidity or basicity), at any rate the composition,[1] of some of their
compounds. This forms the substance of _the conception of the valency
or atomicity of the elements_. Hydrogen is taken as the representative
of the univalent elements, giving compounds, RH, R(OH), R_{2}O, RCl,
R_{3}N, R_{4}C, &c. Oxygen, in that form in which it gives water, is the
representative of the bivalent elements, forming RH_{2}, RO, RCl_{2},
RHCl, R(OH)Cl, R(OH)_{2}, R_{2}C, RCN, &c. Nitrogen in ammonia is the
representative of the trivalent elements, giving compounds RH_{3},
R_{2}O_{3}, R(OH)_{3}, RCl_{3}, RN, RHC, &c. In carbon are exemplified
the properties of the quadrivalent elements, forming RH_{4}, RO_{2},
RO(OH)_{2}, R(OH)_{4}, RHN, RCl_{4}, RHCl_{3}, &c. We meet with these
_forms of combination_, or degrees of union of atoms, in all other
elements, some being analogous to hydrogen, others to oxygen, and others
to nitrogen or to carbon. But besides these quantitative analogies or
resemblances, which are foretold by the law of substitution (Chapter
VI.), there exist among the elements qualitative analogies and relations
which are not fully seen in the compounds of the elements which have
been considered, but are most distinctly exhibited in the formation of
bases, acids, and salts of different types and properties. Therefore,
for a complete study of the nature of the elements and their compounds
it is especially important to become acquainted with the salts, as
substances of a peculiar character, and with the corresponding acids and
bases. Common table salt, or sodium chloride, NaCl, may in every respect
be taken as a type of salts in general, and we will therefore pass to
the consideration of this substance, and of hydrochloric acid, and of
the base sodium hydroxide, formed by the non-metal chlorine and the
metal sodium, which correspond with it.
[1] But it is impossible to foretell all the compounds formed by an
element from its atomicity or valency, because the atomicity of the
elements is variable, and furthermore this variability is not
identical for different elements. In CO_{2}, COX_{2}, CH_{4}, and
the multitude of carbon compounds corresponding with them, the C is
quadrivalent, but in CO either the carbon must be taken as bivalent
or the atomicity of oxygen be accounted as variable. Moreover,
carbon is an example of an element which preserves its atomicity to
a greater degree than most of the other elements. Nitrogen in
NH_{3}, NH_{2}(OH), N_{2}O_{3}, and even in CNH, must be considered
as trivalent, but in NH_{4}Cl, NO_{2}(OH), and in all their
corresponding compounds it is necessarily pentavalent. In N_{2}O,
if the atomicity of oxygen = 2, nitrogen has an uneven atomicity
(1, 3, 5), whilst in NO it is bivalent. If sulphur be bivalent,
like oxygen, in many of its compounds (for example, H_{2}S,
SCl_{2}, KHS, &c.), then it could not be foreseen from this that it
would form SO_{2}, SO_{3}, SCl_{4}, SOCl_{2}, and a series of
similar compounds in which its atomicity must be acknowledged as
greater than 2. Thus SO_{2}, sulphurous anhydride, has many points
in common with CO_{2}, and if carbon be quadrivalent then the S in
SO_{2} is quadrivalent. Therefore the principle of atomicity
(valency) of the elements cannot be considered established as the
basis for the study of the elements, although it gives an easy
method of grasping many analogies. I consider the four following as
the chief obstacles to acknowledging the atomicity of the elements
as a primary conception for the consideration of the properties of
the elements: 1. Such univalent elements as H, Cl, &c., appear in a
free state as molecules H_{2}, Cl_{2}, &c., and are consequently
like the univalent radicles CH_{3}, OH, CO_{2}H, &c., which, as
might be expected, appear as C_{2}H_{6}, O_{2}H_{2},
C_{2}O_{4}H_{2} (ethane, hydrogen peroxide, oxalic acid), whilst on
the other hand, potassium and sodium (perhaps also iodine at a high
temperature) contain only one atom, K, Na, in the molecule in a
free state. Hence it follows that _free affinities_ may exist.
Granting this, nothing prevents the assumption that free affinities
exist in all unsaturated compounds; for example, two free
affinities in NH_{3}. If such instances of free affinities be
admitted, then all the possible advantages to be gained by the
application of the doctrine of atomicity (valency) are lost. 2.
There are instances--for example, Na_{2}H--where univalent elements
are combined in molecules which are more complex than R_{2}, and
form molecules, R_{3}, R_{4}, &c.; this may again be either taken
as evidence of the existence of free affinities, or else
necessitates such primary univalent elements as sodium and hydrogen
being considered as variable in their atomicity. 3. The periodic
system of the elements, with which we shall afterwards become
acquainted, shows that there is a law or rule for the variation of
the forms of oxygen and hydrogen compounds; chlorine is univalent
with respect to hydrogen, and septavalent with respect to oxygen;
sulphur is bivalent to hydrogen, and sexavalent to oxygen;
phosphorus is trivalent to hydrogen and pentavalent in respect to
oxygen--the sum is in every case equal to 8. Only carbon and its
analogues (for example, silicon) are quadrivalent to both hydrogen
and oxygen. Hence the power of the elements to change their
atomicity is an essential part of their nature, and therefore
constant valency cannot he considered as a fundamental property. 4.
Crystallo-hydrates (for instance, NaCl,2H_{2}O, or NaBr,2H_{2}O),
double salts (such as PtCl_{4},2KCl,H_{2}SiF_{6}, &c.), and similar
complex compounds (and, according to Chap. I., solutions also)
demonstrate the capacity not only of the elements themselves, but
also of their saturated and limiting compounds, of entering into
further combination. Therefore the admission of a definite limited
atomicity of the elements includes in itself an admission of
limitation which is not in accordance with the nature of chemical
reactions.
_Sodium chloride_, NaCl, the familiar table salt, occurs, although in
very small quantities, in all the primary formations of the earth's
crust,[2] from which it is washed away by the atmospheric waters; it is
contained in small quantities in all waters flowing through these
formations, and is in this manner conveyed to the oceans and seas. The
immense mass of salt in the oceans has been accumulated by this process
from the remote ages of the earth's creation, because the water has
evaporated from them while the salt has remained in solution. The salt of
sea water serves as the source not only for its direct extraction, but
also for the formation of other masses of workable salt, such as rock
salt, and of saline springs and lakes.[2 bis]
[2] The primary formations are those which do not bear any distinct
traces of having been deposited from water (have not a stratified
formation and contain no remains of animal or vegetable life),
occur under the sedimentary formations of the earth, and are
everywhere uniform in composition and structure, the latter being
generally distinctly crystalline. If it be assumed that the earth
was originally in a molten condition, the first primary formations
are those which formed the first solid crust of the earth. But even
with this hypothesis of the earth's origin, it is necessary to
admit that the first aqueous deposits must have caused a change in
the original crust of the earth, and therefore under the head of
primary formations must be understood the most ancient of the
products of decomposition (mostly by atmospheric, aqueous, and
organic agency, &c.), from which all the rocks and substances of
the earth's surface have arisen. In speaking of the origin of one
or another substance, we can only, on the basis of facts, descend
to the primary formations, of which granite, gneiss, and trachyte
may be taken as examples.
[2 bis] Chloride of sodium has been found to occur in the atmosphere in
the form of a fine dust; in the lower strata it is present in
larger quantities than in the upper, so that the rain water falling
on mountains contains less NaCl than that falling in valleys. Müntz
(1891) found that a litre of rain water collected on the summit of
the Pic du Midi (2,877 metres above the sea level) contained 0·34
milligram of chloride of sodium, while a litre of rain collected
from the valley contained 2·5-7·6 milligrams.
The extraction of salt _from sea water_ is carried on in several ways. In
southern climes, especially on the shores of the Atlantic Ocean and the
Mediterranean and Black Seas, the summer heats are taken advantage of. A
convenient low-lying sea shore is chosen, and a whole series of basins,
communicating with each other, are constructed along it. The upper of
these basins are filled with sea water by pumping, or else advantage is
taken of high tides. These basins are sometimes separated from the sea by
natural sand-banks (limans) or by artificial means, and in spring the
water already begins to evaporate considerably. As the solution becomes
more concentrated, it is run into the succeeding basins, and the upper
ones are supplied with a fresh quantity of sea water, or else an
arrangement is made enabling the salt water to flow by degrees through
the series of basins. It is evident that the beds of the basins should be
as far as possible impervious to water, and for this purpose they are
made of beaten clay. The crystals of salt begin to separate out when the
concentration attains 28 p.c. of salt (which corresponds to 28° of
Baumé's hydrometer). They are raked off, and employed for all those
purposes to which table salt is applicable. In the majority of cases only
the first half of the sodium chloride which can be separated from the sea
water is extracted, because the second half has a bitter taste from the
presence of magnesium salts which separate out together with the sodium
salt. But in certain localities--as, for instance, in the estuary of the
Rhone, on the island of Camarga[3]--the evaporation is carried on to the
very end, in order to obtain those magnesium and potassium salts which
separate out at the end of the evaporation of sea water. Various salts
are separated from sea water in its evaporation. From 100 parts of sea
water there separates out, by natural and artificial evaporation, about
one part of tolerably pure table salt at the very commencement of the
operation; the total amount held in solution being about 2-1/2 p.c. The
remaining portion separates out intermixed with the bitter salts of
magnesium which, owing to their solubility and the small amount in which
they are present (less than 1 p.c.), only separate out, in the first
crystallisations, in traces. Gypsum, or calcium sulphate,
CaSO_{4},2H_{2}O, because of its sparing solubility, separates together
with or even before the table salt. When about half of the latter has
separated, then a mixture of table salt and magnesium sulphate separates
out, and on still further evaporation the chlorides of potassium and
magnesium begin to separate in a state of combination, forming the double
salt KMgCl_{3},6H_{2}O, which occurs in nature as _carnallite_.[4] After
the separation of this salt from sea water, there remains a mother liquor
containing a large amount of magnesium chloride in admixture with various
other salts.[5] The extraction of sea salt is usually carried on for the
purpose of procuring table salt, and therefore directly it begins to
separate mixed with a considerable proportion[6] of magnesium salts (when
it acquires a bitter taste) the remaining liquor is run back into the
sea.
[3] The extraction of the potassium salts (or so-called summer salts)
was carried on at the Isle of Camarga about 1870, when I had
occasion to visit that spot. At the present time the deposits of
Stassfurt provide a much cheaper salt, owing to the evaporation and
separation of the salt being carried on there by natural means and
only requiring a treatment and refining, which is also necessary in
addition for the 'summer salt' obtained from sea-water.
[4] The double salt KCl,MgCl_{2} is a crystallohydrate of KCl and
MgCl_{2}, and is only formed from solutions containing an excess of
magnesium chloride, because water decomposes this double salt,
extracting the more soluble magnesium chloride from it.
[5] Owing to the fundamental property of salts of interchanging their
metals, it cannot be said that sea water contains this or that
salt, but only that it contains certain amounts of certain metals M
(univalent like Na and K, and bivalent like Mg and Ca), and haloids
X (univalent like Cl, Br, and bivalent like SO_{4}, CO_{3}), which
are disposed in every possible kind of grouping; for instance, K as
KCl, KBr, K_{2}SO_{4}, Mg as MgCl_{2}, MgBr_{2}, MgSO_{4}, and so
on for all the other metals. In evaporation different salts
separate out consecutively only because they reach saturation. A
proof of this may be seen in the fact that a solution of a mixture
of sodium chloride and magnesium sulphate (both of which salts are
obtained from sea water, as was mentioned above), when evaporated,
deposits crystals of these salts, but when refrigerated (if the
solution be sufficiently saturated) the salt Na_{2}SO_{4},10H_{2}O
is first deposited because it is the first to arrive at saturation
at low temperatures. Consequently this solution contains MgCl_{2}
and Na_{2}SO_{4}, besides MgSO_{4} and NaCl. So it is with sea
water.
[6] The salt extracted from water is piled up in heaps and left exposed
to the action of rain water, which purifies it, owing to the water
becoming saturated with sodium chloride and then no longer
dissolving it, but washing out the impurities.
The same process which is employed for artificially obtaining salt in a
crystalline form from sea water has been repeatedly accomplished during
the geological evolution of the earth on a gigantic scale; upheavals of
the earth have cut off portions of the sea from the remainder (as the
Dead Sea was formerly a part of the Mediterranean, and the Sea of Aral of
the Caspian), and their water has evaporated and formed (if the mass of
the inflowing fresh water were less than that of the mass evaporated)
deposits of _rock salt_. It is always accompanied by gypsum, because the
latter is separated from sea water with or before the sodium chloride.
For this reason rock salt may always be looked for in those localities
where there are deposits of gypsum. But inasmuch as the gypsum remains on
the spot where it has been deposited (as it is a sparingly soluble salt),
whilst the rock salt (as one which is very soluble) may be washed away by
rain or fresh running water, it may sometimes happen that although gypsum
is still found there may be no salt; but, on the other hand, where there
is rock salt there will always be gypsum. As the geological changes of
the earth's surface are still proceeding at the present day, so in the
midst of the dry land salt lakes are met with, which are sometimes
scattered over vast districts formerly covered by seas now dried up. Such
is the origin of many of the salt lakes about the lower portions of the
Volga and in the Kirghiz steppes, where at a geological epoch preceding
the present the Aralo-Caspian Sea extended. Such are the Baskunchaksky
(in the Government of Astrakhan, 112 square kilometres superficial area),
the Eltonsky (140 versts from the left bank of the Volga, and 200 square
kilometres in superficial area), and upward of 700 other salt lakes lying
about the lower portions of the Volga. In those in which the inflow of
fresh water is less than that yearly evaporated, and in which the
concentration of the solution has reached saturation, the
_self-deposited_ salt is found already deposited on their beds, or is
being yearly deposited during the summer months. Certain limans, or
sea-side lakes, of the Azoff Sea are essentially of the same
character--as, for instance, those in the neighbourhood of Henichesk and
Berdiansk. The saline soils of certain Central Asian steppes, which
suffer from a want of atmospheric fresh water, are of the same origin.
Their salt originally proceeded from the salt of seas which previously
covered these localities, and has not yet been washed away by fresh
water. The main result of the above-described process of nature is the
formation of masses of rock salt, which are, however, being gradually
washed away by the subsoil waters flowing in their neighbourhood, and
afterwards rising to the surface in certain places as _saline springs_,
which indicate the presence of masses of deposited rock salt in the
depths of the earth. If the subsoil water flows along a stratum of salt
for a sufficient length of time it becomes saturated; but in flowing in
its further course along an impervious stratum (clay) it becomes diluted
by the fresh water leaking through the upper soil, and therefore the
greater the distance of a saline spring from the deposit of rock salt,
the poorer will it be in salt. A perfectly saturated brine, however, may
be procured from the depths of the earth by means of bore-holes. The
deposits of rock salt themselves, which are sometimes hidden at great
depths below the earth's strata, may be discovered by the guidance of
bore-holes and the direction of the strata of the district. Deposits of
rock salt, about 35 metres thick and 20 metres below the surface, were
discovered in this manner in the neighbourhood of Brianstcheffky and
Dekonoffky, in the Bakhmut district of the Government of Ekaterinoslav.
Large quantities of most excellent rock salt are now (since 1880)
obtained from these deposits, whose presence was indicated by the
neighbouring salt springs (near Slaviansk and Bakhmut) and by bore-holes
which had been sunk in these localities for procuring strong (saturated)
brines. But the Stassfurt deposits of rock salt near Magdeburg in Germany
are celebrated as being the first discovered in this manner, and for
their many remarkable peculiarities.[7] The plentiful distribution of
saline springs in this and the neighbouring districts suggested the
presence of deposits of rock salt in the vicinity. Deep bore-holes sunk
in this locality did in fact give a richer brine--even quite saturated
with salt. On sinking to a still greater depth, the deposits of salt
themselves were at last arrived at. But the first deposit which was met
with consisted of a bitter salt unfit for consumption, and was therefore
called refuse salt (_Abraumsalz_). On sinking still deeper vast beds of
real rock salt were struck. In this instance the presence of these upper
strata containing salts of potassium, magnesium, and sodium is an
excellent proof of the formation of rock salt from sea water. It is very
evident that not only a case of evaporation to the end--as far, for
instance, as the separation of carnallite--but also the preservation of
such soluble salts as separate out from sea water after the sodium
chloride, must be a very exceptional phenomenon, which is not repeated in
all deposits of rock salt. The Stassfurt deposits therefore are of
particular interest, not only from a scientific point of view, but also
because they form a rich source of potassium salts which have many
practical uses.[7 bis]
[7] When the German savants pointed out the exact locality of the
Stassfurt salt-beds and their depth below the surface, on the basis
of information collected from various quarters respecting
bore-holes and the direction of the strata, and when the borings,
conducted by the Government, struck a salt-bed which was bitter and
unfit for use, there was a great outcry against science, and the
doubtful result even caused the cessation of the further work of
deepening the shafts. It required a great effort to persuade the
Government to continue the work. Now, when the pure salt
encountered below forms one of the important riches of Germany, and
when those 'refuse salts' have proved to be most valuable (as a
source of potassium and magnesium), we should see in the
utilisation of the Stassfurt deposits one of the conquests of
science for the common welfare.
[7 bis] In Western Europe, deposits of rock salt have long been known
at Wieliczka, near Cracow, and at Cardona in Spain. In Russia the
following deposits are known: (_a_) the vast masses of rock salt (3
square kilometres area and up to 140 metres thick) lying directly
on the surface of the earth at Iletzky Zastchit, on the left bank
of the river Ural, in the Government of Orenburg; (_b_) the
Chingaksky deposit, 90 versts from the river Volga, in the
Enotaeffsky district of the Government of Astrakhan; (_c_) the
Kulepinsky (and other) deposits (whose thickness attains 150
metres), on the Araks, in the Government of Erivan in the Caucasus;
(_d_) the Katchiezmansky deposit in the province of Kars; (_e_) the
Krasnovodsky deposit in the Trans-Caspian province; and (_f_) the
Bardymkulsky salt mines in Kokhand.
A saturated brine, formed by the continued contact of subsoil water with
rock salt, is extracted by means of bore-holes, as, for instance, in the
Governments of Perm, Kharkoff, and Ekaterinoslav. Sometimes, as at
Berchtesgaden (and at Hallein) in Austria, spring water is run on to
underground beds of rock salt containing much clay.
[Illustration: FIG. 64.--Graduator for the evaporation of the water of
saline springs.]
If a saline spring, or the salt water pumped from bore-holes, contains
but little salt, then the first concentration of the natural solution
is not carried on by the costly consumption of fuel, but by the cheaper
method of evaporation by means of the wind. For this purpose so-called
graduators are constructed: they consist of long and lofty sheds,
which are sometimes several versts long, and generally extend in a
direction at right angles to that of the usual course of the wind in
the district. These sheds are open at the sides, and are filled with
brushwood as shown in fig. 64. Troughs, A B, C D, into which the salt
water is pumped, run along the top. On flowing from these troughs,
through the openings, _a_, the water spreads over the brushwood and
distributes itself in a thin layer over it, so that it presents a very
large surface for evaporation, in consequence of which it rapidly
becomes concentrated in warm or windy weather. After trickling over the
brushwood, the solution collects in a reservoir under the graduator,
whence it is usually pumped up by the pumps P P´, and again run a
second and third time through the graduator, until the solution reaches
a degree of concentration at which it becomes profitable to extract
the salt by direct heating. Generally the evaporation in the graduator
is not carried beyond a concentration of 12 to 15 parts of salt in
100 parts of solution. Strong natural solutions of salt, and also the
graduated solutions, are evaporated in large shallow metallic vessels,
which are either heated by the direct action of the flame from below
or from above. These vessels are made of boiler plate, and are called
salt-pans. Various means are employed for accelerating the evaporation
and for economising fuel, which are mainly based on an artificial
draught to carry off the steam as it is formed, and on subjecting the
saline solution to a preliminary heating by the waste heat of the
steam and furnace gases. Furthermore, the first portions of the salt
which crystallise out in the salt-pans are invariably contaminated
with gypsum, since the waters of saline springs always contain this
substance. It is only the portions of the salt which separate later that
are distinguished by their great purity. The salt is ladled out as it is
deposited, left to drain on inclined tables and then dried, and in this
manner the so-called bay salt is obtained. Since it has become possible
to discover the saline deposits themselves, the extraction of table salt
from the water of saline springs by evaporation, which previously was
in general use, has begun to be disused, and is only able to hold its
ground in places where fuel is cheap.
In order to understand the full importance of the extraction of salt,
it need only be mentioned that on the average 20 lbs. of table salt are
consumed yearly per head of population, directly in food or for cattle.
In those countries where common salt is employed in technical processes,
and especially in England, almost an equal quantity is consumed in the
production of substances containing chlorine and sodium, and especially
in the manufacture of washing soda, &c., and of chlorine compounds
(bleaching powder and hydrochloric acid). The yearly production of salt
in Europe amounts to as much as 7-1/2 million tons.
Although certain lumps of rock salt and crystals of bay salt sometimes
consist of almost pure sodium chloride, still the ordinary commercial
salt contains various impurities, the most common of which are
magnesium salts. If the salt be pure, its solution gives no precipitate
with sodium carbonate, Na_{2}CO_{3}, showing the absence of magnesium
salts, because magnesium carbonate, MgCO_{3}, is insoluble in water.
Rock salt, which is ground for use, generally contains also a
considerable admixture of clay and other insoluble impurities.[8] For
ordinary use the bulk of the salt obtained can be employed directly
without further purification; but some salts are purified by solution
and crystallisation of the solution after standing, in which case the
evaporation is not carried on to dryness, and the impurities remain
in the _mother liquor_ or in the sediment. When perfectly pure salt
is required for chemical purposes it is best obtained as follows: a
saturated solution of table salt is prepared, and hydrochloric acid
gas is passed through it; this precipitates the sodium chloride (which
is not soluble in a strong solution of hydrochloric acid), while the
impurities remain in solution. By repeating the operation and fusing
the salt (when adhering hydrochloric acid is volatilised) a pure
salt is obtained, which is again crystallised from its solution by
evaporation.[9]
[8] The fracture of rock salt generally shows the presence of
interlayers of impurities which are sometimes very small in weight,
but visible owing to their refraction. In the excellently laid out
salt mines of Briansk I counted (1888), if my memory does not
deceive me, on an average ten interlayers per metre of thickness,
between which the salt was in general very pure, and in places
quite transparent. If this be the case, then there would be 350
interlayers for the whole thickness (about 35 metres) of the bed.
They probably correspond with the yearly deposition of the salt. In
this case the deposition would have extended over more than 300
years. This should be observable at the present day in lakes where
the salt is saturated and in course of deposition.
[9] My own investigations have shown that not only the sulphates, but
also the potassium salts, are entirely removed by this method.
Pure sodium chloride, in well-defined crystals (slowly deposited at the
bottom of the liquid) or in compact masses (in which form rock salt is
sometimes met with), is a colourless and transparent substance
resembling, but more brittle and less hard than, glass.[10] Common salt
always crystallises in the cubic system, most frequently in _cubes_, and
more rarely in octahedra. Large transparent cubes of common salt, having
edges up to 10 centimetres long, are sometimes found in masses of rock
salt.[11] When evaporated in the open the salt often separates out on
the surface[12] as cubes, which grow on to each other in the form of
pyramidal square funnels. In still weather, these clusters are able to
support themselves on the surface of the water for a long time, and
sometimes go on increasing to a considerable extent, but they sink
directly the water penetrates inside them. Salt fuses to a colourless
liquid (sp. gr. 1·602, according to Quincke) at 851° (V. Meyer); if pure
it solidifies to a non-crystalline mass, and if impure to an opaque mass
whose surface is not smooth. In fusing, sodium chloride commences to
volatilise (its weight decreases) and at a white heat it volatilises with
great ease and completely; but at the ordinary temperature it may, like
all ordinary salts, be considered as non-volatile, although as yet no
exact experiments have been made in this direction.
[10] According to the determinations of Klodt, the Briansk rock salt
withstands a pressure of 340 kilograms per square centimetre,
whilst glass withstands 1,700 kilos. In this respect salt is twice
as secure as bricks, and therefore immense masses may be extracted
from underground workings with perfect safety, without having
recourse to brickwork supports, merely taking advantage of the
properties of the salt itself.
[11] To obtain well-formed crystals, a saturated solution is mixed with
ferric chloride, several small crystals of sodium chloride are
placed at the bottom, and the solution is allowed to evaporate
slowly in a vessel with a loose-fitting cover. Octahedral crystals
are obtained by the addition of borax, urea, &c., to the solution.
Very fine crystals are formed in a mass of gelatinous silica.
[12] If a solution of sodium chloride be slowly heated from above,
where the evaporation takes place, then the upper layer will
become saturated before the lower and cooler layers, and therefore
crystallisation will begin on the surface, and the crystals first
formed will float, having also dried from above, on the surface
until they become quite soaked. Being heavier than the solution
the crystals are partially immersed under it, and the following
crystallisation, also proceeding on the surface, will only form
crystals along the side of the original crystals. A funnel is
formed in this manner. It will be borne on the surface like a boat
(if the liquid be quiescent), because it will grow more from the
upper edges. We can thus understand this at first sight strange
funnel form of crystallisation of salt. In explanation why the
crystallisation under the above conditions begins at the surface
and not at the lower layers, it must be mentioned that the
specific gravity of a crystal of sodium chloride = 2·16, and that
of a solution saturated at 25° contains 26·7 p.c. of salt and has
a specific gravity at 25°/4° of 1·2004; at 15° a saturated
solution contains 26·5 p.c. of salt and has a sp. gr. 1·203 at
15°/4°. Hence a solution saturated at a higher temperature is
specifically lighter, notwithstanding the greater amount of salt
it contains. With many substances _surface crystallisation_ cannot
take place because their solubility increases more rapidly with
the temperature than their specific gravity decreases. In this
case the saturated solution will always be in the lower layers,
where also the crystallisation will take place. Besides which it
may be added that as a consequence of the properties of water and
solutions, when they are heated from above (for instance, by the
sun's rays), the warmer layers being the lightest remain above,
whilst when heated from below they rise to the top. For this
reason the water at great depths below the surface is always cold,
which has long been known. These circumstances, as well as those
observed by Soret (Chapter I., Note 19), explain the great
differences of density and temperature, and in the amount of salts
held in the oceans at different latitudes (in polar and tropical
climes) and at various depths.
A saturated[13] solution of table salt (containing 26·4 p.c.) has at the
ordinary temperature a specific gravity of about 1·2. The specific
gravity of the crystals is 2·167 (17°). The salt which separates out at
the ordinary and higher temperatures contains no water of
crystallisation;[14] but if the crystals are formed at a low
temperature, especially from a saturated solution cooled to -12°, then
they present a prismatic form, and contain two equivalents of water,
NaCl,2H_{2}O. At the ordinary temperature these crystals split up into
sodium chloride and its solution.[15] Unsaturated solutions of table salt
when cooled below 0° give[16] crystals of ice, but when the solution has
a composition NaCl,10H_{2}O it solidifies completely at a temperature of
-23°. A solution of table salt saturated at its boiling point boils at
about 109°, and contains about 42 parts of salt per 100 parts of water.
[13] By combining the results of Poggiale, Müller, and Karsten (they
are evidently more accurate than those of Gay-Lussac and others) I
found that a saturated solution at _t_°, from 0° to 108°, contains
35·7 + 0·024_t_ + 0·0002_t_^2 grams of salt per 100 grams of
water. This formula gives a solubility at 0° = 35·7 grams (= 26·3
p.c.), whilst according to Karsten it is 36·09, Poggiale 35·5, and
Müller 35·6 grams.
[14] Perfectly pure _fused_ salt is not hygroscopic, according to
Karsten, whilst the crystallised salt, even when quite pure,
attracts as much as 0·6 p.c. of water from moist air, according to
Stas. (In the Briansk mines, where the temperature throughout the
whole year is about +10°, it may be observed, as Baron Klodt
informed me, that in the summer during damp weather the walls
become moist, while in winter they are dry).
If the salt contain impurities--such as magnesium sulphate, &c.--it
is more hygroscopic. If it contain any magnesium chloride, it
partially deliquesces in a damp atmosphere. The crystallised and
not perfectly pure salt decrepitates when heated, owing to its
containing water. The pure salt, and also the transparent rock
salt, or that which has been once fused, does not decrepitate.
Fused sodium chloride shows a faint alkaline reaction to litmus,
which has been noticed by many observers, and is due to the
presence of sodium oxide (probably by the action of the oxygen of
the atmosphere). According to A. Stcherbakoff very sensitive
litmus (washed in alcohol and neutralised with oxalic acid) shows
an alkaline reaction even with the crystallised salt.
It may be observed that rock salt sometimes contains cavities
filled with a colourless liquid. Certain kinds of rock salt emit
an odour like that of hydrocarbons. These phenomena have as yet
received very little attention.
[15] By cooling a solution of table salt saturated at the ordinary
temperature to -15°, I obtained first of all well-formed tabular
(six-sided) crystals, which when warmed to the ordinary
temperature disintegrated (with the separation of anhydrous sodium
chloride), and then prismatic needles up to 20 mm. long were
formed from the same solution. I have not yet investigated the
reason of the difference in crystalline form. It is known
(Mitscherlich) that NaI,2H_{2}O also crystallises either in plates
or prisms. Sodium bromide also crystallises with 2H_{2}O at the
ordinary temperature.
[16] Notwithstanding the great simplicity (Chapter I., Note 49) of the
observations on the formation of ice from solution, still even for
sodium chloride they cannot yet be considered as sufficiently
harmonious. According to Blagden and Raoult, the temperature of
the formation of ice from a solution containing _c_ grams of salt
per 100 grams of water =-0·6_c_ to _c_ = 10, according to Rosetti
=-0·649_c_ to _c_ = 8·7, according to De Coppet (to _c_ = 10)
=-0·55_c_-0·006_c_^2, according to Karsten (to _c_ = 10)-0·762_c_
+ 0·0084_c_^2, and according to Guthrie a much lower figure. By
taking Rosetti's figure and applying the rule given in Chapter I.,
Note 49 we obtain--
_i_ - 0·649 × 58·5/18·5 = 2·05.
Pickering (1893) gives for _c_ = 1-0·603, for _c_ = 2-1·220; that
is (_c_ up to 2·7) about -(0·600 + 0·005_c_)_c_.
The data for strong solutions are not less contradictory. Thus
with 20 p.c. of salt, ice is formed at -14·4° according to
Karsten,-17° according to Guthrie,-17·6° according to De Coppet.
Rüdorff states that for strong solutions the temperature of the
formation of ice descends in proportion to the contents of the
compound, NaCl,2H_{2}O (per 100 grams of water) by 0°·342 per 1
gram of salt, and De Coppet shows that there is no
proportionality, in a strict sense, for either a percentage of
NaCl or of NaCl,2H_{2}O.
Of all its physical properties the specific gravity of solutions of
sodium chloride is the one which has been the most fully investigated. A
comparison of all the existing determinations of the specific gravity of
solutions of NaCl[17] at 15° (in vacuo, taking water at 4° as 10,000),
with regard to _p_ (the percentage amount of the salt in solution), show
that it is expressed by the equation S_{15} = 9991·6 + 71·17_p_ +
0·2140_p_^2. For instance, for a solution 200H_{2}O + NaCl, in which case
_p_ = 1·6, S_{15} = 1·0106. It is seen from the formula that the addition
of water produces a contraction.[18] The specific gravity[19] at certain
temperatures and concentrations in vacuo referred to water at 4° =
10,000[20] is here given for
0° 15° 30° 100°
_p_ = 5 10372 10353 10307 9922
10 10768 10728 10669 10278
15 11164 11107 11043 10652
20 11568 11501 11429 11043
It should be remarked that Baumé's hydrometer is graduated by taking a
10 p.c. solution of sodium chloride as 10° on the scale, and therefore
it gives approximately the percentage amount of the salt in a solution.
Common salt is somewhat soluble in alcohol,[21] but it is insoluble in
ether and in oils.
[17] A collection of observations on the specific gravity of solutions
of sodium chloride is given in my work cited in Chapter I.,
Note 50.
Solutions of common salt have also been frequently investigated as
regards rate of _diffusion_ (Chapter I.), but as yet there are no
complete data in this respect. It may be mentioned that Graham and
De Vries demonstrated that diffusion in gelatinous masses (for
instance, gelatin jelly, or gelatinous silica) proceeds in the
same manner as in water, which may probably lead to a convenient
and accurate method for the investigation of the phenomena of
diffusion. N. Umoff (Odessa, 1888) investigated the diffusion of
common salt by means of glass globules of definite density. Having
poured water into a cylinder over a layer of a solution of sodium
chloride, he observed during a period of several months the
position (height) of the globules, which floated up higher and
higher as the salt permeated upwards. Umoff found that at a
constant temperature the distances of the globules (that is, the
length of a column limited by layers of definite concentration)
remain constant; that at a given moment of time the concentration,
_q_, of different layers situated at a depth _z_ is expressed by
the equation B-K_z_ = log.(A-_q_), where A, B, and K are
constants; that at a given moment the rate of diffusion of the
different layers is proportional to their depth, &_c._
[18] If _S__{0} be the specific gravity of water, and _S_ the specific
gravity of a solution containing _p_ p.c. of salt, then by mixing
equal weights of water and the solution, we shall obtain a
solution containing 1/2_p_ of the salt, and if it be formed
without contraction, then its specific gravity _x_ will be
determined by the equation 2/_x_ = 1/S_{0} + 1/S, because the
volume is equal to the weight divided by the density. In reality,
the specific gravity is always found to be greater than that
calculated on the supposition of an absence of contraction.
[19] Generally the specific gravity is observed by weighing in air and
dividing the weight in grams by the volume in cubic centimetres,
the latter being found from the weight of water displaced, divided
by its density at the temperature at which the experiment is
carried out. If we call this specific gravity S_{1}, then as a
cubic centimetre of air under the usual conditions weighs about
0·0012 gram, the sp. gr. in a vacuum S = S_{1} + 0·0012 (S_{1}
- 1), if the density of water = 1.
[20] If the sp. gr. S_{2} be found directly by dividing the weight of a
solution by the weight of water at the same temperature and in the
same volume, then the true sp. gr. _S_ referred to water at 4° is
found by multiplying S_{2} by the sp. gr. of water at the
temperature of observation.
[21] According to Schiff 100 grams of alcohol, containing _p_ p.c. by
weight of C_{2}H_{6}O, dissolves at 15°--
_p_ = 10 20 40 60 80
28·5 22·6 13·2 5·9 1·2 grams NaCl.
Common salt gives very few compounds[22] (double salts) and these are
very readily decomposed: it is also decomposed with great difficulty and
its dissociation is unknown.[23] But it is easily decomposed, both when
fused and in solution, by the action of a galvanic current. If the dry
salt be fused in a crucible and an electric current be passed through it
by immersing carbon or platinum electrodes in it (the positive electrode
is made of carbon and the negative of platinum or mercury), it is
_decomposed_: the suffocating gas, chlorine, is liberated at the positive
pole and metallic sodium at the negative pole. Both of them act on the
excess of water at the moment of their evolution; the sodium evolves
hydrogen and forms caustic soda, and the chlorine evolves oxygen and
forms hydrochloric acid, and therefore on passing a current through a
solution of common salt metallic sodium will not be obtained--but oxygen,
chlorine, and hydrochloric acid will appear at the positive pole, and
hydrogen and caustic soda at the negative pole.[23 bis] Thus salt, like
other salts, is decomposed by the action of an electric current into a
metal and a haloid (Chapter III.) Naturally, like all other salts, it may
be formed from the corresponding base and acid with the separation of
water. In fact if we mix caustic soda (base) with hydrochloric acid
(acid), table salt is formed, NaHO + HCl = NaCl + H_{2}O.
[22] Amongst the double salts formed by sodium chloride that obtained
by Ditte (1870) by the evaporation of the solution remaining after
heating sodium iodate with hydrochloric acid until chlorine ceases
to be liberated, is a remarkable one. Its composition is
NaIO_{3},NaCl,14H_{2}O. Rammelsberg obtained a similar (perhaps
the same) salt in well-formed crystals by the direct reaction of
both salts.
[23] But it gives sodium in the flame of a Bunsen's burner (see
Spectrum Analysis), doubtless under the reducing action of the
elements carbon and hydrogen. In the presence of an excess of
hydrochloric acid in the flame (when the sodium would form sodium
chloride), no sodium is formed in the flame and the salt does not
communicate its usual coloration.
[23 bis] There is no doubt, however, but that chloride of sodium is
also decomposed in its aqueous solutions with the separation of
sodium, and that it does not simply enter into double
decomposition with the water (NaCl + H_{2}O = NaHO + HCl). This is
seen from the fact that when a saturated solution of NaCl is
rapidly decomposed by an electric current, a large amount of
chlorine appears at the anode and a sodium amalgam forms at the
mercury cathode, which acts but slowly upon the strong solution of
salt. Castner's process for the electrolysis of brine into
chlorine and caustic soda is an application of this method which
has been already worked in England on an industrial scale.
With resspect to the double decompositions of sodium chloride it should
be observed that they are most varied, and serve as means of obtaining
nearly all the other compounds of sodium and chlorine.
_The double decompositions of sodium chloride_ are almost exclusively
based on the possibility of the metal sodium being exchanged for
hydrogen and other metals. But neither hydrogen nor any other metal can
directly displace the sodium from sodium chloride. This would result
in the separation of metallic sodium, which itself displaces hydrogen
and the majority of other metals from their compounds, and is not,
so far as is known, ever separated by them. The replacement of the
sodium in sodium chloride by hydrogen and various metals can only take
place by the transference of the sodium into some other combination.
If hydrogen or a metal, M, be combined with an element X, then the
double decomposition NaCl + MX = NaX + MCl takes place. Such double
decompositions take place under special conditions, sometimes completely
and sometimes only partially, as we shall endeavour to explain. In order
to acquaint ourselves with the double decompositions of sodium chloride,
we will follow the methods actually employed in practice to procure
compounds of sodium and of chlorine from common salt. For this purpose
we will first describe the treatment of sodium chloride with sulphuric
acid for the preparation of hydrochloric acid and sodium sulphate.
We will then describe the substances obtained from hydrochloric acid
and sodium sulphate. Chlorine itself, and nearly all the compounds of
this element, may be procured from hydrochloric acid, whilst sodium
carbonate, caustic soda, metallic sodium itself and all its compounds,
may be obtained from sodium sulphate.
Even in the animal organism salt undergoes similar changes, furnishing
the sodium, alkali, and hydrochloric acid which take part in the
processes of animal life.
Its necessity as a constituent in the food both of human beings and of
animals becomes evident when we consider that both hydrochloric acid and
salts of sodium are found in the substances which are separated out from
the blood into the stomach and intestines. Sodium salts are found in the
blood and in the bile which is elaborated in the liver and acts on the
food in the alimentary canal, whilst hydrochloric acid is found in the
acid juices of the stomach. Chlorides of the metals are always found in
considerable quantities in the urine, and if they are excreted they must
be replenished in the organism; and for the replenishment of the loss,
substances containing chlorine compounds must be taken in food. Not only
do animals consume those small amounts of sodium chloride which are
found in drinking water or in plants or other animals, but experience
has shown that many wild animals travel long distances in search of salt
springs, and that domestic animals which in their natural condition do
not require salt, willingly take it, and that the functions of their
organisms become much more regular from their doing so.
_The action of sulphuric acid on sodium chloride._--If sulphuric acid
be poured over common salt, then even at the ordinary temperature, as
Glauber observed, an odorous gas, hydrochloric acid, is evolved. The
reaction which takes place consists in the sodium of the salt and the
hydrogen of the sulphuric acid changing places.
NaCl + H_{2}SO_{4} = HCl + NaHSO_{4}
Sodium Sulphuric Hydrochloric Acid sodium
chloride acid acid sulphate
At the ordinary temperature this reaction is not complete, but soon
ceases. When the mixture is heated, the decomposition proceeds until,
if there be sufficient salt present, all the sulphuric acid taken is
converted into acid sodium sulphate. Any excess of acid will remain
unaltered. If 2 molecules of sodium chloride (117 parts) be taken per
molecule of sulphuric acid (98 parts), then on heating the mixture to
a moderate temperature only one-half (58·5) of the salt will suffer
change. Complete decomposition, after which neither hydrogen nor
chlorine is left in the residue, proceeds (when 117 parts of table salt
are taken per 98 parts of sulphuric acid) _at a red heat only_. Then--
2NaCl + H_{2}SO_{4} = 2HCl + Na_{2}SO_{4}
Table Sulphuric Hydrochloric Sodium
salt acid acid sulphate
This double decomposition is the result of the action of the acid
salt, NaHSO_{4}, first formed, on sodium chloride, for the acid salt,
since it contains hydrogen, itself acts like an acid, NaCl + NaHSO_{4}
= HCl + Na_{2}SO_{4}. By adding this equation to the first we obtain
the second, which expresses the ultimate reaction. Hence in the above
reaction, non-volatile or sparingly volatile table salt and sparingly
volatile sulphuric acid are taken, and as the result of their reaction,
after the hydrogen and sodium have exchanged places, there is obtained
non-volatile sodium sulphate and gaseous hydrochloric acid. The fact
of the latter being a gaseous substance forms the main reason for the
reaction proceeding to the very end. The mechanism of this kind of
double decomposition, and the cause of the course of the reaction, are
exactly the same as those we saw in the decomposition of nitre (Chapter
VI.) by the action of sulphuric acid. The sulphuric acid in each case
displaces the other, volatile, acid.
Not only in these two instances, but in every instance, if a volatile
acid can be formed by the substitution of the hydrogen of sulphuric
acid for a metal, then this volatile acid will be formed. From this it
may be concluded that the volatility of the acid should be considered
as the cause of the progress of the reaction; and indeed if the acid be
soluble but not volatile, or if the reaction take place in an enclosed
space where the resulting acid cannot volatilise, or at the ordinary
temperature when it does not pass into the state of elastic vapour--then
the decomposition does not proceed to the end, but only up to a certain
limit. In this respect the explanations given at the beginning of this
century by the French chemist Berthollet in his work 'Essai de Statique
Chimique' are very important. _The doctrine of Berthollet_ starts from
the supposition that the chemical reaction of substances is determined
not only by the degrees of affinity between the different parts, but
also by the relative masses of the reacting substances and by those
physical conditions under which the reaction takes place. Two substances
containing the elements MX and NY, being brought into contact with
each other, form by double decomposition the compounds MY and NX; but
the formation of these two new compounds will not proceed to the end
unless one of them is removed from the sphere of action. But it can
only be removed if it possesses different physical properties from
those of the other substances which are present with it. Either it must
be a gas while the others are liquid or solid, or an insoluble solid
while the others are liquid or soluble. The relative amounts of the
resultant substances, if nothing separates out from their intermixture,
depend only on the relative quantities of the substances MX and NY, and
upon the degrees of attraction existing between the elements M, N, X,
and Y; but however great their mass may be, and however considerable
the attractions, still in any case if nothing separates out from the
sphere of action the decomposition will presently cease, a state of
equilibrium will be established, and instead of two there will remain
four substances in the mass: namely, a portion of the original bodies
MX and NY, and a certain quantity of the newly formed substances MY and
NX, if it be assumed that neither MN or XY nor any other substances are
produced, and this may for the present[24] be admitted in the case of
the double decomposition of salts in which M and X are metals and X and
Y haloids. As the ordinary double decomposition here consists merely
in the exchange of metals, the above simplification is applicable. The
sum total of existing data concerning the double decomposition of salts
leads to the conclusion that from salts MX + NY there always arises
a certain quantity of NX and MY, as should be the case according to
Berthollet's doctrine. A portion of the historical data concerning
this subject will be afterwards mentioned, but we will at once proceed
to point out the observations made by Spring (1888) which show that
_even in a solid state_ salts are subject to a similar interchange of
metals if in a condition of sufficiently close contact (it requires
time, a finely divided state, and intimate mixture). Spring took two
non-hygroscopic salts, potassium nitrate, KNO_{3}, and well-dried sodium
acetate, C_{2}H_{3}NaO_{2}, and left a mixture of their powders for
several months in a desiccator. An interchange of metals took place, as
was seen from the fact that the resultant mass rapidly attracted the
moisture of the air, owing to the formation of sodium nitrate, NaNO_{3},
and potassium acetate, C_{2}H_{3}KO_{2}, both of which are highly
hygroscopic.[24 bis]
[24] If MX and NY represent the molecules of two salts, and if there be
_no third substance_ present (such as water in a solution), the
formation of XY would also be possible; for instance, cyanogen,
iodine, &c. are capable of combining with simple haloids, as well
as with the complex groups which in certain salts play the part of
haloids. Besides which the salts MX and NY or MY with NX may form
double salts. If the number of molecules be unequal, or if the
valency of the elements or groups contained in them be different,
as in NaCl + H_{2}SO_{4}, where Cl is a univalent haloid and
SO_{4} is bivalent, then the matter may be complicated by the
formation of other compounds besides MY and NX, and when a solvent
participates in the action, and especially if present in large
proportion, the phenomena must evidently become still more
complex; and this is actually the case in nature. Hence while
placing before the reader a certain portion of the existing store
of knowledge concerning the phenomena of double saline
decompositions, I cannot consider the theory of the subject as
complete, and have therefore limited myself to a few data, the
completion of which must be sought in more detailed works on the
subject of theoretical chemistry, without losing sight of what has
been said above.
[24 bis] When the mixture of potassium nitrate and sodium acetate was
heated by Spring to 100°, it was completely fused into one mass,
although potassium nitrate fuses at about 340° and sodium nitrate
at about 320°.
When Berthollet enunciated his doctrine the present views of atoms and
molecules had yet to be developed, and it is now necessary to submit the
matter to examination in the light of these conceptions; we will
therefore consider the reaction of salts, taking M and N, X and Y as
equivalent to each other--that is, as capable of replacing each other 'in
toto,' as Na or K,, 1/2Ca or 1/2Mg (bivalent elements) replace hydrogen.
And since, according to Berthollet's doctrine, when _m_MX of one salt
comes into contact with _n_NY of another salt, a certain quantity _x_MY
and _x_NX is formed, there remains _m_-_x_ of the salt MX, and _n_-_x_
of the salt NY. If _m_ be greater than _n_, then the maximum interchange
could lead to _x_ = _n_, whilst from the salts taken there would be
formed _n_MY + _n_NX + (_m_-_n_)MX--that is, a portion of one only of
the salts taken would remain unchanged because the reaction could only
proceed between _n_MX and _n_NY. If _x_ were actually equal to _n_, the
mass of the salt MX would not have any influence on the _modus operandi_
of the reaction, which is equally in accordance with the teaching of
Bergmann, who supposed double reactions to be independent of the mass
and determined by affinity only. If M had more affinity for X than for
Y, and N more affinity for Y than for X, then according to Bergmann
there would be no decomposition whatever, and _x_ would equal 0. If
the affinity of M for Y and of N for X were greater than those in the
original grouping, then the affinity of M for X and of N for Y would be
overcome, and, according to Bergmann's doctrine, complete interchange
would take place--_i.e._ _x_ would equal _n_. According to Berthollet's
teaching, a distribution of M and N between X and Y will take place
in every case, not only in proportion to the degrees of affinity, but
also in proportion to the masses, so that with a small affinity and a
large mass the same action can be produced as with a large affinity
and a small mass. Therefore, (1) _x_ will always be less than _n_ and
their ratio _x_/_n_ less than unity--that is, the decomposition will be
expressed by the equation, _m_MX + _n_NY = (_m_-_x_)MX + (_n_-_x_)NY
+ _x_MY + _x_NX; (2) by increasing the mass _m_ we increase the
decomposition--that is, we increase _x_ and the ratio _x_/(_n_-_x_),
until with an infinitely large quantity m the fraction _x_/_n_ will
equal 1, and the decomposition will be complete, however small the
affinities uniting MY and NX may be; and (3) if _m_ = _n_, by taking MX
+ NY or MY + NX we arrive at one and the same system _in either case_:
(_n_-_x_)MX + (_n_-_x_)NY + _x_MY + _x_NX. These direct consequences of
Berthollet's teaching are verified by experience. Thus, for example,
a mixture of solutions of sodium nitrate and potassium chloride in
all cases has entirely the same properties as a mixture of solutions
of potassium nitrate and sodium chloride, of course on condition that
the mixed solutions are of identical elementary composition. But this
identity of properties might either proceed from one system of salts
passing entirely into the other (Bergmann's hypothesis) in conformity
with the predominating affinities (for instance, from KCl + NaNO_{3}
there might arise KNO_{3} + NaCl, if it be admitted that the affinities
of the elements as combined in the latter system are greater than in
the former); or, on the other hand, it might be because both systems by
the interchange of a portion of their elements give one and the same
state of equilibrium, as according to Berthollet's teaching. Experiment
proves the latter hypothesis to be the true one. But before citing
the most historically important experiments verifying Berthollet's
doctrine, we must stop to consider the conception _of the mass_ of the
reacting substances. Berthollet understood by mass the actual relative
quantity of a substance; but now it is impossible to understand this
term otherwise than as the number of molecules, for they act as chemical
units, and in the special case of double saline decompositions it is
better to take it as the number of equivalents. Thus in the reaction
NaCl + H_{2}SO_{4} the salt is taken in one equivalent and the acid in
two. If 2NaCl + H_{2}SO_{4} act, then the number of equivalents are
equal, and so on. The _influence of mass_ on the amount of decomposition
_x_/_n_ forms the root of Berthollet's doctrine, and therefore we will
first of all turn our attention to the establishment of this principle
in relation to the double decomposition of salts.
About 1840 H. Rose[25] showed that water decomposes metallic sulphides
like calcium sulphide, CaS, forming hydrogen sulphide, H_{2}S,
notwithstanding the fact that the affinity of hydrogen sulphide, as an
acid, for lime, CaH_{2}O_{2}, as a base, causes them to react on each
other, forming calcium sulphide and water, CaS + 2H_{2}O. Furthermore,
Rose showed that the greater the amount of water acting on the calcium
sulphide, the more complete is the decomposition. The results of this
reaction are evident from the fact that the hydrogen sulphide formed
may be expelled from the solution by heating, and that the resulting
lime is sparingly soluble in water. Rose clearly saw from this that such
feeble agents, in a chemical sense, as carbonic anhydride and water, by
acting in a mass and for long periods of time in nature on the durable
rocks, which resist the action of the most powerful acids, are able
to bring about chemical change--to extract, for example, from rocks
the bases, lime, soda, potash. The influence of the mass of water on
antimonious chloride, bismuth nitrate, &c., is essentially of the same
character. These substances give up to the water a quantity of acid
which is greater in proportion as the mass of the water acting on them
is greater.[25 bis]
[25] H. Rose is more especially known for his having carefully studied
and perfected several methods for the exact chemical analysis of
many mineral substances. His predecessor in this branch of
research was Berzelius, and his successor Fresenius.
[25 bis] Historically the influence of the mass of water was the first
well-observed phenomenon in support of Berthollet's teaching, and
it should not now be forgotten. In double decompositions taking
place in dilute solutions where the mass of water is large, its
influence, notwithstanding the weakness of affinities, must he
great, according to the very essence of Berthollet's doctrine.
As explaining the action of the mass of water, the experiments of
Pattison Muir (1879) are very instructive. These experiments
demonstrate that the decomposition of bismuth chloride is the more
complete the greater the relative quantity of water, and the less
the mass of hydrochloric acid forming one of the products of the
reaction.
Barium sulphate, BaSO_{4}, which is insoluble in water, when fused with
sodium carbonate, Na_{2}CO_{3}, gives, but not completely, barium
carbonate, BaCO_{3}, (also insoluble), and sodium sulphate, Na_{2}SO_{4}.
If a solution of sodium carbonate acts on precipitated barium sulphate,
the same decomposition is also effected (Dulong, Rose), but it is
restricted by a limit and requires time. A mixture of sodium carbonate
and sulphate is obtained in the solution and a mixture of barium
carbonate and sulphate in the precipitate. If the solution be decanted
off and a fresh solution of sodium carbonate be poured over the
precipitate, then a fresh portion of the barium sulphate passes into
barium carbonate, and so by increasing the mass of sodium carbonate it is
possible to entirely convert the barium sulphate into barium carbonate.
If a definite quantity of sodium sulphate be added to the solution of
sodium carbonate, then the latter will have no action whatever on the
barium sulphate, because then a system in equilibrium determined by the
reverse action of the sodium sulphate on the barium carbonate and by the
presence of both sodium carbonate and sulphate in the solution, is at
once arrived at. On the other hand, if the mass of the sodium sulphate in
the solution be great, then the barium carbonate is reconverted into
sulphate until a definite state of equilibrium is attained between the
two opposite reactions, producing barium carbonate by the action of the
sodium carbonate and barium sulphate by the action of the sodium
sulphate.
Another most important principle of Berthollet's teaching is the
existence of _a limit of exchange decomposition_, or _the attainment
of a state of equilibrium_. In this respect the determinations of
Malaguti (1857) are historically the most important. He took a mixture
of solutions of equivalent quantities of two salts, MX and NY, and
judged the amount of the resulting exchange from the composition of the
precipitate produced by the addition of alcohol. When, for example, zinc
sulphate and sodium chloride (ZnSO_{4} and 2NaCl) were taken, there
were produced by exchange sodium sulphate and zinc chloride. A mixture
of zinc sulphate and sodium sulphate was precipitated by an excess of
alcohol, and it appeared from the composition of the precipitate that
72 per cent. of the salts taken had been decomposed. When, however, a
mixture of solutions of sodium sulphate and zinc chloride was taken, the
precipitate presented the same composition as before--that is, about 28
per cent. of the salts taken had been subjected to decomposition. In
a similar experiment with a mixture of sodium chloride and magnesium
sulphate, 2NaCl + MgSO_{4} or MgCl_{2} + Na_{2}SO_{4}, about half of
the metals underwent the decomposition, which may be expressed by
the equation 4NaCl + 2MgSO_{4} = 2NaCl + MgSO_{4} + Na_{2}SO_{4} +
MgCl_{2} = 2Na_{2}SO_{4} + 2MgCl_{2}. A no less clear limit expressed
itself in another of Malaguti's researches when he investigated the
above-mentioned reversible reactions of the insoluble salts of barium.
When, for example, barium carbonate and sodium sulphate (BaCO_{3} +
Na_{2}SO_{4}) were taken, then about 72 per cent. of the salts were
decomposed, that is, were converted into barium sulphate and sodium
carbonate. But when the two latter salts were taken, then about 19
per cent. of them passed into barium carbonate and sodium sulphate.
Probably the end of the reaction was not reached in either case, because
this would require a considerable time and a uniformity of conditions
attainable with difficulty.
Gladstone (1855) took advantage of the colour of solutions of different
ferric salts for determining the measure of exchange between metals.
Thus a solution of ferric thiocyanate has a most intense red colour, and
by making a comparison between the colour of the resulting solutions and
the colour of solutions of known strength it was possible to judge to a
certain degree the quantity of the thiocyanate formed. This colorimetric
method of determination has an important significance as being the
first in which a method was applied for determining the composition
of a solution without the removal of any of its component parts. When
Gladstone took equivalent quantities of ferric nitrate and potassium
thiocyanate--Fe(NO_{3})_{3} + 3KCNS--only 13 per cent. of the salts
underwent decomposition. On increasing the mass of the latter salt the
quantity of ferric thiocyanate formed increased, but even when more than
300 equivalents of potassium thiocyanate were taken a portion of the
iron still remained as nitrate. It is evident that the affinity acting
between Fe and NO_{3} and between K and CNS on the one hand, is greater
than the affinity acting between Fe and CNS, together with the affinity
of K for NO_{3}, on the other hand. The investigation of the variation
of the fluorescence of quinine sulphate, as well as the variation of the
rotation of the plane of polarisation of nicotine, gave in the hands
of Gladstone many proofs of the entire applicability of Berthollet's
doctrine, and in particular demonstrated the influence of mass which
forms the chief distinctive feature of the teaching of Berthollet,
teaching little appreciated in his own time.
At the beginning of the year 1860, the doctrine of the limit of reaction
and of the influence of mass on the process of chemical transformations
received a very important support in the researches of Berthelot and
P. de Saint-Gilles on the formation of the ethereal salts RX from the
alcohols ROH and acids HX, when water is also formed. This conversion is
essentially very similar to the formation of salts, but differs in that
it proceeds slowly at the ordinary temperature, extending over whole
years, and is not complete--that is, it has a distinct limit determined
by a reverse reaction; thus an ethereal salt RX with water gives an
alcohol ROH and an acid HX--up to that limit generally corresponding
with two-thirds of the alcohol taken, if the action proceed between
molecular quantities of alcohol and acid. Thus common alcohol,
C_{2}H_{5}OH, with acetic acid, HC_{2}H_{3}O_{2}, gives the following
system rapidly when heated, or slowly at the ordinary temperature, ROH
+ HX + 2RX + 2H_{2}O, whether we start from 3RHO + 3HX or from 3RX +
3H_{2}O. The process and completion of the reaction in this instance
are very easily observed, because the quantity of free acid is easily
determined from the amount of alkali requisite for its saturation, as
neither alcohol nor ethereal salt acts on litmus or other reagent for
acids. Under the influence of an increased mass of alcohol the reaction
proceeds further. If two molecules of alcohol, RHO, be taken for every
one molecule of acetic acid, HX, then instead of 66 p.c., 83 p.c. of
the acid passes into ethereal salt, and with fifty molecules of RHO
nearly all the acid is etherised. The researches of Menschutkin in their
details touched on many important aspects of the same subject, such as
the influence of the composition of the alcohol and acid on the limit
and rate of exchange--but these, as well as other details, must be
looked for in special treatises on organic and theoretical chemistry.
In any case the study of etherification has supplied chemical mechanics
with clear and valuable data, which directly confirm the two fundamental
propositions of Berthollet; the influence of mass, and the limit of
reaction--that is, the equilibrium between opposite reactions. The study
of numerous instances of dissociation which we have already touched on,
and shall again meet with on several occasions, gave the same results.
With respect to double saline decompositions, it is also necessary to
mention the researches of Wiedemann on the decomposing action of a mass
of water on the ferric salts, which could be determined by measuring the
magnetism of the solutions, because the ferric oxide (soluble colloid)
set free by the water is less magnetic than the ferric salts.
A very important epoch in the history of Berthollet's doctrine was
attained when, in 1867, the Norwegian chemists, Guldberg and Waage,
expressed it as an algebraical formula. They defined the active mass
as the number of molecules contained in a given volume, and assumed,
as follows from the spirit of Berthollet's teaching, that the action
between the substances was equal to the product of the masses of the
reacting substances. Hence if the salts MX and NY be taken in equivalent
quantities (_m_ = 1 and _n_ = 1) and the salts MY and NX are not
added to the mixture but proceed from it, then if _k_ represent the
coefficient of the rate of the action of MX on NY and if _k_´ represent
the same coefficient for the pair MY and NX, then we shall have at the
moment when the decomposition equals x a measure of action for the first
pair: _k_(1-_x_)(1-_x_) and for the second pair _k´xx_, and a state of
equilibrium or limit will be reached when _k_(1-_x_)^2 = _k_´_x_^2,
whence the ratio _k_/_k_´ = [_x_/(1-_x_)]^2. Therefore in the case of
the action of alcohol on an acid, when _x_ = 2/3, the magnitude _k_/_k_´
= 4, that is, the reaction of the alcohol on the acid is four times as
fast as that of the ethereal salt on water. If the ratio _k_/_k_´ be
known, then the influence of mass may be easily determined from it. Thus
if instead of one molecule of alcohol two be taken, then the equation
will be _k_(2-_x_)(1-_x_) = _k´xx_, whence _x_ = 0·85 or 85 percent.,
which is close to the result of experiment. If 300 molecules of alcohol
be taken, then x proves to be approximately 100 per cent., which is also
found to be the case by experiment.[26]
[26] From the above it follows that an excess of acid should influence
the reaction like an excess of alcohol. It is in fact shown by
experiment that if two molecules of acetic acid be taken to one
molecule of alcohol, 84 p.c. of alcohol is etherified. If with a
large preponderance of acid or of alcohol certain discrepancies
are observed, their cause must be looked for in the incomplete
correspondence of the conditions and external influences.
But it is impossible to subject the formation of salts to any process
directly analogous to that which is so conveniently effected in
etherification. Many efforts have, however, been made to solve the
problem of the measure of reaction in this case also. Thus, for example,
Khichinsky (1866), Petrieff (1885), and many others investigated the
distribution of metals and haloid groups in the case of one metal and
several haloids taken in excess, as acids; or conversely with an excess
of bases, the distribution of these bases with relation to an acid; in
cases where a portion of the substances forms a precipitate and a portion
remains in solution. But such complex cases, although they in general
confirm Berthollet's teaching (for instance, a solution of silver nitrate
gives some silver oxide with lead oxide, and a solution of nitrate of
lead precipitates some lead oxide under the action of silver oxide, as
Petrieff demonstrated), still, owing to the complexity of the phenomena
(for instance, the formation of basic and double salts), they cannot give
simple results. But much more instructive and complete are researches
like those made by Pattison Muir (1876), who took the simple case of the
precipitation of calcium carbonate, CaCO_{3}, from the mixture of
solutions of calcium chloride and sodium or potassium carbonate, and
found in this case that not only was the rate of action (for example, in
the case of CaCl_{2} + Na_{2}CO_{3}, 75 per cent. of CaCO_{3} was
precipitated in five minutes, 85 per cent. in thirty minutes, and 94 per
cent. in two days) determined by the temperature, relative mass, and
amount of water (a large mass of water decreases the rate), but that the
limit of decomposition was also dependent on these influences. However,
even in researches of this kind the conditions of reaction are
complicated by the non-uniformity of the media, inasmuch as a portion of
the substance is obtained or remains in the form of a precipitate, so
that the system is heterogeneous. The investigation of double saline
decompositions offers many difficulties which cannot be considered as yet
entirely overcome. Although many efforts have long since been made, the
majority of the researches were carried on in aqueous solutions, and as
water is itself a saline compound and able to combine with salts and
enter into double decomposition with them, such reactions taking place in
solutions in reality present very complex cases.[27] In this sense the
reaction between alcohols and acids is much more simple, and therefore
its significance in confirmation of Berthollet's doctrine is of
particular importance. The only cases which can be compared with these
reactions for simplicity are those exchange decompositions investigated
by G. G. Gustavson, which take place between CCl_{4} and RBr_{n} on the
one hand, and CBr_{4} and RCl_{n} on the other. This case is convenient
for investigation inasmuch as the RCl_{n} and RBr_{n} taken (such as
BCl_{3}, SiCl_{4}, TiCl_{4}, POCl_{3}, and SnCl_{4}) belong to those
substances which are decomposed by water, whilst CCl_{4} and CBr_{4} are
not decomposed by water; and therefore, by heating, for instance, a
mixture of CCl_{4} + SiBr_{4} it is possible to arrive at a conclusion as
to the amount of interchange by treating the product with water, which
decomposes the SiBr_{4} left unchanged and the SiCl_{4} formed by the
exchange, and therefore by determining the composition of the product
acted on by the water it is possible to form a conclusion as to the
amount of decomposition. The mixture was always formed with equivalent
quantities--for instance, 4BCl_{3} + 3CBr_{4}. It appeared that there was
no exchange whatever on simple intermixture, but that it proceeded
slowly, when the mixture was heated (for example, with the mixture above
mentioned at 123° 4·86 per cent. of Cl was replaced by Br after 14 days'
heating, and 6·83 per cent. after 28 days, and 10·12 per cent. when
heated at 150° for 60 days). A limit was always reached which
corresponded with that of the complemental system; in the given instance
the system 4BBr_{3} + 3CCl_{4}. In this last 89·97 per cent. of bromine
in the BBr_{3} was replaced by chlorine; that is, there were obtained
89·97 molecules of BCl_{3} and there remained 10·02 molecules of BBr_{3},
and therefore the same state of equilibrium was reached as that given by
the system 4BCl_{3} + 3CBr_{4}. Both systems gave one and the same state
of equilibrium at the limit, which is in agreement with Berthollet's
doctrine.[28]
[27] As an example two methods may be mentioned, Thomsen's and
Ostwald's. Thomsen (1869) applied a thermochemical method to
exceedingly dilute solutions without taking the water into further
consideration. He took solutions of caustic soda containing
100H_{2}O per NaHO, and sulphuric acid containing 1/2H_{2}SO_{4} +
100H_{2}O. In order that these solutions may be mixed in such
quantities that atomic proportions of acid and alkali would act,
for forty grams of caustic soda (which answers to its equivalent)
there should be employed 49 grams of sulphuric acid, and then
+15,689 heat units would be evolved. If the normal sodium sulphate
so formed be mixed with _n_ equivalents of sulphuric acid, a
certain amount of heat is absorbed, namely a quantity equal to
(_n_.1650)/(_n_ + 0·8) heat units. An equivalent of caustic soda,
in combining with an equivalent of nitric acid, evolves +13,617
units of heat, and the augmentation of the amount of nitric acid
entails an absorption of heat for each equivalent equal to -27
units; so also in combining with hydrochloric acids +13,740 heat
units are absorbed, and for each equivalent of hydrochloric acid
beyond this amount there are absorbed -32 heat units. Thomsen mixed
each one of three neutral salts, sodium sulphate, sodium chloride
and sodium nitrate, with an acid which is not contained in it; for
instance, he mixed a solution of sodium sulphate with a solution
of nitric acid and determined the number of heat units then
absorbed. An absorption of heat ensued because a normal salt was
taken in the first instance, and the mixture of all the above
normal salts with acid produces an absorption of heat. The amount
of heat absorbed enabled him to obtain an insight into the process
taking place in this mixture, for sulphuric acid added to sodium
sulphate absorbs a considerable quantity of heat, whilst
hydrochloric and nitric acids absorb a very small amount of heat
in this case. By mixing an equivalent of sodium sulphate with
various numbers of equivalents of nitric acid, Thomsen observed
that the amount of heat absorbed increased more and more as the
amount of nitric acid was increased; thus when HNO_{3} was taken
per 1/2Na_{2}SO_{4}, 1,752 heat units were absorbed per equivalent
of soda contained in the sodium sulphate. When twice as much
nitric acid was taken, 2,026 heat units, and when three times as
much, 2,050 heat units were absorbed. Had the double decomposition
been complete in the case where one equivalent of nitric acid was
taken per equivalent of Na_{2}SO_{4} then according to calculation
from similar data there should have been absorbed -2,989 units of
heat, while in reality only -1,752 units were absorbed. Hence
Thomsen concluded that a displacement of only about two-thirds of
the sulphuric acid had taken place--that is, the ratio _k_ : _k_´
for the reaction 1/2Na_{2}SO_{4} + HNO_{3} and NaNO_{3} +
1/2H_{2}SO_{4} is equal, as for ethereal salts, to 4. By taking
this figure and admitting the above supposition, Thomsen found
that for all mixtures of soda with nitric acid, and of sodium
nitrate with sulphuric acid, the amounts of heat followed Guldberg
and Waage's law; that is, the limit of decomposition reached was
greater the greater the mass of acid added. The relation of
hydrochloric to sulphuric acid gave the same results. Therefore
the researches of Thomsen fully confirm the hypotheses of Guldberg
and Waage and the doctrine of Berthollet.