conductivity of the alloys of antimony and copper with lead, came
to the conclusion that only two definite compounds of antimony and
copper exist, whilst the other alloys are either alloys of these
two together or with antimony or with copper. These compounds are
Cu_{2}Sb and Cu_{4}Sb--one corresponds with the maximum, and the
other with the minimum, electrical resistance. In general, the
resistance offered to an electrical current forms one of the
methods by which the composition of definite alloys (for example,
Pb_{2}Zn_{7}) is often established, whilst the electromotive force
of alloys affords (Laurie, 1888) a still more accurate method--for
instance, several definite compounds were discovered by this
method among the alloys of copper with zinc and tin; but we will
not enter into any details of this subject, because we avoid all
references to electricity, although the reader is recommended to
make himself acquainted with this branch of science, which has
many points in common with chemistry. The study of alloys regarded
as solid solutions should, in my opinion, throw much light upon
the question of solutions, which is still obscure in many aspects
and in many branches of chemistry.
Copper, in its cuprous compounds, is so analogous to _silver_, that were
there no cupric compounds, or if silver gave stable compounds of the
higher oxide, AgO, the resemblance would be as close as that between
chlorine and bromine or zinc and cadmium; but silver compounds
corresponding to AgO are quite unknown. Although silver peroxide--which
was regarded as AgO, but which Berthelot (1880) recognised as the
sesquioxide Ag_{2}O_{3}--is known, still it does not form any true salts,
and consequently cannot be placed along with cupric oxide. In distinction
to copper, silver as a metal does not oxidise under the influence of
heat; and its oxides, Ag_{2}O and Ag_{2}O_{3}, easily lose oxygen (_see_
Note 8 tri). Silver does _not oxidise_ in air at the ordinary pressure,
and is therefore classed among the so-called _noble metals_. It has a
white colour, which is much purer than that of any other known metal,
especially when the metal is chemically pure. In the arts silver is
always used alloyed, because chemically-pure silver is so soft that it
wears exceedingly easily, whilst when fused with a small amount of
copper, it becomes very hard, without losing its colour.[13]
[13] There are not many soft metals; lead, tin, copper, silver, iron,
and gold are somewhat soft, and potassium and sodium very soft.
The metals of the alkaline earths are sonorous and hard, and many
other metals are even brittle, especially bismuth and antimony.
But the very slight significance which these properties have in
determining the fundamental chemical properties of substances
(although, however, of immense importance in the practical
applications of metals) is seen from the example shown by zinc,
which is hard at the ordinary temperature, soft at 100°, and
brittle at 200°.
[Illustration: FIG. 95.--Cupel for silver assaying.]
[Illustration: FIG. 96.--Clay muffle.]
[Illustration: FIG. 97.--Portable muffle furnace.]
As the value of silver depends exclusively on its purity, and as
there is no possibility of telling the amount of impurities
alloyed with it from its external appearance, it is customary in
most countries to mark an article with the amount of pure silver
it contains after an accurately-made analysis known as the assay
of the silver. In France the assay of silver shows the amount of
pure silver in 1,000 parts by weight; in Russia the amount of pure
silver in 96 parts--that is, the assay shows the number of
zolotniks (4·26 grams) of pure silver in one pound (410 grams) of
alloyed silver. Russian silver is generally 84 assay--that is,
contains 84 parts by weight of pure silver and 12 parts of copper
and other metals. French money contains 90 p.c. (in the Russian
system this will be 86·4 assay) by weight of silver [English coins
and jewellery contain 92·5 p.c. of silver]; the silver rouble is
of 83-1/3 assay--that is, it contains 86·8 p.c. of silver--and the
smaller Russian silver coinage is of 48 assay, and therefore
contains 50 p.c. of silver. Silver ornaments and articles are
usually made in Russia of 84 and 72 assay. As the alloys of silver
and copper, especially after being subjected to the action of
heat, are not so white as pure silver, they generally undergo a
process known as 'blanching' (or 'pickling') after being worked
up. This consists in removing the copper from the surface of the
article by subjecting it to a dark-red heat and then immersing it
in dilute acid. During the calcination the copper on the surface
is oxidised, whilst the silver remains unchanged; the dilute acid
then dissolves the copper oxides formed, and pure silver is left
on the surface. The surface is dull after this treatment, owing to
the removal of a portion of the metal by the acid. After being
polished the article acquires the desired lustre and colour, so as
to be indistinguishable from a pure silver object. In order to
test a silver article, a portion of its mass must be taken, not
from the surface, but to a certain depth. The methods of assay
used in practice are very varied. The commonest and most often
used is that known as _cupellation_. It is based on the difference
in the oxidisability of copper, lead, and silver. The cupel is a
porous cup with thick sides, made by compressing bone ash. The
porous mass of bone ash absorbs the fused oxides, especially the
lead oxide, which is easily fusible, but it does not absorb the
unoxidised metal. The latter collects into a globule under the
action of a strong heat in the cupel, and on cooling solidifies
into a button, which may then be weighed. Several cupels are
placed in a muffle. A muffle is a semi-cylindrical clay vessel,
shown in the accompanying drawing. The sides of the muffle are
pierced with several orifices, which allow the access of air into
it. The muffle is placed in a furnace, where it is strongly
heated. Under the action of the air entering the muffle the copper
of the silver alloy is oxidised, but as the oxide of copper is
infusible, or, more strictly speaking, difficultly fusible, a
certain quantity of lead is added to the alloy; the lead is also
oxidised by the air at the high temperature of the muffle, and
gives the very fusible lead oxide. The copper oxide then fuses
with the lead oxide, and is absorbed by the cupel, whilst the
silver remains as a bright white globule. If the weight of the
alloy taken and of the silver left on the cupel be determined, it
is possible to calculate the composition of the alloy. Thus the
essence of cupellation consists in the separation of the
oxidisable metals from silver, which does not oxidise under the
action of heat. A more accurate method, based on the precipitation
of silver from its solutions in the form of silver chloride, is
described in detail in works on analytical chemistry.
Silver occurs in _nature_, both in a native state and in certain
compounds. Native silver, however, is of rather rare occurrence. A far
greater quantity of silver occurs in combination with sulphur, and
especially in the form of _silver sulphide_, Ag_{2}S, with lead sulphide
or copper sulphide, or the ores of various other metals. The largest
amount of silver is extracted from the lead in which it occurs. If this
lead be calcined in the presence of air, it oxidises, and the resultant
lead oxide, PbO ('litharge' or 'silberglätte,' as it is called), melts
into a mobile liquid, which is easily removed. The silver remains in an
unoxidised metallic state.[14] This process is called _cupellation_.
[14] In America, whence the largest amount of silver is now obtained,
ores are worked containing not more than 1/5 p.c. of silver,
whilst at 1/2 p.c. its extraction is very profitable. Moreover,
the extraction of silver from ores containing not more than 0·01
p.c. of this metal is sometimes profitable. The majority of the
lead smelted from galena contains silver, which is extracted from
it. Thus near Arras, in France, an ore is worked which contains
about 65 parts of lead and 0·088 part of silver in 100 parts of
ore, which corresponds with 136 parts of silver in 100,000 parts
of lead. At Freiberg, in Saxony, the ore used (enriched by
mechanical dressing) contains about 0·9 of silver, 160 of lead,
and 2 of copper in 10,000 parts. In every case the lead is first
extracted in the manner described in Chapter XVIII., and this lead
will contain all the silver. Not unfrequently other ores of silver
are mixed with lead ores, in order to obtain an argentiferous lead
as the product. The extraction of small quantities of silver from
lead is facilitated by the fact (Pattinson's process) that molten
argentiferous lead in cooling first deposits crystals of pure
lead, which fall to the bottom of the cooling vessel, whilst the
proportion of silver in the unsolidified mass increases owing to
the removal of the crystals of lead. The lead is enriched in this
manner until it contains 1/400 part of silver, and is then
subjected to cupellation on a larger scale. According to Park's
process, zinc is added to the molten argentiferous lead, and the
alloy of Pb and Zn, which first separates out on cooling, is
collected. This alloy is found to contain all the silver
previously contained in the lead. The addition of 0·5 p.c. of
aluminium to the zinc (Rossler and Edelman) facilitates the
extraction of the Ag from the resultant alloy besides preventing
oxidation; for, after re-melting, nearly all the lead easily runs
off (remains fluid), and leaves an alloy containing about 30 p.c.
Ag and about 70 p.c. Zn. This alloy may be used as an anode in a
solution of ZnCl_{2}, when the Zn is deposited on the cathode,
leaving the silver with a small amount of Pb, &c. behind. The
silver can be easily obtained pure by treating it with dilute
acids and cupelling.
The ores of silver which contain a larger amount of it are: silver
glance, Ag_{2}S (sp. gr. 7·2); argentiferous-copper glance, CuAgS;
horn silver or chloride of silver, AgCl; argentiferous grey copper
ore; polybasite, M_{9}RS_{6} (where M = Ag, Cu, and R = Sb, As),
and argentiferous gold. The latter is the usual form in which gold
is found in alluvial deposits and ores. The crystals of gold from
the Berezoffsky mines in the Urals contain 90 to 95 of gold and 5
to 9 of silver, and the Altai gold contains 50 to 65 of gold and
36 to 38 of silver. The proportion of silver in native gold varies
between these limits in other localities. Silver ores, which
generally occur in veins, usually contain native silver and
various sulphur compounds. The most famous mines in Europe are in
Saxony (Freiberg), which has a yearly output of as much as 26 tons
of silver, Hungary, and Bohemia (41 tons). In Russia, silver is
extracted in the Altai and at Nerchinsk (17 tons). The richest
silver mines known are in America, especially in Chili (as much as
70 tons), Mexico (200 tons), and more particularly in the Western
States of North America. The richness of these mines may be judged
from the fact that one mine in the State of Nevada (Comstock, near
Washoe and the cities of Gold Hill and Virginia), which was
discovered in 1859, gave an output of 400 tons in 1866. In place
of cupellation, chlorination may also be employed for extracting
silver from its ores. The method of chlorination consists in
converting the silver in an ore into silver chloride. This is
either done by a wet or by a dry method, roasting the ore with
NaCl. When the silver chloride is formed, the extraction of the
metal is also done by two methods. The first consists in the
silver chloride being reduced to metal by means of iron in
rotating barrels, with the subsequent addition of mercury which
dissolves the silver, but does not act on the other metals. The
mercury holding the silver in solution is distilled, when the
silver remains behind. This method is called _amalgamation_. The
other method is less frequently used, and consists in dissolving
the silver chloride in sodium chloride or in sodium thiosulphate,
and then precipitating the silver from the solution. The
amalgamation is then carried on in rotating barrels containing the
roasted ore mixed with water, iron, and mercury. The iron reduces
the silver chloride by taking up the chlorine from it. The
technical details of these processes are described in works on
metallurgy. The extraction of AgCl by the wet method is carried on
(Patera's process) by means of a solution of hyposulphite of
sodium which dissolves AgCl (_see_ Note 23), or by lixiviating
with a 2 p.c. solution of a double hyposulphite of Na and Cu
(obtained by adding CuSO_{4} to Na_{2}S_{2}O_{3}). The resultant
solution of AgCl is first treated with soda to precipitate
PbCO_{3}, and then with Na_{2}S, which precipitates the Ag and Au.
The process should be carried on rapidly to prevent the
precipitation of Cu_{2}S from the solution of CuSO_{4} and
Na_{2}S_{2}O_{3}.
Commercial silver generally contains copper, and, more rarely, other
metallic impurities also. Chemically _pure silver_ is obtained either by
cupellation or by subjecting ordinary silver to the following treatment.
The silver is first dissolved in nitric acid, which converts it and the
copper into nitrates, Cu(NO_{3})_{2} and AgNO_{3}; hydrochloric acid is
then added to the resultant solution (green, owing to the presence of the
cupric salt), which is considerably diluted with water in order to retain
the lead chloride in solution if the silver contained lead. The copper
and many other metals remain in solution, whilst the silver is
precipitated as silver chloride. The precipitate is allowed to settle,
and the liquid is decanted off; the precipitate is then washed and fused
with sodium carbonate. A double decomposition then takes place, sodium
chloride and silver carbonate being formed; but the latter decomposes
into metallic silver, because the silver oxide is decomposed by heat:
Ag_{2}CO_{3} = Ag_{2} + O + CO_{2}. The silver chloride may also be mixed
with metallic zinc, sulphuric acid, and water, and left for some time,
when the zinc removes the chlorine from the silver chloride and
precipitates the silver as a powder. This finely-divided silver is called
'molecular silver.'[15]
[15] There is another practical method which is also suitable for
separating the silver from the solutions obtained in photography,
and consists in precipitating the silver by oxalic acid. In this
case the amount of silver in the solution must be known, and 23
grams of oxalic acid dissolved in 400 grams of water must be added
for every 60 grams of silver in solution in a litre of water. A
precipitate of silver oxalate, Ag_{2}C_{2}O_{4}, is then obtained,
which is insoluble in water but soluble in acids. Hence, if the
liquid contain any free acid it must be previously freed from it
by the addition of sodium carbonate. The resultant precipitate of
silver oxalate is dried, mixed with an equal weight of dry sodium
carbonate, and thrown into a gently-heated crucible. The
separation of the silver then proceeds without an explosion,
whilst the silver oxalate if heated alone decomposes with
explosion.
According to Stas, the best method for obtaining silver from its
solutions is by the reduction of silver chloride dissolved in
ammonia by means of an ammoniacal solution of cuprous
thiosulphate; the silver is then precipitated in a crystalline
form. A solution of ammonium sulphite may be used instead of the
cuprous salt.
Chemically-pure silver has an exceeding pure white colour, and a specific
gravity of 10·5. Solid silver is lighter than the molten metal, and
therefore a piece of silver floats on the latter. The fusing-point of
silver is about 950° C., and at the high temperature attained by the
combustion of detonating gas it volatilises.[16] By employing silver
reduced from silver chloride by milk sugar and caustic potash, and
distilling it, Stas obtained silver purer than that obtained by any other
means; in fact, this was perfectly pure silver. The vapour of silver has
a very beautiful green colour, which is seen when a silver wire is placed
in an oxyhydrogen flame.[17]
[16] Silver is very malleable and ductile; it may be beaten into leaves
0·002 mm. in thickness. Silver wire may be made so fine that 1
gram is drawn into a wire 2-1/2 kilometres long. In this respect
silver is second only to gold. A wire of 2 mm. diameter breaks
under a strain of 20 kilograms.
[17] In melting, silver absorbs a considerable amount of oxygen, which
is disengaged on solidifying. One volume of molten silver absorbs
as much as 22 volumes of oxygen. In solidifying, the silver forms
cavities like the craters of a volcano, and throws off metal,
owing to the evolution of the gas; all these phenomena recall a
volcano on a miniature scale (Dumas). Silver which contains a
small quantity of copper or gold, &c., does not show this property
of dissolving oxygen.
The absorption of oxygen by molten silver is, however, an
oxidation, but it is at the same time a phenomenon of solution.
One cubic centimetre of molten silver can dissolve twenty-two
cubic centimetres of oxygen, which, even at 0°, only weighs 0·03
gram, whilst 1 cubic centimetre of silver weighs at least 10
grams, and therefore it is impossible to suppose that the
absorption of the oxygen is attended by the formation of any
definite compound (rich in oxygen) of silver and oxygen (about 45
atoms of silver to 1 of oxygen) in any other but a dissociated
form, and this is the state in which substances in solution must
be regarded (Chapter I.)
Le Chatelier showed that at 300° and 15 atmospheres pressure
silver absorbs so much oxygen that it may be regarded as having
formed the compound Ag_{4}O, or a mixture of Ag_{2} and Ag_{2}O.
Moreover, silver oxide, Ag_{2}O, only decomposes at 300° under low
pressures, whilst at pressures above 10 atmospheres there is no
decomposition at 300° but only at 400°.
Stas showed that silver is oxidised by air in the presence of
acids. V. d. Pfordten confirmed this, and showed that an acidified
solution of potassium permanganate rapidly dissolves silver in the
presence of air.
It has long been known (Wöhler) that when nitrate of silver, AgNO_{3},
reacts as an oxidising agent upon citrates and tartrates, it is able
under certain conditions to give either a salt of suboxide of silver (see
Note 19) or a red solution, or to give a precipitate of metallic silver
reduced at the expense of the organic substances. In 1889 Carey Lea, in
his researches on this class of reactions, showed that _soluble silver_
is here formed, which he called _allotropic silver_. It may be obtained
by taking 200 c.c. of a 10 per cent. solution of AgNO_{3} and quickly
adding a mixture (neutralised with NaHO) of 200 c.c. of a 30 per cent.
solution of FeSO_{4} and 200 c.c. of a 40 per cent. solution of sodium
citrate. A lilac precipitate is obtained, which is collected on a filter
(the precipitate becomes blue) and washed with a solution of
NH_{4}NO_{3}. It then becomes soluble in pure water, forming a red
perfectly transparent solution from which the dissolved silver is
precipitated on the addition of many soluble foreign bodies. Some of the
latter--for instance, NH_{4}NO_{3}, alkaline sulphates, nitrates, and
citrates--give a precipitate which redissolves in pure water, whilst
others--for instance, MgSO_{4}, FeSO_{4}, K_{2}Cr_{2}O_{7}, AgNO_{3},
Ba(NO_{3})_{2} and many others--convert the precipitated silver into a
new variety, which, although no longer soluble in water, regains its
solubility in a solution of borax and is soluble in ammonia. Both the
soluble and insoluble silver are rapidly converted into the ordinary
grey-metallic variety by sulphuric acid, although nothing is given off in
the reaction; the same change takes place on ignition, but in this case
CO_{2} is disengaged; the latter is formed from the organic substances
which remain (to the amount of 3 per cent.) in the modified silver (they
are not removed by soaking in alcohol or water). If the precipitated
silver be slightly washed and laid in a smooth thin layer on paper or
glass, it is seen that the soluble variety is red when moist and a fine
blue colour when dry, whilst the insoluble variety has a blue reflex.
Besides these, under special conditions[18] a golden yellow variety may
be obtained, which gives a brilliant golden yellow coating on glass; but
it is easily converted into the ordinary grey-metallic state by friction
or trituration. There is no doubt[18 bis] that there is the same relation
between ordinary silver which is perfectly insoluble in water and the
varieties of silver obtained by Carey Lea[18 tri] as there is between
quartz and soluble silica or between CuS and As_{2}S_{2} in their
ordinary insoluble forms and in the state of the colloid solution of
their hydrosols (_see_ Chapter I., Note 57, and Chapter XVII., Note 25
bis). Here, however, an important step in advance has been made in this
respect, that we are dealing with the solution of a simple body, and
moreover of a metal--_i.e._ of a particularly characteristic state of
matter. And as boron, gold, and certain other simple bodies have already
been obtained in a soluble (colloid) form, and as numerous organic
compounds (albuminous substances, gum, cellulose, starch, &c.) and
inorganic substances are also known in this form, it might be said that
the colloid state (of hydrogels and hydrosols) can be acquired, if not by
every substance, at all events by substances of most varied chemical
character under particular conditions of formation from solutions. And
this being the case, we may hope that a further study of soluble colloid
compounds, which apparently present various transitions towards
emulsions, may throw a new light upon the complex question of solutions,
which forms one of the problems of the present epoch of chemical science.
Moreover, we may remark that Spring (1890) clearly proved the colloid
state of soluble silver by means of dialysis as it did not pass through
the membrane.
[18] When solutions of AgNO_{3}, FeSO_{4}, sodium citrate, and NaHO are
mixed together in the manner described above, they throw down a
precipitate of a beautiful lilac colour; when transferred to a
filter paper the precipitate soon changes colour, and becomes dark
blue. To obtain the substance as pure as possible it is washed
with a 5-10 p.c. solution of ammonium nitrate; the liquid is
decanted, and 150 c.c. of water poured over the precipitate. It
then dissolves entirely in the water. A small quantity of a
saturated solution of ammonium nitrate is added to the solution,
and the silver in solution again separates out as a precipitate.
These alternate solutions and precipitations are repeated seven or
eight times, after which the precipitate is transferred to a
filter and washed with 95 p.c. alcohol until the filtrate gives no
residue on evaporation. An analysis of the substance so obtained
showed that it contained from 97·18 p.c. to 97·31 p.c. of metallic
silver. It remained to discover what the remaining 2-3 p.c. were
composed of. Are they merely impurities, or is the substance some
compound of silver with oxygen or hydrogen, or does it contain
citric acid in combination which might account for its solubility?
The first supposition is set aside by the fact that no gases are
disengaged by the precipitate of silver, either under the action
of gases or when heated. The second supposition is shown to be
impossible by the fact that there is no definite relation between
the silver and citric acid. A determination of the amount of
silver in solution showed that the amount of citric acid varies
greatly for one and the same amount of silver, and there is no
simple ratio between them. Among other methods of preparing
soluble silver given by Carey Lea, we may mention the method
published by him in 1891. AgNO_{3} is added to a solution of
dextrine in caustic soda or potash; at first a precipitate of
brown oxide of silver is thrown down, but the brown colour then
changes into a reddish chocolate, owing to the reduction of the
silver by the dextrine, and the solution turns a deep red. A few
drops of this solution turn water bright red, and give a perfectly
transparent liquid. The dextrine solution is prepared by
dissolving 40 grams of caustic soda and the same amount of
ordinary brown dextrine in two litres of water. To this solution
is gradually added 28 grams of AgNO_{3} dissolved in a small
quantity of water.
The insoluble allotropic silver is obtained, as was mentioned
above, from a solution of silver prepared in the manner described,
by the addition of sulphate of copper, iron, barium, magnesium,
&c. In one experiment Lea succeeded in obtaining the insoluble
allotropic Ag in a crystalline form. The red solution, described
above, after standing several weeks, deposits crystals
spontaneously in the form of short black needles and thin prisms,
the liquid becoming colourless. This insoluble variety, when
rubbed upon paper, has the appearance of bright shining green
flakes, which polarise light.
The gold variety is obtained in a different manner to the two
other varieties. A solution is prepared containing 200 c.c. of a
10 p.c. solution of nitrate of silver, 200 c.c. of a 20 p.c.
solution of Rochelle salt, and 800 c.c. of water. Just as in the
previous case the reaction consisted in the reduction of the
citrate of silver, so in this case it consists in the reduction of
the tartrate, which here first forms a red, and then a black
precipitate of allotropic Ag, which, when transferred to the
filter, appears of a beautiful bronze colour. After washing and
drying, this precipitate acquires the lustre and colour peculiar
to polished gold, and this is especially remarked where the
precipitate comes into contact with glass or china. An analysis of
the golden variety gave a percentage composition of 98·750 to
98·749 Ag. Both the insoluble varieties (the blue and gold) have a
different specific gravity from ordinary silver. Whilst that of
fused silver is 10·50, and of finely-divided silver 10·62, the
specific gravity of the blue insoluble variety is 9·58, and of the
gold variety 8·51. The gold variety passes into ordinary Ag with
great ease. This transition may even be remarked on the filter in
those places which have accidentally not been moistened with
water. A simple shock, and therefore friction of one particle upon
another, is enough to convert the gold variety into normal white
silver. Carey Lea sent samples of the gold variety for a long
distance by rail packed in three tubes, in which the silver
occupied about the quarter of their volume; in one tube only he
filled up this space with cotton-wool. It was afterwards found
that the shaking of the particles of Ag had completely converted
it into ordinary white silver, and that only the tube containing
the cotton-wool had preserved the golden variety intact.
The soluble variety of Ag also passes into the ordinary state with
great ease, the heat of conversion being, as Prange showed in
1890, about +60 calories.
[18 bis] The opinion of the nature of soluble silver given below was
first enunciated in the _Journal of the Russian Chemical Society_,
February 1, 1890, Vol. XXII., Note 73. This view is, at the
present time, generally accepted, and this silver is frequently
known as the 'colloid' variety. I may add that Carey Lea observed
the solution of ordinary molecular silver in ammonia without the
access of air.
[18 tri] It is, however, noteworthy that ordinary metallic lead has
long been considered soluble in water, that boron has been
repeatedly obtained in a brown solution, and that observations
upon the development of certain bacteria have shown that the
latter die in water which has been for some time in contact with
metals. This seems to indicate the passage of small quantities of
metals into water (however, the formation of peroxide of hydrogen
may be supposed to have some influence in these cases).
As regards the capacity of silver for chemical reactions, it is
remarkable for its small capacity for combination with oxygen and for its
considerable energy of combination with sulphur, iodine, and certain
kindred non-metals. _Silver does not oxidise_ at any temperature, and its
oxide, Ag_{2}O, is decomposed by heat. It is also a very important fact
that silver is not oxidised by oxygen either in the presence of alkalis,
even at exceedingly high temperatures, or in the presence of acids--at
least, of dilute acids--which properties render it a very important metal
in chemical industry for the fusion of alkalis, and also for many
purposes in everyday life; for instance, for making spoons, salt-cellars,
&c. Ozone, however, oxidises it. Of all acids nitric acid has the
greatest action on silver. The reaction is accompanied by the formation
of oxides of nitrogen and silver nitrate, AgNO_{3}, which dissolves in
water and does not, therefore, hinder the further action of the acid on
the metal. The halogen acids, especially hydriodic acid, act on silver,
hydrogen being evolved; but this action soon stops, owing to the halogen
compounds of silver being insoluble in water and only very slightly
soluble in acids; they therefore preserve the remaining mass of metal
from the further action of the acid; in consequence of this the action of
the halogen acids is only distinctly seen with finely-divided silver.
Sulphuric acid acts on silver in the same manner that it does on copper,
only it must be concentrated and at a higher temperature. Sulphurous
anhydride, and not hydrogen, is then evolved, but there is no action at
the ordinary temperature, even in the presence of air. Among the various
salts, sodium chloride (in the presence of moisture, air, and carbonic
acid) and potassium cyanide (in the presence of air) act on silver more
decidedly than any others, converting it respectively into silver
chloride and a double cyanide.
Although silver does not directly combine with oxygen, still three
different grades of combination with oxygen may be obtained indirectly
from the salts of silver. They are all, however, unstable, and decompose
into oxygen and metallic silver when ignited. These three oxides of
silver have the following composition: _silver suboxide_, Ag_{4}O,[19]
corresponding with the (little investigated) suboxides of the alkali
metals; _silver oxide_, Ag_{2}O, corresponding with the oxides of the
alkali metals and the ordinary salts of silver, AgX; and _silver
peroxide_, AgO,[19 bis] or, judging from Berthelot's researches,
Ag_{2}O_{3}. _Silver oxide_ is obtained as a brown precipitate (which
when dried does not contain water) by adding potassium hydroxide to a
solution of a silver salt--for example, of silver nitrate. The
precipitate formed seems, however, to be an hydroxide, AgHO, _i.e._
AgNO_{3} + KHO = KNO_{3} + AgHO, and the formation of the anhydrous
oxide, 2AgHO = Ag_{2}O + H_{2}O, may be compared with the formation of
the anhydrous cupric oxide by the action of potassium hydroxide on hot
cupric solutions. Silver hydroxide decomposes into water and silver
oxide, even at low temperatures; at least, the hydroxide no longer exists
at 60°, but forms the anhydrous oxide, Ag_{2}O.[19 tri] Silver oxide is
almost insoluble in water; but, nevertheless, it is undoubtedly a rather
powerful basic oxide, because it displaces the oxides of many metals from
their soluble salts, and saturates such acids as nitric acid, forming
with them neutral salts, which do not act on litmus paper.[20]
Undoubtedly water dissolves a small quantity of silver oxide, which
explains the possibility of its action on solutions of salts--for
example, on solutions of cupric salts. Water in which silver oxide is
shaken up has a distinctly alkaline reaction. The oxide is distinguished
by its great instability when heated, so that it loses all its oxygen
when slightly heated. Hydrogen reduces it at about 80°.[20 bis] The
feebleness of the affinity of silver for oxygen is shown by the fact that
silver oxide decomposes under the action of light, so that it must be
kept in opaque vessels. The silver _salts_ are colourless and decompose
when heated, leaving metallic silver if the elements of the acid are
volatile.[20 tri] They have a peculiar metallic taste, and are
exceedingly poisonous; the majority of them are acted on by light,
especially in the presence of organic substances, which are then
oxidised. The alkaline carbonates give a white precipitate of silver
carbonate, Ag_{2}CO_{3}, which is insoluble in water, but soluble in
ammonia and ammonium carbonate. Aqueous ammonia, added to a solution of a
normal silver salt, first acts like potassium hydroxide, but the
precipitate dissolves in an excess of the reagent, like the precipitate
of cupric hydroxide.[21] Silver oxalate and the halogen compounds of
silver are insoluble in water; hydrochloric acid and soluble chlorides
give, as already repeatedly observed, a white precipitate of silver
chloride in solutions of silver salts. Potassium iodide gives a yellowish
precipitate of silver iodide. Zinc separates all the silver in a metallic
form from solutions of silver salts. Many other metals and reducing
agents--for example, organic substances--also reduce silver from the
solutions of its salts.
[19] Silver suboxide (Ag_{4}O) or argentous oxide is obtained from
argentic citrate by heating it to 100° in a stream of hydrogen.
Water and argentous citrate are then formed, and the latter,
although but slightly soluble in water, gives a reddish-brown
solution of colloid silver (Note 18), and when boiled this
solution becomes colourless and deposits metallic silver, the
argentic salt being again formed. Wöhler, who discovered this
oxide, obtained it as a black precipitate by adding potassium
hydroxide to the above solution of argentous citrate. With
hydrochloric acid the suboxide gives a brown compound, Ag_{2}Cl.
Since the discovery of soluble silver the above data cannot be
regarded as perfectly trustworthy; it is probable that a mixture
of Ag_{2} and Ag_{2}O was being dealt with, so that the actual
existence of Ag_{4}O is now doubtful, but there can be no doubt as
to the formation of a subchloride, Ag_{2}Cl (_see_ Note 25),
corresponding to the suboxide. The same compound is obtained by
the action of light on the higher chloride. Other acids do not
combine with silver suboxide, but convert it into an argentic salt
and metallic silver. In this respect cuprous oxide presents a
certain resemblance to these suboxides. But copper forms a
suboxide of the composition Cu_{4}O, which is obtained by the
action of an alkaline solution of stannous oxide on cupric
hydroxide, and is decomposed by acids into cupric salts and
metallic copper. The problems offered by the suboxides, as well as
by the peroxides, cannot be considered as fully solved.
[19 bis] _Silver peroxide_, AgO or Ag_{2}O_{3}, is obtained by the
decomposition of a dilute (10 p.c.) solution of silver nitrate by
the action of a galvanic current (Ritter). On the positive pole,
where oxygen is usually evolved in the decomposition of salts,
brittle grey needles with a metallic lustre, which occasionally
attain a somewhat considerable size, are then formed. They are
insoluble in water, and decompose with the evolution of oxygen
when they are dried and heated, especially up to 150°, and, like
lead dioxide, barium peroxide, &c., their action is strongly
oxidising. When treated with acids, oxygen is evolved and a salt
of the oxide formed. Silver peroxide absorbs sulphurous anhydride
and forms silver sulphate. Hydrochloric acid evolves chlorine;
ammonia reduces the silver, and is itself oxidised, forming water
and gaseous nitrogen. Analyses of the above-mentioned crystals
show that they contain silver nitrate, peroxide, and water.
According to Fisher, they have the composition
4AgO,AgNO_{3},H_{2}O, and, according to Berthelot,
4Ag_{2}O_{5},2AgNO_{3},H_{2}O.
[19 tri] According to Carey Lea, however, oxide of silver still retains
water even at 100°, and only parts with it together with the
oxygen. Oxide of silver is used for colouring glass yellow.
[20] The reaction of Pb(OH)_{2} upon AgHO in the presence of NaHO leads
to the formation of a compound of both oxides, PbO_n_Ag_{2}O, from
which the oxide of lead cannot be removed by alkalies (Wöhler,
Leton). Wöhler, Welch, and others obtained crystalline double
salts, R_{2}AgX_{3}, by the action of strong solutions of RX of
the halogen salts of the alkaline metals upon AgX, where R = Cs,
Rb, K.
[20 bis] According to Müller, ferric oxide is reduced by hydrogen
(_see_ Chapter XXII., Note 5) at 295° (into what ?), cupric oxide
at 140°, Ni_{2}O_{3} at 150°; nickelous oxide, NiO, is reduced to
the suboxide, Ni_{2}O, at 195°, and to nickel at 270°; zinc oxide
requires so high a temperature for its reduction that the glass
tube in which Müller conducted the experiment did not stand the
heat; antimony oxide requires a temperature of 215° for its
reduction; yellow mercuric oxide is reduced at 130° and the red
oxide at 230°; silver oxide at 85°, and platinum oxide even at the
ordinary temperature.
[20 tri] A silica compound, Ag_{2}OSiO_{2} is obtained by fusing
AgNO_{3} with silica; this salt is able to decompose with the
evolution of oxygen, leaving Ag + SiO_{2}.
[21] If a solution of a silver salt be precipitated by sodium
hydroxide, and aqueous ammonia is added drop by drop until the
precipitate is completely dissolved, the liquid when evaporated
deposits a violet mass of crystalline silver oxide. If moist
silver oxide be left in a strong solution of ammonia it gives a
black mass, which easily decomposes with a loud explosion,
especially when struck. This black substance is called fulminating
silver. Probably this is a compound like the other compounds of
oxides with ammonia, and in exploding the oxygen of the silver
oxide forms water with the hydrogen of the ammonia, which is
naturally accompanied by the evolution of heat and formation of
gaseous nitrogen, or, as Raschig states, fulminating silver
contains NAg_{3} or one of the amides (for instance, NHAg_{2} =
NH_{3} + Ag_{2}O - H_{2}O). Fulminating silver is also formed when
potassium hydroxide is added to a solution of silver nitrate in
ammonia. The dangerous explosions which are produced by this
compound render it needful that great care be taken when salts of
silver come into contact with ammonia and alkalis (_see_ Chapter
XVI., Note 26).
_Silver nitrate_, AgNO_{3}, is known by the name of lunar caustic (or
_lapis infernalis_); it is obtained by dissolving metallic silver in
nitric acid. If the silver be impure, the resultant solution will contain
a mixture of the nitrates of copper and silver. If this mixture be
evaporated to dryness and the residue carefully fused at an incipient red
heat, all the cupric nitrate is decomposed, whilst the greater part of
the silver nitrate remains unchanged. On treating the fused mass with
water the latter is dissolved, whilst the cupric oxide remains insoluble.
If a certain amount of silver oxide be added to the solution containing
the nitrates of silver and copper, it displaces all the cupric oxide. In
this case it is of course not necessary to take pure silver oxide, but
only to pour off some of the solution and to add potassium hydroxide to
one portion, and to mix the resultant precipitate of the hydroxides,
Cu(OH)_{2} and AgOH, with the remaining portion.[22] By these methods all
the copper can be easily removed and pure silver nitrate obtained (its
solution is colourless, while the presence of Cu renders it blue), which
may be ultimately purified by crystallisation. It crystallises in
colourless transparent prismatic plates, which are not acted on by air.
They are anhydrous. Its sp. gr. is 4·34; it dissolves in half its weight
of water at the ordinary temperature.[22 bis] The pure salt is not acted
on by light, but it easily acts in an oxidising manner on the majority of
organic substances, which it generally blackens. This is due to the fact
that the organic substance is oxidised by the silver nitrate, which is
reduced to metallic silver; the latter is thus obtained in a
finely-divided state, which causes the black stain. This peculiarity is
taken advantage of for marking linen. Silver nitrate is for the same
reason used for _cauterising wounds_ and various growths on the body.
Here again it acts by virtue of its oxidising capacity in destroying the
organic matter, which it oxidises, as is seen from the separation of a
coating of black metallic powdery silver from the part cauterised.[22
tri] From the description of the preparation of silver nitrate it will
have been seen that this salt, which fuses at 218°, does not decompose at
an incipient red heat; when cast into sticks it is usually employed for
cauterising. On further heating, the fused salt undergoes decomposition,
first forming silver nitrite and then metallic silver. With ammonia,
silver nitrate forms, on evaporation of the solution, colourless crystals
containing AgNO_{3},2HN_{3} (Marignac). In general the salts of silver,
like cuprous, cupric, zinc, &c. salts, are able to give several compounds
with ammonia; for example, silver nitrate in a dry state absorbs three
molecules (Rose). The ammonia is generally easily expelled from these
compounds by the action of heat.
[22] So that we here encounter the following phenomena: copper
displaces silver from the solutions of its salts, and silver oxide
displaces copper oxide from cupric salts. Guided by the
conceptions enunciated in Chapter XV., we can account for this in
the following manner: The atomic volume of silver = 10·3, and of
copper = 7·2, of silver oxide = 32, and of copper oxide = 13. A
greater contraction has taken place in the formation of cupric
oxide, CuO, than in the formation of silver oxide, Ag_{2}O, since
in the former (13 - 7 = 6) the volume after combination with the
oxygen has increased by very little, whilst the volume of silver
oxide is considerably greater than that of the metal it contains
[32 - (2 × 10·3) = 11·4]. Hence silver oxide is less compact than
cupric oxide, and is therefore less stable; but, on the other
hand, there are greater intervals between the atoms in silver
oxide than in cupric oxide, and therefore silver oxide is able to
give more stable compounds than those of copper oxide. This is
verified by the figures and data of their reactions. It is
impossible to calculate for cupric nitrate, because this salt has
not yet been obtained in an anhydrous state; but the sulphates of
both oxides are known. The specific gravity of copper sulphate in
an anhydrous state is 3·53, and of silver sulphate 5·36; the
molecular volume of the former is 45, and of the latter 58. The
group SO_{3} in the copper occupies, as it were, a volume 45 - 13
= 32, and in the silver salt a volume 58 - 32 = 26; hence a
smaller contraction has taken place in the formation of the copper
salt from the oxide than in the formation of the silver salt, and
consequently the latter should be more stable than the former.
Hence silver oxide is able to decompose the salt of copper oxide,
whilst with respect to the metals both salts have been formed with
an almost identical contraction, since 58 volumes of the silver
salt contain 21 volumes of metal (difference = 37), and 45 volumes
of the copper salt contain 7 volumes of copper (difference = 38).
Besides which, it must be observed that copper oxide displaces
iron oxide, just as silver oxide displaces copper oxide. Silver,
copper, and iron, in the form of oxides, displace each other in
the above order, but in the form of metals in a reverse order
(iron, copper, silver). The cause of this order of the
displacement of the oxides lies, amongst other things, in their
composition. They have the composition Ag_{2}O, Cu_{2}O_{2},
Fe_{2}O_{3}; the oxide containing a less proportion of oxygen
displaces that containing a larger proportion, because the basic
character diminishes with the increase of contained oxygen.
Copper also displaces mercury from its salts. It may here be
remarked that Spring (1888), on leaving a mixture of dry mercurous
chloride and copper for two hours, observed a distinct reduction,
which belongs to the category of those phenomena which demonstrate
the existence of a mobility of parts (_i.e._ atoms and molecules)
in solid substances.
[22 bis] The reaction of 1 part by weight of AgNO_{3} requires
(according to Kremers) the following amounts of water: at 0°, 0·82
part, at 19°·5, 0·41 part, at 54°, 0·20 part, at 110°, 0·09 part,
and, according to Tilden, at 125°, 0·0617 part, and at 133°,
0·0515 part.
[22 tri] It may be remarked that the black stain produced by the
reduction of metallic silver disappears under the action of a
solution of mercuric chloride or of potassium cyanide, because
these salts act on finely-divided silver.
Nitrate of silver easily forms double salts like AgNO_{3}2NaNO_{3} and
AgNO_{3}KNO_{3}. Silver nitrate under the action of water and a halogen
gives nitric acid (_see_ Vol. I. p. 280, formation of N_{2}O_{5}), a
halogen salt of silver, and a silver salt of an oxygen acid of the
halogen. Thus, for example, a solution of chlorine in water, when mixed
with a solution of silver nitrate, gives silver chloride and chlorate. It
is here evident that the reaction of the silver nitrate is identical with
the reaction of the caustic alkalis, as the nitric acid is all set free
and the silver oxide only reacts in exactly the same way in which aqueous
potash acts on free chlorine. Hence the reaction may be expressed in the
following manner: 6AgNO_{3} + 3Cl_{2} + 3H_{2}O = 5AgCl + AgClO_{3} +
6NHO_{3}.
Silver nitrate, like the nitrates of the alkalis, does not contain any
water of crystallisation. Moreover the other salts of silver almost
always separate out without any water of crystallisation. The silver
salts are further characterised by the fact that they _give neither basic
nor acid salts_, owing to which the formation of silver salts generally
forms the means of determining the true composition of acids--thus, to
any acid H_{n}X there corresponds a salt Ag_{n}X--for instance,
Ag_{3}PO_{4} (Chapter XIX., Note 15).
_Silver_ gives insoluble and exceedingly stable _compounds with the
halogens_. They are obtained by double decomposition with great facility
whenever a silver salt comes in contact with halogen salts. Solutions of
nitrate, sulphate, and all other kindred salts of silver give a
precipitate of silver chloride or iodide in solutions of chlorides and
iodides and of the halogen acids, because the halogen salts of silver are
insoluble both in water[23] and in other acids. _Silver chloride_, AgCl,
is then obtained as a white flocculent precipitate, silver bromide forms
a yellowish precipitate, and silver iodide has a very distinct yellow
colour. These halogen compounds sometimes occur in nature; they are
formed by a dry method--by the action of halogen compounds on silver
compounds, especially under the influence of heat. Silver chloride easily
fuses at 451° on cooling from a molten state; it forms a somewhat soft
horn-like mass which can be cut with a knife and is known as _horn
silver_. It volatilises at a higher temperature. Its ammoniacal solution,
on the evaporation of the ammonia, deposits crystalline chloride of
silver, in octahedra. Bromide and iodide of silver also appear in forms
of the regular system, so that in this respect the halogen salts of
silver resemble the halogen salts of the alkali metals.[24]
[23] Silver chloride is almost perfectly insoluble in water, but is
somewhat soluble in water containing sodium chloride or
hydrochloric acid, or other chlorides, and many salts, in
solution. Thus at 100°, 100 parts of water saturated with sodium
chloride dissolve 0·4 part of silver chloride. Bromide and iodide
of silver are less soluble in this respect, as also in regard to
other solvents. It should be remarked that _silver chloride
dissolves in solutions of ammonia, potassium cyanide, and of
sodium thiosulphate_, Na_{2}S_{2}O_{3}. Silver bromide is almost
perfectly analogous to the chloride, but silver iodide is nearly
insoluble in a solution of ammonia. Silver chloride even absorbs
dry ammonia gas, forming very unstable ammoniacal compounds. When
heated, these compounds (Vol. I. p. 250, Note 8) evolve the
ammonia, as they also do under the action of all acids. Silver
chloride enters into double decomposition with potassium cyanide,
forming a soluble double cyanide, which we shall presently
describe; it also forms a soluble double salt, NaAgS_{2}O_{3},
with sodium thiosulphate.
Silver chloride offers different modifications in the structure of
its molecule, as is seen in the variations in the consistency of
the precipitate, and in the differences in the action of light
which partially decomposes AgCl (_see_ Note 25). Stas and Carey
Lea investigated this subject, which has a particular importance
in photography, because silver bromide also gives _photo-salts_.
There is still much to be discovered in this respect, since Abney
showed that perfectly dry AgCl placed in a vacuum in the dark is
not in the least acted upon when subsequently exposed to light.
[24] _Silver bromide_ and _iodide_ (which occur as the minerals bromite
and iodite) resemble the chloride in many respects, but the degree
of affinity of silver for iodine is greater than that for chlorine
and bromine, although less heat is evolved (_see_ Note 28 bis).
Deville deduced this fact from a number of experiments. Thus
silver chloride, when treated with hydriodic acid, evolves
hydrochloric acid, and forms silver iodide. Finely-divided silver
easily liberates hydrogen when treated with hydriodic acid; it
produces the same decomposition with hydrochloric acid, but in a
considerably less degree and only on the surface. The difference
between silver chloride and iodide is especially remarkable, since
the formation of the former is attended with a greater contraction
than that of the latter. The volume of AgCl = 26; of chlorine 27,
of silver 10, the sum = 37, hence a contraction has ensued; and in
the formation of silver iodide an expansion takes place, for the
volume of Ag is 10, of I 26, and of AgI 39 instead of 36 (density,
AgCl, 5·59; AgI, 5·67). The atoms of chlorine have united with the
atoms of silver without moving asunder, whilst the atoms of iodine
must have moved apart in combining with the silver. It is
otherwise with respect to the metal; the distance between its
atoms in the metal = 2·2, in silver chloride = 3·0, and in silver
iodide = 3·5; hence its atoms have moved asunder considerably in
both cases. It is also very remarkable, as Fizeau observed, that
the density of silver iodide increases with a rise of
temperature--that is, a contraction takes place when it is heated
and an expansion when it is cooled.
In order to explain the fact that in silver compounds the iodide
is more stable than the chloride and oxide, Professor N. N.
Beketoff, in his 'Researches on the Phenomena of Substitutions'
(Kharkoff, 1865), proposed the following original hypothesis,
which we will give in almost the words of the author:--In the case
of aluminium, the oxide, Al_{2}O_{3}, is more stable than the
chloride, Al_{2}Cl_{6}, and the iodide, Al_{2}I_{6}. In the oxide
the amount of the metal is to the amount of the element combined
with it as 54·8 (Al = 27·3) is to 48, or in the ratio 112 : 100;
for the chloride the ratio is = 25 : 100; for the iodide it = 7 :
100. In the case of silver the oxide (ratio = 1350 : 100) is less
stable than the chloride (ratio = 304 : 100), and the iodide
(ratio of the weight of metal to the weight of the halogen = 85 :
100) is the most stable. From these and similar examples it
follows that the most stable compounds are those in which the
weights of the combined substances are equal. This may be partly
explained by the attraction of similar molecules even after their
having passed into combination with others. This attraction is
proportional to the product of the acting masses. In silver oxide
the attraction of Ag_{2} for Ag_{2} = 216 × 216 = 46,656, and the
attraction of Ag_{2} for O = 216 × 16 = 3,456. The attraction of
like molecules thus counteracts the attraction of the unlike
molecules. The former naturally does not overcome the latter,
otherwise there would be a disruption, but it nevertheless
diminishes the stability. In the case of an equality or proximity
of the magnitude of the combining masses, the attraction of the
like parts will counteract the stability of the compound to the
least extent--in other words, with an inequality of the combined
masses, the molecules have an inclination to return to an
elementary state, to decompose, which does not exist to such an
extent where the combined masses are equal. There is, therefore, a
tendency for large masses to combine with large, and for small
masses to combine with small. Hence Ag_{2}O + 2KI gives K_{2}O +
2AgI. The influence of an equality of masses on the stability is
seen particularly clearly in the effect of a rise of temperature.
Argentic, mercuric, auric and other oxides composed of unequal
masses, are somewhat readily decomposed by heat, whilst the oxides
of the lighter metals (like water) are not so easily decomposed by
heat. Silver chloride and iodide approach the condition of
equality, and are not decomposed by heat. The most stable oxides
under the action of heat are those of magnesium, calcium, silicon,
and aluminium, since they also approach the condition of equality.
For the same reason hydriodic acid decomposes with greater
facility than hydrochloric acid. Chlorine does not act on magnesia
or alumina, but it acts on lime and silver oxide, &c. This is
partially explained by the fact that by considering heat as a mode
of motion, and knowing that the atomic heats of the free elements
are equal, it must be supposed that the amount of the motion of
atoms (their _vis viva_) is equal, and as it is equal to the
product of the mass (atomic weight) into the square of the
velocity, it follows that the greater the combining weight the
smaller will be the square of the velocity, and if the combining
weights be nearly equal, then the velocities also will be nearly
equal. Hence the greater the difference between the weights of the
combined atoms the greater will be the difference between their
velocities. The difference between the velocities will increase
with the temperature, and therefore the temperature of
decomposition will be the sooner attained the greater be the
original difference--that is, the greater the difference of the
weights of the combined substances. The nearer these weights are
to each other, the more analogous the motion of the unlike atoms,
and consequently, the more stable the resultant compound.
The instability of cupric chloride and nitric oxide, the absence
of compounds of fluorine with oxygen, whilst there are compounds
of oxygen with chlorine, the greater stability of the oxygen
compounds of iodine than those of chlorine, the stability of boron
nitride, and the instability of cyanogen, and a number of similar
instances, where, judging from the above argument, one would
expect (owing to the closeness of the atomic weights) a stability,
show that Beketoff's addition to the mechanical theory of chemical
phenomena is still far from sufficient for explaining the true
relations of affinities. Nevertheless, in his mode of explaining
the relative stabilities of compounds, we find an exceedingly
interesting treatment of questions of primary importance. Without
such efforts it would be impossible to generalise the complex data
of experimental knowledge.
_Fluoride of silver_, AgF, is obtained by dissolving Ag_{2}O or
Ag_{2}CO_{3} in hydrofluoric acid. It differs from the other
halogen salts of silver in being soluble in water (1 part of salt
in 0·55 of water). It crystallises from its solution in prisms,
AgFH_{2}O (Marignac), or AgF_{2}H_{2}O (Pfaundler), which lose
their water in vacuo. Güntz (1891), by electrolising a saturated
solution of Ag_{2}F, obtained _polyfluoride of silver_, Ag_{2}F,
which is decomposed by water into AgF + Ag. It is also formed by
the action of a strong solution of AgF upon finely-divided
(precipitated) silver.
Silver chloride may be decomposed, with the separation of silver oxide,
by heating it with a solution of an alkali, and if an organic substance
be added to the alkali the chloride can easily be reduced o metallic
silver, the silver oxide being reduced in the oxidation of the organic
substance. Iron, zinc, and many other metals reduce silver chloride in
the presence of water. Cuprous and mercurous chlorides and many organic
substances are also able to reduce the silver from chloride of silver.
This shows the rather easy decomposability of the halogen compounds of
silver. Silver iodide is much more stable in this respect than the
chloride. The same is also observed with respect to the _action of light_
upon moist AgCl. White silver chloride soon acquires a violet colour when
exposed to the action of light, and especially under the direct action of
the sun's rays. After being acted upon by light it is no longer entirely
soluble in ammonia, but leaves metallic silver undissolved, from which it
might be assumed that the action of light consisted in the decomposition
of the silver chloride into chlorine and metallic silver and in fact the
silver chloride becomes in time darker and darker. Silver bromide and
iodide are much more slowly acted on by light, and, according to certain
observations, when pure they are even quite unacted on; at least they do
not change in weight,[24 bis] so that if they are acted on by light, the
change they undergo must be one of a change in the structure of their
parts and not of decomposition, as it is in silver chloride. The silver
chloride under the action of light changes in weight, which indicates the
formation of a volatile product, and the deposition of metallic silver on
dissolving in ammonia shows the loss of chlorine. The change does
actually occur under the action of light, but the decomposition does not
go as far as into chlorine and silver, but only to the formation of a
subchloride of silver, Ag_{2}Cl, which is of a brown colour and is easily
decomposed into metallic silver and silver chloride, Ag_{2}Cl = AgCl +
Ag. This change of the chemical composition and structure of the halogen
salts of silver under the action of light forms the basis of
_photography_, because the halogen compounds of silver, after having been
exposed to light, give a precipitate of finely-divided silver, of a black
colour, when treated with reducing agents.[25]
[24 bis] The changes brought about by the action of light necessitate
distinguishing the photo-salts of silver.
[25] In photography these are called 'developers.' The most common
developers are: solutions of ferrous sulphate, pyrogallol, ferrous
oxalate, hydroxylamine, potassium sulphite, hydroquinone (the last
acts particularly well and is very convenient to use), &c. The
chemical processes of photography are of great practical and
theoretical interest; but it would be impossible in this work to
enter into this special branch of chemistry, which has as yet been
very little worked out from a theoretical point of view.
Nevertheless, we will pause to consider certain aspects of this
subject which are of a purely chemical interest, and especially
the facts concerning _subchloride of silver_, Ag_{2}Cl (_see_ Note
19), and the photo-salts (Note 23). There is no doubt that under
the action of light, AgCl becomes darker in colour, decreases in
weight, and probably forms a mixture of AgCl, Ag_{2}Cl, and Ag.
But the isolation of the subchloride has only been recently
accomplished by Güntz by means of the Ag_{2}F, discovered by him
(_see_ Note 24). Many chemists (and among them Hodgkinson) assumed
that an oxychloride of silver was formed by the decomposition of
AgCl under the action of light. Carey Lea's (1889) and A.
Richardson's (1891) experiments showed that the product formed
does not, however, contain any oxygen at all, and the change in
colour produced by the action of light upon AgCl is most probably
due to the formation of Ag_{2}Cl. This substance was isolated by
Güntz (1891) by passing HCl over crystals of Ag_{2}F. He also
obtained Ag_{2}I in a similar manner by passing HI, and Ag_{2}S by
passing H_{2}S over Ag_{2}F. Ag_{2}Cl is best prepared by the
action of phosphorus trichloride upon Ag_{2}F. At the temperature
of its formation Ag_{2}Cl has an easily changeable tint, with
shades of violet red to violet black. Under the action of light a
similar (isomeric) substance is obtained, which splits up into
AgCl + Ag when heated. With potassium cyanide Ag_{2}Cl gives Ag +
AgCN + KCl, whence it is possible to calculate the heat of
formation of Ag_{2}Cl; it = 29·7, whilst the heat of formation of
AgCl = 29·2--_i.e._ the reaction 2AgCl = Ag_{2}Cl + Cl corresponds
to an absorption of 28·7 major calories. If we admit the formation
of such a compound by the action of light, it is evident that the
energy of the light is consumed in the above reaction. Carey Lea
(1892) subjected AgCl, AgBr, and AgI to a pressure (of course in
the dark) of 3,000 atmospheres, and to trituration with water in a
mortar, and observed a change of colour indicating incipient
decomposition, which is facilitated under the action of light by
the molecular currents set up (Lermontoff, Egoroff). The change of
colour of the halogen salts of silver under the action of light,
and their faculty of subsequently giving a visible photographic
image under the action of 'developers,' must now be regarded as
connected with the decomposition of AgX, leading to the formation
of Ag_{2}X, and the different tinted photo-salts must be
considered as systems containing such Ag_{2}X's. Carey Lea
obtained photo-salts of this kind not only by the action of light
but also in many other ways, which we will enumerate to prove that
they contain the products of an incomplete combination of Ag with
the halogens, (for the salts Ag_{2}X must be regarded as such).
The photo-salts have been obtained (1) by the imperfect
chlorination of silver; (2) by the incomplete decomposition of
Ag_{2}O or Ag_{2}CO_{3} by alternately heating and treating with a
halogen acid; (3) by the action of nitric acid or Na_{2}S_{2}O_{3}
upon Ag_{2}Cl; (4) by mixing a solution of AgNO_{3} with the
hydrates of FeO, MnO and CrO, and precipitating by HCl; (5) by the
action of HCl upon the product obtained by the reduction of
citrate of silver in hydrogen (Note 19), and (6) by the action of
milk sugar upon AgNO_{3} together with soda and afterwards
acidulating with HCl. All these reactions should lead to the
formation of products of imperfect combination with the halogens
and give photo-salts of a similar diversity of colour to those
produced by the action of developers upon the halogen salts of
silver after exposure to light.
The insolubility of the halogen compounds of silver forms the basis of
many methods used in practical chemistry. Thus by means of this reaction
it is possible to obtain salts of other acids from a halogen salt of a
given metal, for instance, RCl_{2} + 2AgNO_{3} = R(NO_{3})_{2} + 2AgCl.
The formation of the halogen compounds of silver is very frequently used
in the investigation of organic substances; for example, if any product
of metalepsis containing iodine or chlorine be heated with a silver salt
or silver oxide, the silver combines with the halogen and gives a halogen
salt, whilst the elements previously combined with the silver replace the
halogen. For instance, ethylene dibromide, C_{2}H_{4}Br_{2}, is
transformed into ethylene diacetate, C_{2}H_{4}(C_{2}H_{3}O_{2})_{2}, and
silver bromide by heating it with silver acetate, 2C_{2}H_{3}O_{2}Ag. The
insolubility of the halogen compounds of silver is still more frequently
taken advantage of in determining the amount of silver and halogen in a
given solution. If it is required, for instance, to determine the
quantity of chlorine present in the form of a metallic chloride in a
given solution, a solution of silver nitrate is added to it so long as it
gives a precipitate. On _shaking or stirring_ the liquid, the silver
chloride easily settles in the form of heavy flakes. It is possible in
this way to precipitate the whole of the chlorine from a solution,
without adding an excess of silver nitrate, since it can be easily seen
whether the addition of a fresh quantity of silver nitrate produces a
precipitate in the clear liquid. In this manner it is possible to add to
a solution containing chlorine, as much silver as is required for its
entire precipitation, and to calculate the amount of chlorine previously
in solution from the amount of the solution of silver nitrate consumed,
if the quantity of silver nitrate in this solution has been previously
determined.[25 bis] The atomic proportions and preliminary experiments
with a pure salt--for example, with sodium chloride--will give the amount
of chlorine from the quantity of silver nitrate. Details of these methods
will be found in works on analytical chemistry.[25 tri]
[25 bis] In order to determine when the reaction is at an end, a few
drops of a solution of K_{2}CrO_{4} are added to the solution of
the chloride. Before all the chlorine is precipitated as AgCl, the
precipitate (after shaking) is white (since Ag_{2}CrO_{4} with
2RCl gives 2AgCl); but when all the chlorine is thrown down
Ag_{2}CrO_{4} is formed, which colours the precipitate
reddish-brown. In order to obtain accurate results the liquid
should be neutral to litmus.
[25 tri] _Silver cyanide_, AgCN, is closely analogous to the haloid
salts of silver. It is obtained, in similar manner to silver
chloride, by the addition of potassium cyanide to silver nitrate.
A white precipitate is then formed, which is almost insoluble in
boiling water. It is also, like silver chloride, insoluble in
dilute acids. However, it is dissolved when heated with nitric
acid, and both hydriodic and hydrochloric acids act on it,
converting it into silver chloride and iodide. Alkalis, however,
do not act on silver cyanide, although they act on the other
haloid salts of silver. Ammonia and solutions of the cyanides of
the alkali metals dissolve silver cyanide, as they do the
chloride. In the latter case double cyanides are formed--for
example, KAgC_{2}N_{2}. This salt is obtained in a crystalline
state on evaporating a solution of silver cyanide in potassium
cyanide. It is much more stable than silver cyanide itself. It has
a neutral reaction, does not change in the air, and does not smell
of hydrocyanic acid. Many acids, in acting on a solution of this
double salt, precipitate the insoluble silver cyanide. Metallic
silver dissolves in a solution of potassium cyanide in the
presence of air, with formation of the same double salt and
potassium hydroxide, and when silver chloride dissolves in
potassium cyanide it forms potassium chloride, besides the salt
KAgC_{2}N_{2}. This double salt of silver is used in silver
plating. For this purpose potassium cyanide is added to its
solution, as otherwise silver cyanide, and not metallic silver, is
deposited by the electric current. If two electrodes--one positive
(silver) and the other negative (copper)--be immersed in such a
solution, silver will be deposited upon the latter, and the silver
of the positive electrode will be dissolved by the liquid, which
will thus preserve the same amount of metal in solution as it
originally contained. If instead of the negative electrode a
copper object be taken, well cleaned from all dirt, the silver
will be deposited in an even coating; this, indeed, forms the mode
of _silver plating by the wet method_, which is most often used in
practice. A solution of one part of silver nitrate in 30 to 50
parts of water, and mixed with a sufficient quantity of a solution
of potassium cyanide to redissolve the precipitate of silver
cyanide formed, gives a dull coating of silver, but if twice as
much water be used the same mixture gives a bright coating.
Silver plating in the wet way has now replaced to a considerable
extent the old process of _dry silvering_, because this process,
which consists in dissolving silver in mercury and applying the
amalgam to the surface of the objects, and then vaporising the
mercury, offers the great disadvantage of the poisonous mercury
fumes. Besides these, there is another method of silver plating,
based on the direct displacement of silver from its salts by other
metals--for example, by copper. The copper reduces the silver from
its compounds, and the silver separated is deposited upon the
copper. Thus a solution of silver chloride in sodium thiosulphate
deposits a coating of silver upon a strip of copper immersed in
it. It is best for this purpose to take pure _silver sulphite_.
This is prepared by mixing a solution of silver nitrate with an
excess of ammonia, and adding a saturated solution of sodium
sulphite and then alcohol, which precipitates silver sulphite from
the solution. The latter and its solutions are very easily
decomposed by copper. Metallic iron produces the same
decomposition, and iron and steel articles may be very readily
silver-plated by means of the thiosulphate solution of silver
chloride. Indeed, copper and similar metals may even be
silver-plated by means of silver chloride; if the chloride of
silver, with a small amount of acid, be rubbed upon the surface of
the copper, the latter becomes covered with a coating of silver,
which it has reduced.
Silver plating is not only applicable to metallic objects, but
also to glass, china, &c. Glass is silvered for various
purposes--for example, glass globes silvered internally are used
for ornamentation, and have a mirrored surface. Common
looking-glass silvered upon one side forms a mirror which is
better than the ordinary mercury mirrors, owing to the truer
colours of the image due to the whiteness of the silver. For
optical instruments--for example, telescopes--concave mirrors are
now made of silvered glass, which has first been ground and
polished into the required form. The _silvering of glass_ is based
on the fact that silver which is reduced from certain solutions
deposits itself uniformly in a perfectly homogeneous and
continuous but very thin layer, forming a bright reflecting
surface. Certain organic substances have the property of reducing
silver in this form. The best known among these are certain
aldehydes--for instance, ordinary acetaldehyde, C_{2}H_{4}O, which
easily oxidises in the air and forms acetic acid, C_{2}H_{4}O_{2}.
This oxidation also easily takes place at the expense of silver
oxide, when a certain amount of ammonia is added to the mixture.
The oxide of silver gives up its oxygen to the aldehyde, and the
silver reduced from it is deposited in a metallic state in a
uniform bright coating. The same action is produced by certain
saccharine substances and certain organic acids, such as tartaric
acid, &c.
Accurate experiments, and more especially the _researches of Stas_ at
Brussels, show the proportion in which silver reacts with metallic
chlorides. These researches have led to the determination of the
_combining weights_ of silver, sodium, potassium, chlorine, bromine,
iodine, and other elements, and are distinguished for their model
exactitude, and we will therefore describe them in some detail. As sodium
chloride is the chloride most generally used for the precipitation of
silver, since it can most easily be obtained in a pure state, we will
here cite the quantitative observations made by Stas for showing the
co-relation between the quantities of chloride of sodium and silver which
react together. In order to obtain perfectly pure sodium chloride, he
took pure rock salt, containing only a small quantity of magnesium and
calcium compounds and a small amount of potassium salts. This salt was
dissolved in water, and the saturated solution evaporated by boiling. The
sodium chloride separated out during the boiling, and the mother liquor
containing the impurities was poured off. Alcohol of 65 p.c. strength and
platinic chloride were added to the resultant salt, in order to
precipitate all the potassium and a certain part of the sodium salts. The
resultant alcoholic solution, containing the sodium and platinum
chlorides, was then mixed with a solution of pure ammonium chloride in
order to remove the platinic chloride. After this precipitation, the
solution was evaporated in a platinum retort, and then separate portions
of this purified sodium chloride were collected as they crystallised. The
same salt was prepared from sodium sulphate, tartrate, nitrate, and from
the platinochloride, in order to have sodium chloride prepared by
different methods and from different sources, and in this manner ten
samples of sodium chloride thus prepared were purified and investigated
in their relation to silver. After being dried, weighed quantities of all
ten samples of sodium chloride were dissolved in water and mixed with a
solution in nitric acid of a weighed quantity of perfectly pure silver. A
slightly greater quantity of silver was taken than would be required for
the decomposition of the sodium chloride, and when, after pouring in all
the silver solution, the silver chloride had settled, the amount of
silver remaining in excess was determined by means of a solution of
sodium chloride of known strength. This solution of sodium chloride was
added so long as it formed a precipitate. In this manner Stas determined
how many parts of sodium chloride correspond to 100 parts by weight of
silver. The result of ten determinations was that for the entire
precipitation of 100 parts of silver, from 54·2060 to 54·2093 parts of
sodium chloride were required. The difference is so inconsiderable that
it has no perceptible influence on the subsequent calculations. The mean
of ten experiments was that 100 parts of silver react with 54·2078 parts
of sodium chloride. In order to learn from this the relation between the
chlorine and silver, it was necessary to determine the quantity of
chlorine contained in 54·2078 parts of sodium chloride, or, what is the
same thing, the quantity of chlorine which combines with 100 parts of
silver. For this purpose Stas made a series of observations on the
quantity of silver chloride obtained from 100 parts of silver. Four
syntheses were made by him for this purpose. The first synthesis
consisted in the formation of silver chloride by the action of chlorine
on silver at a red heat. This experiment showed that 100 parts of silver
give 132·841, 132·843 and 132·843 of silver chloride. The second method
consisted in dissolving a given quantity of silver in nitric acid and
precipitating it by means of gaseous hydrochloric acid passed over the
surface of the liquid; the resultant mass was evaporated in the dark to
drive off the nitric acid and excess of hydrochloric acid, and the
remaining silver chloride was fused first in an atmosphere of
hydrochloric acid gas and then in air. In this process the silver
chloride was not washed, and therefore there could be no loss from
solution. Two experiments made by this method showed that 100 parts of
silver give 132·849 and 132·846 parts of silver chloride. A third series
of determinations was also made by precipitating a solution of silver
nitrate with a certain excess of gaseous hydrochloric acid. The amount of
silver chloride obtained was altogether 132·848. Lastly, a fourth
determination was made by precipitating dissolved silver with a solution
of ammonium chloride, when it was found that a considerable amount of
silver (0·3175) had passed into solution in the washing; for 100 parts of
silver there was obtained altogether 132·8417 of silver chloride. Thus
from the mean of seven determinations it appears that 100 parts of silver
give 132·8445 parts of silver chloride--that is, that 32·8445 parts of
chlorine are able to combine with 100 parts of silver and with that
quantity of sodium which is contained in 54·2078 parts of sodium
chloride. These observations show that 32·8445 parts of chlorine combine
with 100 parts of silver and with 21·3633 parts of sodium. From these
figures expressing the relation between the combining weights of
chlorine, silver, and sodium, it would be possible to determine their
atomic weights--that is, the combining quantity of these elements with
respect to one part by weight of hydrogen or 16 parts of oxygen, if there
existed a series of similarly accurate determinations for the reactions
between hydrogen or oxygen and one of these elements--chlorine, sodium,
or silver. If we determine the quantity of silver chloride which is
obtained from silver chlorate, AgClO_{3}, we shall know the relation
between the combining weights of silver chloride and oxygen, so that,
taking the quantity of oxygen as a constant magnitude, we can learn from
this reaction the combining weight of silver chloride, and from the
preceding numbers the combining weights of chlorine and silver. For this
purpose it was first necessary to obtain pure silver chlorate. This Stas
did by acting on silver oxide or carbonate, suspended in water, with
gaseous chlorine.[26]
[26] The phenomenon which then takes place is described by Stas as
follows, in a manner which is perfect in its clearness and
accuracy: if silver oxide or carbonate be suspended in water, and
an excess of water saturated with chlorine be added, all the
silver is converted into chloride, just as is the case with oxide
or carbonate of mercury, and the water then contains, besides the
excess of chlorine, only pure hypochlorous acid without the least
trace of chloric or chlorous acid. If a stream of chlorine be
passed into water containing _an excess of silver oxide_ or silver
carbonate while the liquid is continually agitated, the reaction
is the same as the preceding; silver chloride and hypochlorous
acid are formed. But this acid does not long remain in a free
state: it gradually acts on the silver oxide and gives silver
hypochlorite, _i.e._ AgClO. If, after some time, the current of
chlorine be stopped but the shaking continued, the liquid loses
its characteristic odour of hypochlorous acid, while preserving
its energetic decolorising property, because the silver
hypochlorite which is formed is easily soluble in water. In the
presence of an excess of silver oxide this salt can be kept for
several days without decomposition, but it is exceedingly unstable
when no excess of silver oxide or carbonate is present. So long as
the solution of silver hypochlorite is shaken up with the silver
oxide, it preserves its transparency and bleaching property, but
directly it is allowed to stand, and the silver oxide settles, it
becomes rapidly cloudy and deposits large flakes of silver
chloride, so that the black silver oxide which had settled becomes
covered with the white precipitate. The liquid then loses its
bleaching properties and contains silver chlorate, _i.e._
AgClO_{3}, in solution, which has a slightly alkaline reaction,
owing to the presence of a small amount of dissolved oxide. In
this manner the reactions which are consecutively accomplished may
be expressed by the equations:
6Cl_{2} + 3Ag_{2}O + 3H_{2}O = 6AgCl + 6HClO;
6HClO + 3Ag_{2}O = 3H_{2}O + 6AgClO;
6AgClO = 4AgCl + 2AgClO_{3}.
Hence, Stas gives the following method for the preparation of
silver chlorate: A slow current of chlorine is caused to act on
oxide of silver, suspended in water which is kept in a state of
continual agitation. The shaking is continued after the supply of
chlorine has been stopped, in order that the free hypochlorous
acid should pass into silver hypochlorite, and the resultant
solution of the hypochlorite is drawn off from the sediment of the
excess of silver oxide. This solution decomposes spontaneously
into silver chloride and chlorate. The pure silver chlorate,
AgClO_{3}, does not change under the action of light. The salt is
prepared for further use by drying it in dry air at 150°. It is
necessary during drying to prevent the access of any organic
matter; this is done by filtering the air through cotton wool, and
passing it over a layer of red-hot copper oxide.
The decomposition of the silver chlorate thus obtained was accomplished
by the action of a solution of sulphurous anhydride on it. The salt was
first fused by carefully heating it at 243°. The solution of sulphurous
anhydride used was one saturated at 0°. Sulphurous anhydride in dilute
solutions is oxidised at the expense of silver chlorate, even at low
temperatures, with great ease if the liquid be continually shaken,
sulphuric acid and silver chloride being formed: AgClO_{3} + 3SO_{2} +
3H_{2}O = AgCl + 3H_{2}SO_{4}. After decomposition, the resultant liquid
was evaporated, and the residue of silver chloride weighed. Thus the
process consisted in taking a known weight of silver chlorate, converting
it into silver chloride, and determining the weight of the latter. The
analysis conducted in this manner gave the following results, which, like
the preceding, designate the weight in a vacuum calculated from the
weights obtained in air: In the first experiment it appeared that
138·7890 grams of silver chlorate gave 103·9795 parts of silver chloride,
and in the second experiment that 259·5287 grains of chlorate gave
194·44515 grams of silver chloride, and after fusion 194·4435 grams. The
mean result of both experiments, converted into percentages, shows that
100 parts of silver chlorate contain 74·9205 of silver chloride and
25·0795 parts of oxygen. From this it is possible to calculate the
combining weight of silver chloride, because in the decomposition of
silver chlorate there are obtained three atoms of oxygen and one molecule
of silver chloride: AgClO_{3} = AgCl + 3O. Taking the weight of an atom
of oxygen to be 16, we find from the mean result that the equivalent
weight of silver chloride is equal to 143·395. Thus if O = 16, AgCl =
143·395, and as the preceding experiments show that silver chloride
contains 32·8445 parts of chlorine per 100 parts of silver, the weight of
the atom of silver[26 bis] must be 107·94 and that of chlorine 35·45. The
weight of the atom of sodium is determined from the fact that 21·3633
parts of sodium chloride combine with 32·8445 parts of chlorine;
consequently Na = 23·05. This conclusion, arrived at by the analysis of
silver chlorate, was verified by means of the analysis of potassium
chlorate by decomposing it by heat and determining the weight of the
potassium chloride formed, and also by effecting the same decomposition
by igniting the chlorate in a stream of hydrochloric acid. The combining
weight of potassium chloride was thus determined, and another series of
determinations confirmed the relation between chlorine, potassium, and
silver, in the same manner as the relation between sodium, chlorine, and
silver was determined above. Consequently, the combining weights of
sodium, chlorine, and potassium could be deduced by combining these data
with the analysis of silver chlorate and the synthesis of silver
chloride. The agreement between the results showed that the
determinations made by the last method were perfectly correct, and did
not depend in any considerable degree on the methods which were employed
in the preceding determinations, as the combining weights of chlorine and
silver obtained were the same as before. There was naturally a
difference, but so small a one that it undoubtedly depended on the errors
incidental to every process of weighing and experiment. The atomic weight
of silver was also determined by Stas by means of the synthesis of silver
sulphide and the analysis of silver sulphate. The combining weight
obtained by this method was 107·920. The synthesis of silver iodide and
the analysis of silver iodate gave the figure 107·928. The synthesis of
silver bromide with the analysis of silver bromate gave the figure
107·921. The synthesis of silver chloride and the analysis of silver
chlorate gave a mean result of 107·937. Hence there is no doubt that the
combining weight of silver is at least as much as 107·9--greater than
107·90 and less than 107·95, and probably equal to the mean = 107·92.
Stas determined the combining weights of many other elements in this
manner, such as lithium, potassium, sodium, bromine, chlorine, iodine,
and also nitrogen, for the determination of the amount of silver nitrate
obtained from a given amount of silver gives directly the combining
weight of nitrogen. Taking that of oxygen as 16, he obtained the
following combining weights for these elements: nitrogen 14·04, silver
107·93, chlorine 35·46, bromine 79·95, iodine 126·85, lithium 7·02,
sodium 23·04, potassium 39·15. These figures differ slightly from those
which are usually employed in chemical investigations. They must be
regarded as the result of the best observations, whilst the figures
usually used in practical chemistry are only approximate--are, so to
speak, round numbers for the atomic weights which differ so little from
the exact figures (for instance, for Ag 108 instead of 107·92, for Na 23
instead of 23·04) that in ordinary determinations and calculations the
difference falls within the limits of experimental error inseparable from
such determinations.
[26 bis] The results given by Stas' determinations have recently
been recalculated and certain corrections have been introduced. We
give in the context the average results of van der Plaats and
Thomsen's calculations, as well as in Table III. neglecting the
doubtful thousandths.
The exhaustive investigations conducted by Stas on the atomic weights
of the above-named elements have great significance in the solution of
the problem as to whether the atomic weights of the elements can be
expressed in whole numbers if the unit taken be the atomic weight of
hydrogen. Prout, at the beginning of this century, stated that this was
the case, and held that the atomic weights of the elements are multiples
of the atomic weight of hydrogen. The subsequent determinations of
Berzelius, Penny, Marchand, Marignac, Dumas, and more especially of Stas,
proved this conclusion to be untenable; since a whole series of elements
proved to have fractional atomic weights--for example, chlorine, about
35·5. On account of this, Marignac and Dumas stated that the atomic
weights of the elements are expressed in relation to hydrogen, either by
whole numbers or by numbers with simple fractions of the magnitudes 1/2
and 1/4. But Stas's researches refute this supposition also. Even between
the combining weight of hydrogen and oxygen, there is not, so far as is
yet known, that simple relation which is required by _Prout's
hypothesis_,[27] _i.e._, taking O = 16, the atomic weight of hydrogen is
equal not to 1 but to a greater number somewhere between 1·002 and 1·008
or mean 1·005. Such a conclusion arrived at by direct experiment cannot
but be regarded as having greater weight than Prout's supposition
(hypothesis) that the atomic weights of the elements are in multiple
proportion to each other, which would give reason for surmising (but not
asserting) a complexity of nature in the elements, and their common
origin from a single primary material, and for expecting their mutual
conversion into each other. All such ideas and hopes must now, thanks
more especially to Stas, be placed in a region void of any experimental
support whatever, and therefore not subject to the discipline of the
positive data of science.
[27] This hypothesis, for the establishment or refutation of which so
many researches have been made, is exceedingly important, and
fully deserves the attention which has been given to it. Indeed,
if it appeared that the atomic weights of all the elements could
be expressed in whole numbers with reference to hydrogen, or if
they at least proved to be commensurable with one another, then it
could be affirmed with confidence that the elements, with all
their diversity, were formed of one material condensed or grouped
in various manners into the stable, and, under known conditions,
undecomposable groups which we call the atoms of the elements. At
first it was supposed that all the elements were nothing else but
condensed hydrogen, but when it appeared that the atomic weights
of the elements could not be expressed in whole numbers in
relation to hydrogen, it was still possible to imagine the
existence of a certain material from which both hydrogen and all
the other elements were formed. If it should transpire that four
atoms of this material form an atom of hydrogen, then the atom of
chlorine would present itself as consisting of 142 atoms of this
substance, the weight of whose atom would be equal to 0·25. But in
this case the atoms of all the elements should be expressed in
whole numbers with respect to the weight of the atom of this
original material. Let us suppose that the atomic weight of this
material is equal to unity, then all the atomic weights should be
expressible in whole numbers relatively to this unit. Thus the
atom of one element, let us suppose, would weigh _m_, and of
another _n_, but, as both _m_ and _n_ must be whole numbers, it
follows that the atomic weights of all the elements would be
commensurable. But it is sufficient to glance over the results
obtained by Stas, and to be assured of their accuracy, especially
for silver, in order to entirely destroy, or at least strongly
undermine, this attractive hypothesis. We must therefore refuse
our assent to the doctrine of the building up from a single
substance of the elements known to us. This hypothesis is not
supported either by any known transformation (for one element has
never been converted into another element), or by the
commensurability of the atomic weights of the elements. Although
the hypothesis of the formation of all the elements from a single
substance (for which Crookes has suggested the name protyle) is
most attractive in its comprehensiveness, it can neither be denied
nor accepted for want of sufficient data. Marignac endeavoured,
however, to overcome Stas's conclusions as to the
incommensurability of the atomic weights by supposing that in his,
as in the determinations of all other observers, there were
unperceived errors which were quite independent of the mode of
observation--for example, silver nitrate might be supposed to be
an unstable substance which changes, under the heatings,
evaporations, and other processes to which it is subjected in the
reactions for the determination of the combining weight of silver.
It might be supposed, for instance, that silver nitrate contains
some impurity which cannot be removed by any means; it might also
be supposed that a portion of the elements of the nitric acid are
disengaged in the evaporation of the solution of silver nitrate
(owing to the decomposing action of water), and in its fusion, and
that we have not to deal with normal silver nitrate, but with a
slightly basic salt, or perhaps an excess of nitric acid which
cannot be removed from the salt. In this case the observed
combining weight will not refer to an actually definite chemical
compound, but to some mixture for which there does not exist any
perfectly exact combining relations. Marignac upholds this
proposition by the fact that the conclusions of Stas and other
observers respecting the combining weights determined with the
greatest exactitude very nearly agree with the proposition of the
commensurability of the atomic weights--for example, the combining
weight of silver was shown to be equal to 107·93, so that it only
differs by 0·08 from the whole number 108, which is generally
accepted for silver. The combining weight of iodine proved to be
equal to 126·85--that is, it differs from 127 by 0·15. The
combining weights of sodium, nitrogen, bromine, chlorine, and
lithium are still nearer to the whole or round numbers which are
generally accepted. But Marignac's proposition will hardly bear
criticism. Indeed if we express the combining weights of the
elements determined by Stas in relation to hydrogen, the
approximation of these weights to whole numbers disappears,
because one part of hydrogen in reality does not combine with 16
parts of oxygen, but with 15·92 parts, and therefore we shall
obtain, taking H = 1, not the above-cited figures, but for silver
107·38, for bromine 79·55, magnitudes which are still further
removed from whole numbers. Besides which, if Marignac's
proposition were true the combining weight of silver determined by
one method--_e.g._ by the analysis of silver chlorate combined
with the synthesis of silver chloride--would not agree well with
the combining weight determined by another method--_e.g._ by means
of the analysis of silver iodate and the synthesis of silver
iodide. If in one case a basic salt could be obtained, in the
other case an acid salt might be obtained. Then the analysis of
the acid salt would give different results from that of the basic
salt. Thus Marignac's arguments cannot serve as a support for the
vindication of Prout's hypothesis.
In conclusion, I think it will not be out of place to cite the
following passage from a paper I read before the Chemical Society
of London in 1889 (Appendix II.), referring to the hypothesis of
the complexity of the elements recognised in chemistry, owing to
the fact that many have endeavoured to apply the periodic law to
the justification of this idea 'dating from a remote antiquity,
when it was found convenient to admit the existence of many gods
but only one matter.'
'When we try to explain the origin of the idea of a unique primary
matter, we easily trace that, in the absence of deductions from
experiment, it derives its origin from the scientifically
philosophical attempt at discovering some kind of unity in the
immense diversity of individualities which we see around. In
classical times such a tendency could only be satisfied by
conceptions about the immaterial world. As to the material world,
our ancestors were compelled to resort to some hypothesis, and
they adopted the idea of unity in the formative material, because
they were not able to evolve the conception of any other possible
unity in order to connect the multifarious relations of matter.
Responding to the same legitimate scientific tendency, natural
science has discovered throughout the universe a unity of plan, a
unity of forces, and a unity of matter; and the convincing
conclusions of modern science compel every one to admit these
kinds of unity. But while we admit unity in many things, we none
the less must also explain the individuality and the apparent
diversity which we cannot fail to trace everywhere. It was said of
old "Give us a fulcrum and it will become easy to displace the
earth." So also we must say, "Give us something that is
individualised, and the apparent diversity will be easily
understood." Otherwise, how could unity result in a multitude.
'After a long and painstaking research, natural science has
discovered the individualities of the chemical elements, and
therefore it is now capable, not only of analysing, but also of
synthesising; it can understand and grasp generality and unity, as
well as the individualised and multifarious. The general and
universal, like time and space, like force and motion, vary
uniformly. The uniform admit of interpolations, revealing every
intermediate phase; but the multitudinous, the
individualised--such as ourselves, or the chemical elements, or
the members of a peculiar periodic function of the elements, or
Dalton's multiple proportions--is characterised in another way. We
see in it--side by side with a general connecting
principle--leaps, breaks of continuity, points which escape from
the analysis of the infinitely small--an absence of complete
intermediate links. Chemistry has found an answer to the question
as to the causes of multitudes, and while retaining the conception
of many elements, all submitted to the discipline of a general
law, it offers an escape from the Indian Nirvana--the absorption
in the universal--replacing it by the individualised. However, the
place for individuality is so limited by the all-grasping,
all-powerful universal, that it is merely a point of support for
the understanding of multitude in unity.'
Among the platinum metals ruthenium, rhodium, and palladium, by their
atomic weights and properties, approach silver, just as iron and its
analogues (cobalt and nickel) approach copper in all respects. _Gold_
stands in exactly the same position in relation to the heavy platinum
metals, osmium, iridium, and platinum, as copper and silver do to the two
preceding series. The atomic weight of gold is nearly equal to their
atomic weights;[28] it is dense like these metals. It also gives various
grades of oxidation, which are feeble, both in a basic and an acid sense.
Whilst near to osmium, iridium, and platinum, gold at the same time is
able, like copper and silver, to form compounds which answer to the type
RX--that is, oxides of the composition R_{2}O. Cuprous chloride, CuCl,
silver chloride, AgCl, and aurous chloride, AuCl, are substances which
are very much alike in their physical and chemical properties.[28 bis]
They are insoluble in water, but dissolve in hydrochloric acid and
ammonia, in potassium cyanide, sodium thiosulphate, &c. Just as copper
forms a link between the iron metals and zinc, and as silver unites the
light platinum metals with cadmium, so also gold presents a transition
from the heavy platinum metals to mercury. Copper gives saline compounds
of the types CuX and CuX_{2}, silver of the type AgX, whilst gold,
besides compounds of the type AuX, very easily and most frequently forms
those of the type AuCl_{3}. The compounds of this type frequently pass
into those of the lower type, just as PtX_{4} passes into PtX_{2}, and
the same is observable in the elements which, in their atomic weights,
follow gold. Mercury gives HgX_{2} and HgX, thallium gives TlX_{3} and
TlX, lead gives PbX_{4} and PbX_{2}. On the other hand, gold in a
qualitative respect differs from silver and copper in the _extreme ease_
with which all its compounds are _reduced to metal_ by many means. This
is not only accomplished by many reducing agents, but also by the action
of heat. Thus its chlorides and oxides lose their chlorine and oxygen
when heated, and, if the temperature be sufficiently high, these elements
are entirely expelled and metallic gold alone remains. Its compounds,
therefore, act as oxidising agents.[29]
[28] It might be expected from the periodic law and analogies with the
series iron, cobalt, nickel, copper, zinc, that the atomic weights
of the elements of the series osmium, iridium, platinum, gold,
mercury, would rise in this order, and at the time of the
establishment of the periodic law (1869), the determinations of
Berzelius, Rose, and others gave the following values for the
atomic weights: Os = 200, Ir = 197, Pt = 198, Au = 196, Hg = 200.
The fulfilment of the expectations of the periodic law was given
in the first place by the fresh determinations (Seubert, Dittmar,
and Arthur) of the atomic weight of platinum, which proved to be
nearly 196, if O = 16 (as Marignac, Brauner, and others propose);
in the second place, by the fact that Seubert proved that the
atomic weight of osmium is really less than that of platinum, and
approximately Os = 191; and, in the third place, by the fact that
after the researches of Krüss, Thorpe, and Laurie there was no
doubt that the atomic weight of gold is greater than that of
platinum--namely, nearly 197.
[28 bis] In Chapter XXII., Note 40, we gave the thermal data for
certain of the compounds of copper of the type CuX_{2}; we will
now cite certain data for the cuprous compounds of the type CuX,
which present an analogy to the corresponding compounds AgX and
AuX, some of which were investigated by Thomsen in his classical
work, 'Thermochemische Untersuchungen' (Vol. iii., 1883). The data
are given in the same manner as in the above-mentioned note:
R = Cu Ag Au
R + Cl +33 +29 +6
R + Br +25 +23 0
R + I +16 +14 -6
R + O +41 + 6 -?
Thus we see in the first place that gold, which possesses a much
smaller affinity than Ag, evolves far less heat than an equivalent
amount of copper, giving the same compound, and in the second
place that the combination of copper with one atom of oxygen
disengages more heat than its combination with one atom of a
halogen, whilst with silver the reverse is the case. This is
connected with the fact that Cu_{2}O is more stable under the
action of heat than Ag_{2}O.
[29] Heavy atoms and molecules, although they may present many points
of analogy, are more easily isolated; thus C_{16}H_{32}, although,
like C_{2}H_{4}, it combines with Br_{2}, and has a similar
composition, yet reacts with much greater difficulty than
C_{2}H_{4}, and in this it resembles gold; the heavy atoms and
molecules are, so to say, inert, and already saturated by
themselves. Gold in its higher grade of oxidation, Au_{2}O_{3},
presents feeble basic properties and weakly-developed acid
properties, so that this oxide of gold, Au_{2}O_{3}, may be
referred to the class of feeble acid oxides, like platinic oxide.
This is not the case in the highest known oxides of copper and
silver. But in the lower grade of oxidation, aurous oxide,
Au_{2}O, gold, like silver and copper, presents basic properties,
although they are not very pronounced. In this respect it stands
very close in its properties, although not in its types of
combination (AuX and AuX_{3}), to platinum (PtX_{2} and PtX_{4})
and its analogues.
As yet the general chemical characteristics of gold and its
compounds have not been fully investigated. This is partly due to
the fact that very few researches have been undertaken on the
compounds of this metal, owing to its inaccessibility for working
in large quantities. As the atomic weight of gold is high (Au =
197), the preparation of its compounds requires that it should be
taken in large quantities, which forms an obstacle to its being
fully studied. Hence the facts concerning the history of this
metal are rarely distinguished by that exactitude with which many
facts have been established concerning other elements more
accessible, and long known in use.
_In nature_ gold occurs in the primary and chiefly in quartzose rocks,
and especially in quartz veins, as in the Urals (at Berezoffsk), in
Australia, and in California. The native gold is extracted from these
rocks by subjecting them to a mechanical treatment consisting of crushing
and washing.[29 bis] Nature has already accomplished a similar
disintegration of the hard rocky matter containing gold.[30] These
disintegrated rocks, washed by rain and other water, have formed
gold-bearing deposits, which are known as _alluvial gold deposits_.
Gold-bearing soil is sometimes met with on the surface and sometimes
under the upper soil, but more frequently along the banks of dried-up
water-courses and running streams. The sand of many rivers contains,
however, a very small amount of gold, which it is not profitable to work;
for example, that of the Alpine rivers contains 5 parts of gold in
10,000,000 parts of sand. The richest gold deposits are those of Siberia,
especially in the southern parts of the Government of Yeniseisk, the
South Urals, Mexico, California, South Africa, and Australia, and then
the comparatively poorer alluvial deposits of many countries (Hungary,
the Alps, and Spain in Europe). The extraction of the gold from alluvial
deposits is based on the principle of levigation; the earth is washed,
while constantly agitated, by a stream of water, which carries away the
lighter portion of the earth, and leaves the coarser particles of the
rock and heavier particles of the gold, together with certain substances
which accompany it, in the washing apparatus. The extraction of this
_washed_ gold only necessitates mechanical appliances,[31] and it is not
therefore surprising that gold was known to savages and in the most
remote period of history. It sometimes occurs in crystals belonging to
the regular system, but in the majority of cases in nuggets or grains of
greater or less magnitude. It always contains silver (from very small
quantities up to 30 p.c., when it is called 'electrum') and certain other
metals, among which lead and rhodium are sometimes found.
[29 bis] Sonstadt (1872) showed that sea water, besides silver, always
contains gold. Munster (1892) showed that the water of the
Norwegian fiords contains about 5 milligrams of gold per ton (or 5
milliardths)--_i.e._ a quantity deserving practical attention, and
I think it may be already said that, considering the immeasurable
amount of sea water, in time means will be discovered for
profitably extracting gold from sea water by bringing it into
contact with substances capable of depositing gold upon their
surface. The first efforts might be made upon the extraction of
salt from sea water, and as the total amount of sea water may be
taken as about 2,000,000,000,000,000,000 tons, it follows that it
contains about 10,000 million tons of gold. The yearly production
of gold is about 200 tons for the whole world, of which about one
quarter is extracted in Russia. It is supposed that gold is
dissolved in sea water owing to the presence of iodides, which,
under the action of animal organisms, yield free iodine. It is
thought (as Professor Konovaloff mentions in his work upon 'The
Industries of the United States,' 1894) that iodine facilitates
the solution of the gold, and the organic matter its
precipitation. These facts and considerations to a certain extent
explain the distribution of gold in veins or rock fissures,
chiefly filled with quartz, because there is sufficient reason for
supposing that these rocks once formed the ocean bottom. R.
Dentrie, and subsequently Wilkinson, showed that organic
matter--for instance, cork--and pyrites are able to precipitate
gold from its solutions in that metallic form and state in which
it occurs in quartz veins, where (especially in the deeper parts
of vein deposits) gold is frequently found on the surface of
pyrites, chiefly arsenical pyrites. Kazantseff (in Ekaterinburg,
1891) even supposes, from the distribution of the gold in these
pyrites, that it occurred in solution as a compound of sulphide of
gold and sulphide of arsenic when it penetrated into the veins. It
is from such considerations that the origin of vein and pyritic
gold is, at the present time, attributed to the reaction of
solutions of this metal, the remains of which are seen in the gold
still present in sea water.
[30] However, in recent times, especially since about 1870, when
chlorine (either as a solution of the gas or as bleaching powder)
and bromine began to be applied to the extraction of
finely-divided gold from poor ores (previously roasted in order to
drive off arsenic and sulphur, and oxidise the iron), the
extraction of gold from quartz and pyrites, by the wet method,
increases from year to year, and begins to equal the amount
extracted from alluvial deposits. Since the nineties the _cyanide
process_ (Chapter XIII., Note 13 bis) has taken an important place
among the wet methods for extracting gold from its ores. It
consists in pouring a dilute solution of cyanide of potassium
(about 500 parts of water and 1 to 4 parts of cyanide of potassium
per 1,000 parts of ore, the amount of cyanide depending upon the
richness of the ore) and a mixture of it with NaCN, (_see_ Chapter
XIII., Note 12) over the crushed ore (which need not be roasted,
whilst roasting is indispensable in the chlorination process, as
otherwise the chlorine is used up in oxidising the sulphur,
arsenic, &c.) The gold is dissolved very rapidly even from
pyrites, where it generally occurs on the surface in such fine and
adherent particles that it either cannot be mechanically washed
away, or, more frequently is carried away by the stream of water,
and cannot be caught by mechanical means or by the mercury used
for catching the gold in the sluices. Chlorination had already
given the possibility of extracting the finest particles of gold;
but the cyanide process enables such pyrites to be treated as
could be scarcely worked by other means. The treatment of the
crushed ore by the KCN is carried on in simple wooden vats (coated
with paraffin or tar) with the greatest possible rapidity (in
order that the KCN solution should not have time to change) by a
method of systematic lixiviation, and is completed in 10 to 12
hours. The resultant solution of gold, containing AuK(CN)_{2}, is
decomposed either with freshly-made zinc filings (but when the
gold settles on the Zn, the cyanide solution reacts upon the Zn
with the evolution of H_{2} and formation of ZnH_{2}O_{2}) or by
sodium amalgam prepared at the moment of reaction by the action of
an electric current upon a solution of NaHO poured into a vessel
partially immersed in mercury (the NaCN is renewed continually by
this means). The silver in the ore passes into solution, together
with the gold, as in amalgamation.
[31] But the particles of gold are sometimes so small that a large
amount is lost during the washing. It is then profitable to have
recourse to the extraction by chlorine and KCN (Note 30).
In speaking of the extraction of gold the following remarks may
not be out of place:
In California advantage is taken of water supplied from high
altitudes in order to have a powerful head of water, with which
the rocks are directly washed away, thus avoiding the greater
portion of the mechanical labour required for the exploitation of
these deposits.
The last residues of gold are sometimes extracted from sand by
washing them with mercury, which dissolves the gold. The sand
mixed with water is caused to come into contact with mercury
during the washing. The mercury is then distilled.
Many sulphurous ores, even pyrites, contain a small amount of
gold. Compounds of gold with bismuth, BiAu_{2}, tellurium,
AuTe_{2} (calverite), &c., have been found, although rarely.
Among the minerals which accompany gold, and from which the
presence of gold may be expected, we may mention white quartz,
titanic and magnetic iron ores, and also the following, which are
of rarer occurrence: zircon, topaz, garnet, and such like. The
concentrated gold washings first undergo a mechanical treatment,
and the impure gold obtained is treated for pure gold by various
methods. If the gold contain a considerable amount of foreign
metals, especially lead and copper, it is sometimes cupelled, like
silver, so that the oxidisable metals may be absorbed by the cupel
in the form of oxides, but in every case the gold is obtained
together with silver, because the latter metal also is not
oxidised. Sometimes the gold is extracted by means of mercury,
that is, by amalgamation (and the mercury subsequently driven off
by distillation), or by smelting it with lead (which is afterwards
removed by oxidation) and processes like those employed for the
extraction of silver, because gold, like silver, does not oxidise,
is dissolved by lead and mercury, and is non-volatile. If copper
or any other metal contain gold and it be employed as an anode,
pure copper will be deposited upon the cathode, while all the gold
will remain at the anode as a slime. This method often amply
repays the whole cost of the process, since it gives, besides the
gold, a pure electrolytic copper.
_The separation of the silver_ from gold is generally carried on with
great precision, as the presence of the silver in the gold does not
increase its value for exchange, and it can be substituted by other less
valuable metals, so that the extraction of the silver, as a precious
metal, from its alloy with gold, is a profitable operation. This
separation is conducted by different methods. Sometimes the argentiferous
gold is melted in crucibles, together with a mixture of common salt and
powdered bricks. The greater portion of the silver is thus converted into
the chloride, which fuses and is absorbed by the slags, from which it may
be extracted by the usual methods. The silver is also extracted from gold
by treating it with boiling sulphuric acid, which does not act on the
gold but dissolves the silver. But if the alloy does not contain a large
proportion of silver it cannot be extracted by this method or at all
events the separation will be imperfect, and therefore a fresh amount of
silver is added (by fusion) to the gold, in such quantity that the alloy
contains twice as much silver as gold. The silver which is added is
preferably such as contains gold, which is very frequently the case. The
alloy thus formed is poured in a thin stream into water, by which means
it is obtained in a granulated form; it is then boiled with strong
sulphuric acid, three parts of acid being used to one part of alloy. The
sulphuric acid extracts all the silver without acting on the gold. It is
best, however, to pour off the first portion of the acid, which has
dissolved the silver, and then treat the residue of still imperfectly
pure gold with a fresh quantity of sulphuric acid. The gold is thus
obtained in the form of powder, which is washed with water until it is
quite free from silver. The silver is precipitated from the solution by
means of copper, so that cupric sulphate and metallic silver are
obtained. This process is carried out in many countries, as in Russia, at
the Government mints.
Gold is generally used alloyed with copper; since pure gold, like pure
silver, is very soft, and therefore soon worn away. In assaying or
determining the amount of pure gold in such an alloy it is usual to add
silver to the gold in order to make up an alloy containing three parts of
silver to one of gold (this is known as quartation because the alloy
contains 1/4 of gold), and the resultant alloy is treated with nitric
acid. If the silver be not in excess over the gold, it is not all
dissolved by the nitric acid, and this is the reason for the quartation.
The amount of pure gold (assay) is determined by weighing the gold which
remains after this treatment. English gold (= 22 carats) coinage is
composed of an alloy containing 91·66 p.c. of gold, but for many articles
gold is frequently used containing a larger amount of foreign metals.
_Pure gold_ may be obtained from gold alloys by dissolving in aqua
regia, and then adding ferrous sulphate to the solution or heating it
with a solution of oxalic acid. These deoxidising agents reduce the gold,
but not the other metals. The chlorine combined with the gold then acts
like free chlorine. The gold, thus reduced, is precipitated as an
exceedingly fine brown powder.[31 bis] It is then washed with water, and
fused with nitre or borax. Pure gold reflects a yellow light, and in the
form of very thin sheets (gold leaf), into which it can be hammered and
rolled,[31 tri] it transmits a bluish-green light. The specific gravity
of gold is about 19·5, the sp. gr. of gold coin is about 17·1. It fuses
at 1090°--at a higher temperature than silver--and can be drawn into
exceedingly fine wires or hammered into thin sheets. With its softness
and ductility, gold is distinguished for its tenacity, and a gold wire
two millimetres thick breaks only under a load of 68 kilograms. Gold
vaporises even at a furnace heat, and imparts a greenish colour to a
flame passing over it in a furnace. Gold alloys with copper almost
without changing its volume.[32] In its chemical aspect, gold presents,
as is already seen from its general characteristics given above, an
example of the so-called noble metals--_i.e._ it is incapable of being
oxidised at any temperature, and its oxide is decomposed when calcined.
Only chlorine and bromine combine directly with it at the ordinary
temperature, but many other metals and non-metals combine with it at a
red heat--for example, sulphur, phosphorus, and arsenic. Mercury
dissolves it with great ease. It dissolves in potassium cyanide in the
presence of air; a mixture of sulphuric acid with nitric acid dissolves
it with the aid of heat, although in small quantity. It is also soluble
in aqua regia and in selenic acid. Sulphuric, hydrochloric, nitric, and
hydrofluoric acids and the caustic alkalis do not act on gold, but a
mixture of hydrochloric acid with such oxidising agents as evolve
chlorine naturally dissolves it like aqua regia.[32 bis]
[31 bis] Schottländer (1893) obtained gold in a soluble colloid form
(the solution is violet) by the action of a mixture of solutions
of cerium acetate and NaHO upon a solution of AuCl_{3}. The gold
separates out from such a solution in exactly the same manner as
Ag does from the solution of colloid silver mentioned above. There
always remains a certain amount of a higher oxide of cerium,
CeO_{2}, in the solution--_i.e._ the gold is reduced by converting
the cerium into a higher grade of oxidation. Besides which Krüss
and Hofmann showed that sulphide of gold precipitated by the
action of H_{2}S upon a solution of AuKCy_{2} mixed with HCl
easily passes into a colloid solution after being properly washed
(like As_{2}S_{3}, CuS, &c., Chapter I., Note 57).
[31 tri] Gold-leaf is used for gilding wood (leather, cardboard, and
suchlike, upon which it is glued by means of varnish, &c.), and is
about 0·003 millimetre thick. It is obtained from thin sheets
(weighing at first about 1/4 grm. to a square inch), rolled
between gold rollers, by gradually hammering them (in packets of a
number at once) between sheets of moist (but not wet) parchment,
and then, after cutting them into four pieces, between a specially
prepared membrane, which, when at the right degree of moisture,
does not tear or stick together under the blows of the hammer.
[32] The formation of the alloys Cu + Zn, Cu + Sn, Cu + Bi, Cu + Sb,
Pb + Sb, Ag + Pb, Ag + Sn, Au + Zn, Au + Sn, &c., is accompanied
by a contraction (and evolution of heat). The formation of the
alloys Fe + Sb, Fe + Pb, Cu + Pb, Pb + Sn, Pb + Sb, Zn + Sb, Ag +
Cu, Au + Cu, Au + Pb, takes place with a certain increase in
volume. With regard to the alloys of gold, it may be mentioned
that gold is only slightly dissolved by mercury (about 0·06 p.c.,
Dudley, 1890); the remaining portion forms a granular alloy, whose
composition has not been definitely determined. Aluminium (and
silicon) also have the capacity of forming alloys with gold. The
presence of a small amount of aluminium lowers the melting point
of gold considerably (Roberts-Austen, 1892); thus the addition of
4 p.c. of aluminium lowers it by 14°·28, the addition of 10 p.c.
Al by 41°·7. The latter alloy is white. The alloy AuAl_{2} has a
characteristic purple colour, and its melting point is 32°·5 above
that of gold, which shows it to be a definite compound of the two
metals. The melting points of alloys richer in Al gradually fall
to 660°--that is, below that of aluminium (665°).
Heycock and Neville (1892), in studying the triple alloys of Au,
Cd, and Sn, observed a tendency in the gold to give compounds with
Cd, and by sealing a mixture of Au and Cd in a tube, from which
the air had been exhausted, and heating it, they obtained a grey
crystalline brittle definite alloy AuCd.
[32 bis] Calderon (1892), at the request of some jewellers,
investigated the cause of a peculiar alteration sometimes found on
the surface of dead-gold articles, there appearing brownish and
blackish spots, which widen and alter their form in course of
time. He came to the conclusion that these spots are due to the
appearance and development of peculiar micro-organisms
(Aspergillus niger and Micrococcus cimbareus) on the gold, spores
of which were found in abundance on the cotton-wool in which the
gold articles had been kept.
As regards the compounds of gold, they belong, as was said above, to the
types AuX_{3} and AuX. _Auric chloride_ or _gold trichloride_, AuCl_{3},
which is formed when gold is dissolved in aqua regia, belongs to the
former and higher of these types. The solution of this substance in water
has a yellow colour, and it may be obtained pure by evaporating the
solution in aqua regia to dryness, but not to the point of decomposition.
If the evaporation proceed to the point of crystallisation, a compound of
gold chloride and hydrochloric acid, AuHCl_{4}, is obtained, like the
allied compounds of platinum; but it easily parts with the acid and
leaves auric chloride, which fuses into a red-brown liquid, and then
solidifies to a crystalline mass. If dry chlorine be passed over gold in
powder it forms a mixture of aurous and auric chlorides, but the aurous
chloride is also decomposed by water into gold and auric chloride. Auric
chloride crystallises from its solutions as AuCl_{3},2H_{2}O, which
easily loses water, and the dry chloride loses two-thirds of its chlorine
at 185°, forming aurous chloride, whilst above 300° the latter chloride
also loses its chlorine and leaves metallic gold. Auric chloride is the
usual form in which gold occurs in solutions, and in which its salts are
used in the arts and for chemical purposes. It is soluble in water,
alcohol, and ether. Light has a reducing action on these solutions, and
after a time metallic gold is deposited upon the sides of vessels
containing the solution. Hydrogen when nascent, and even in a gaseous
form, reduces gold from this solution to a metallic state. The reduction
is more conveniently and usually effected by ferrous sulphate, and in
general by the action of ferrous salts.[33]
[33] Stannous chloride as a reducing agent also acts on auric chloride,
and gives a red precipitate known as _purple of Cassius_. This
substance, which probably contains a mixture or compound of aurous
oxide and tin oxide, is used as a red pigment for china and glass.
Oxalic acid, on heating, reduces metallic gold from its salts, and
this property may be taken advantage of for separating it from its
solutions. The oxidation which then takes place in the presence of
water may be expressed by the following equation: 2AuCl_{3} +
3C_{2}H_{2}O_{4} = 2Au + 6HCl + 6CO_{2}. Nearly all organic
substances have a reducing action on gold, and solutions of gold
leave a violet stain on the skin.
Auric chloride, like platinic chloride, is distinguished for its
clearly-developed property of forming double salts. These double
salts, as a rule, belong to the type AuMCl_{4}. The compound of
auric chloride with hydrochloric acid mentioned above evidently
belongs to the same type. The compounds 2KAuCl_{4},5H_{2}O,
NaAuCl_{4},2H_{2}O, AuNH_{4}Cl_{4},H_{2}O,
Mg(AuCl_{4})_{2},2H_{2}O, and the like are easily crystallised in
well-formed crystals. Wells, Wheeler, and Penfield (1892) obtained
RbAuCl_{4} (reddish yellow) and CsAuCl_{4} (golden yellow), and
corresponding bromides (dark coloured). AuBr_{3} is extremely like
the chloride. Auric cyanide is obtained easily in the form of a
double salt of potassium, KAu(CN)_{4} by mixing saturated and hot
solutions of potassium cyanide with auric chloride and then
cooling.
If a solution of potassium hydroxide be added to a solution of auric
chloride, a precipitate is first formed, which re-dissolves in an excess
of the alkali. On being evaporated under the receiver of an air-pump,
this solution yields yellow crystals, which present the same composition
as the double salts AuMCl_{4}, with the substitution of the chlorine by
oxygen--that is to say, _potassium aurate_, AuKO_{2}, is formed in
crystals containing 3H_{2}O. The solution has a distinctly alkaline
reaction. _Auric oxide_, Au_{2}O_{3}, separates when this alkaline
solution is boiled with an excess of sulphuric acid. But it then still
retains some alkali; however, it may be obtained in a pure state as a
brown powder by dissolving in nitric acid and diluting with water. The
brown powder decomposes below 250° into gold and oxygen. It is insoluble
in water and in many acids, but it dissolves in alkalis, which shows the
acid character of this oxide. An hydroxide, Au(OH)_{3} may be obtained as
a brown powder by adding magnesium oxide to a solution of auric chloride
and treating the resultant precipitate of magnesium aurate with nitric
acid. This hydroxide loses water at 100°, and gives auric oxide.[34]
[34] If ammonia be added to a solution of auric chloride, it forms a
yellow precipitate of the so-called fulminating gold, which
contains gold, chlorine, hydrogen, nitrogen, and oxygen, but its
formula is not known with certainty. It is probably a sort of
ammonio-metallic compound, Au_{2}O_{3},4NH_{3}, or amide (like the
mercury compound). This precipitate explodes at 140°, but when
left in the presence of solutions containing ammonia it loses all
its chlorine and becomes non-explosive. In this form the
composition Au_{2}O_{3},2NH_{3},H_{2}O is ascribed to it, but this
is uncertain. Auric sulphide, Au_{2}S_{3}, is obtained by the
action of hydrogen sulphide on a solution of auric chloride, and
also directly by fusing sulphur with gold. It has an acid
character, and therefore dissolves in sodium and ammonium
sulphides.
The starting-point of the compounds of the type AuX[35] is _gold
monochloride_ or _aurous chloride_, AuCl, which is formed, as mentioned
above, by heating auric chloride at 185°. Aurous chloride forms a
yellowish-white powder; this, when heated with water, is decomposed into
metallic gold and auric chloride, which passes into solution: 3AuCl =
AuCl_{3} + 2Au. This decomposition is accelerated by the action of light.
Hence it is obvious that the compounds corresponding with aurous oxide
are comparatively unstable. But this only refers to the simple compounds
AuX; some of the complex compounds, on the contrary, form the most stable
compounds of gold. Such, for example, is the cyanide of gold and
potassium, AuK(CN)_{2}. It is formed, for instance, when finely-divided
gold dissolves in the presence of air in a solution of potassium cyanide:
4KCN + 2Au + H_{2}O + O = 2KAu(CN)_{2} + 2KHO (this reaction also
proceeds with solid pieces of gold, although very slowly). The same
compound is formed in solution when many compounds of gold are mixed with
potassium cyanide, because if a higher compound of gold be taken, it is
reduced by the potassium cyanide into aurous oxide, which dissolves in
potassium cyanide and forms KAu(CN)_{2}. This substance is soluble in
water, and gives a colourless solution, which can be kept for a long
time, and is employed in electro-gilding--that is, for coating other
metallic objects with a layer of gold, which is deposited if the object
be connected with the negative pole of a battery and the positive pole
consist of a gold plate. When an electric current is passed between them,
the gold from the latter will dissolve, whilst a coating of gold from the
solution will be deposited on the object.
[35] Many double salts of suboxide of gold belong to the type AuX--for
instance, the cyanide corresponding to the type AuKX_{2}, like
PtK_{2}X_{4}, with which we became acquainted in the last chapter.
We will enumerate several of the representatives of this class of
compounds. If auric chloride, AuCl_{3}, be mixed with a solution
of sodium thiosulphate, the gold passes into a colourless
solution, which deposits colourless crystals, containing a double
thiosulphate of gold and sodium, which are easily soluble in water
but are precipitated by alcohol. The composition of this salt is
Na_{3}Au(S_{2}O_{3})_{2},2H_{2}O. If the sodium thiosulphate be
represented as NaS_{2}O_{3}Na, the double salt in question will be
AuNa(S_{2}O_{3}Na)_{2},2H_{2}O, according to the type AuNaX_{2}.
The solution of this colourless and easily crystallisable salt has
a sweet taste, and the gold is not separated from it either by
ferrous sulphate or oxalic acid. This salt, which is known as
_Fordos and Gelis's salt_, is used in medicine and photography. In
general, aurous oxide exhibits a distinct inclination to the
formation of similar double salts, as we saw also with
PtX_{2}--for example, it forms similar salts with sulphurous acid.
Thus if a solution of sodium sulphite be gradually added to a
solution of oxide of gold in sodium hydroxide, the precipitate at
first formed re-dissolves to a colourless solution, which contains
the double salt Na_{3}Au(SO_{3})_{2} = AuNa(SO_{3}Na)_{2}. The
solution of this salt, when mixed with barium chloride, first
forms a precipitate of barium sulphite, and then a red barium
double salt which corresponds with the above sodium salt.
The oxygen compound of the type AuX, _aurous oxide_, Au_{2}O, is
obtained as a greenish violet powder on mixing aurous chloride
with potassium chloride in the cold. With hydrochloric acid this
oxide gives gold and auric chloride, and when heated it easily
splits up into oxygen and metallic gold.
APPENDIX I