Chapter X: , Note 2 (1)

[2] Carnallite belongs to the number of double salts which are directly
decomposed by water, and it only crystallises from solutions which
contain an excess of magnesium chloride. It may be prepared
artificially by mixing strong solutions of potassium and magnesium
chlorides, when colourless crystals of sp. gr. 1·60 separate,
whilst the Stassfurt salt is usually of a reddish tint, owing to
traces of iron. At the ordinary temperature sixty-five parts of
carnallite are soluble in one hundred parts of water in the
presence of an excess of the salt. It deliquesces in the air,
forming a solution of magnesium chloride and leaving potassium
chloride. The quantity of carnallite produced at Stassfurt is now
as much as 100,000 tons a year.

[3] The method of separating sodium chloride from potassium chloride
has been described in Chapter I. On evaporation of a mixture of the
saturated solutions, sodium chloride separates; and then, on
cooling, potassium chloride separates, owing to the difference of
rate of variation of their solubilities with the temperature. The
following are the most trustworthy figures for the solubility of
_potassium chloride_ in one hundred parts of water (for sodium
chloride, _see_ Chapter X., Note 13):--

10° 20° 40° 60° 100°
32 35 40 46 57

When mixed with solutions of other salts the solubility of
potassium chloride naturally varies, but not to any great extent.

[3 bis] The specific gravity of the solid salt is 1·99--that is, less
than that of sodium chloride. All the salts of sodium are
specifically heavier than the corresponding salts of potassium, as
are also their solutions for equal percentage compositions. If the
specific gravity of water at 4° = 10,000, then at 15° the specific
gravity of a solution of _p_ p.c. potassium chloride = 9,992 +
63·29_p_ + 0·226_p_^2, and therefore for 10 p.c. = 1·0647, 20 p.c.
= 1·1348, &c.

Potassium chloride combines with iodine trichloride to form a
compound KCl + ICl_{3} = KICl_{4}, which has a yellow colour, is
fusible, loses iodine trichloride at a red heat, and gives
potassium iodate and hydrochloric acid with water. It is not only
formed by direct combination, but also by many other methods; for
instance, by passing chlorine into a solution of potassium iodide
so long as the gas is absorbed, KI + 2Cl_{2} = KCl,ICl_{3}.
Potassium iodide, when treated with potassium chlorate and strong
hydrochloric acid, also gives this compound; another method for its
formation is given by the equation KClO_{3} + I + 6HCl =
KCl,ICl_{3} + 3Cl + 3H_{2}O. This is a kind of salt corresponding
with KIO_{2} (unknown) in which the oxygen is replaced by chlorine.
If valency be taken as the starting-point in the study of chemical
compounds, and the elements considered as having a constant
atomicity (number of bonds)--that is, if K, Cl, and I be taken as
univalent elements--then it is impossible to explain the formation
of such a compound because, according to this view, univalent
elements are only able to form dual compounds with each other; such
as, KCl, ClI, KI, &c., whilst here they are grouped together in the
molecule KICl_{4}. Wells, Wheeler, and Penfield (1892) obtained a
large number of such poly-haloid salts. They may all be divided
into two large classes: the tri-haloid and the penta-haloid salts.
They have been obtained not only for K but also for Rb and Cs, and
partially also for Na and Li. The general method of their formation
consists in dissolving the ordinary halogen salt of the metal in
water, and treating it with the requisite amount of free halogen.
The poly-haloid salt separates out after evaporating the solution
at a more or less low temperature. In this manner, among the
tri-haloid salts, may be obtained: KI_{3}, KBr_{2}I, KCl_{2}I, and
the corresponding salts of rubidium and cæsium, for instance,
CsI_{3}, CsBrI_{2}, CsBr_{2}I, CsClBrI,CsCl_{2}I, CsBr_{3},
CsClBr_{2}, CsCl_{2}Br, and in general MX_{3} where X is a halogen.
The colour of the crystals varies according to the halogen, thus
CsI_{3} is black, CrBr_{3} yellowish red, CrBrI_{2} reddish brown,
CsBr_{2}I red, CsCl_{2}Br yellow. The cæsium salts are the most
stable, and those of potassium least so, as also those which
contain Br and I separately or together; for cæsium no compounds
containing Cl and I were obtained. The penta-haloid salts form a
smaller class; among these salts potassium forms KCl_{4}I, rubidium
RbCl_{4}I, cæsium CsI_{5}, CsBr, CsCl_{4}I, lithium LiCl_{4}I (with
4H_{2}O) and sodium NaCl_{4}I (with 2H_{2}O). The most stable are
those salts containing the metal with the greatest atomic
weight--cæsium (_see_ Chapter XI., Note 63).

[4] It is possible to extract the compounds of potassium directly from
the primary rocks which are so widely distributed over the earth's
surface and so abundant in some localities. From a chemical point
of view this problem presents no difficulty; for instance, by
fusing powdered orthoclase with lime and fluor spar (Ward's method)
and then extracting the alkali with water (on fusion the silica
gives an insoluble compound with lime), or by treating the
orthoclase with hydrofluoric acid (in which case silicon fluoride
is evolved as a gas) it is possible to transfer the alkali of the
orthoclase to an aqueous solution, and to separate it in this
manner from the other insoluble oxides. However, as yet there is no
profit in, nor necessity for, recourse to this treatment, as
carnallite and potash form abundant materials for the extraction of
potassium compounds by cheaper methods. Furthermore, the salts of
potassium are now in the majority of chemical reactions replaced by
salts of sodium, especially since the preparation of sodium
carbonate has been facilitated by the Leblanc process. The
replacement of potassium compounds by sodium compounds not only has
the advantage that the salts of sodium are in general cheaper than
those of potassium, but also that a smaller quantity of a sodium
salt is needed for a given reaction than of a potassium salt,
because the combining weight of sodium (23) is less than that of
potassium (39).

[5] It has been shown by direct experiment on the cultivation of plants
in artificial soils and in solutions that under conditions
(physical, chemical, and physiological) otherwise identical plants
are able to thrive and become fully developed in the entire absence
of sodium salts, but that their development is impossible without
potassium salts.

The primary rocks contain an almost equal proportion of potassium and sodium. But in sea water the compounds of the latter metal predominate. It may be asked, what became of the compounds of potassium in the disintegration of the primary rocks, if so small a quantity went to the sea water? They remained with the other products of the decomposition of the primary rocks. When granite or any other similar rock formation is disintegrated, there are formed, besides the soluble substances, also insoluble substances--sand and finely-divided clay, containing water, alumina, and silica. This clay is carried away by the water, and is then deposited in strata. It, and especially its admixture with vegetable remains, retain compounds of potassium in a greater quantity than those of sodium. This has been proved with absolute certainty to be the case, and is due to the _absorptive power of the soil_. If a dilute solution of a potassium compound be filtered through common mould used for growing plants, containing clay and the remains of vegetable decomposition, this mould will be found to have retained a somewhat considerable percentage of the potassium compounds. If a salt of potassium be taken, then during the filtration an equivalent quantity of a salt of calcium--which is also found, as a rule, in soils--is set free. Such a process of filtration through finely divided earthy substances proceeds in nature, and the compounds of potassium are everywhere retained by the friable earth in considerable quantity. This explains the presence of so small an amount of potassium salts in the water of rivers, lakes, streams, and oceans, where the lime and soda have accumulated. The compounds of potassium retained by the friable mass of the earth are absorbed as an aqueous solution by the roots of _plants_. Plants, as everyone knows, when burnt leave an ash, and this ash, besides various other substances, without exception contains compounds of potassium. Many land plants contain a very small amount of sodium compounds,[6] whilst potassium and its compounds occur in all kinds of vegetable ash. Among the generally cultivated plants, grass, potatoes, the turnip, and buckwheat are particularly rich in potassium compounds. The ash of plants, and especially of herbaceous plants, buckwheat straw, sunflower and potato leaves are used in practice for the extraction of potassium compounds. There is no doubt that potassium occurs in the plants themselves in the form of complex compounds, and often as salts of organic acids. In certain cases such salts of potassium are even extracted from the juice of plants. Thus, sorrel and oxalis, for example, contain in their juices the acid oxalate of potassium, C_{2}HKO_{4}, which is employed for removing ink stains. Grape juice contains the so-called cream of tartar, which is the acid tartrate of potassium, C_{4}H_{5}KO_{6}.[7] This salt also separates as a sediment from wine. When the plants, containing one or more of the salts of potassium, are burnt, the carbonaceous matter is oxidised, and in consequence the potassium is obtained in the ash as carbonate, K_{2}CO_{3}, which is generally known as _potashes_. Hence potashes occur ready prepared in the ash of plants, and therefore the ash of land plants is employed as a source for the extraction of potassium compounds. Potassium carbonate is extracted by lixiviating the ash with water.[8] Potassium carbonate may also be obtained from the chloride by a method similar to that by which sodium carbonate is prepared from sodium chloride.[8 bis] There is no difficulty in obtaining any salt of potassium--for example, the sulphate,[9] bromide, and iodide[10]--by the action of the corresponding acid on KCl and especially on the carbonate, whilst the hydroxide, _caustic potash_, KHO, which is in many respects analogous to caustic soda, is easily obtained by means of lime in exactly the same manner in which sodium hydroxide is prepared from sodium carbonate.[11] Therefore, in order to complete our knowledge of the alkali metals, we will only describe two salts of potassium which are of practical importance, and whose analogues have not been described in the preceding chapter, potassium cyanide and potassium nitrate.

[6] If herbaceous plants contain much sodium salts, it is evident that
these salts mainly come from the sodium compounds in the water
absorbed by the plants.

[7] As plants always contain mineral substances and cannot thrive in
a medium which does not contain them, more especially in one which
is free from the salts of the four basic oxides, K_{2}O, CaO, MgO,
and Fe_{2}O_{3}, and of the four acid oxides, CO_{2}, N_{2}O_{5},
P_{2}O_{5}, and SO_{3}, and as the amount of ash-forming substances
in plants is small, the question inevitably arises as to what part
these play in the development of plants. With the existing chemical
data only one answer is possible to this question, and it is still
only a hypothesis. This answer was particularly clearly expressed
by Professor Gustavson of the Petroffsky Agricultural Academy.
Starting from the fact (Chapter XI., Note 55) that a small quantity
of aluminium renders possible or facilitates the reaction of
bromine on hydrocarbons at the ordinary temperature, it is easy to
arrive at the conclusion, which is very probable and in accordance
with many data respecting the reactions of organic compounds, that
the addition of mineral substances to organic compounds lowers the
temperature of reaction and in general facilitates chemical
reactions in plants, and thus aids the conversion of the most
simple nourishing substances into the complex component parts of
the plant organism. The province of chemical reactions proceeding
in organic substances in the presence of a small quantity of
mineral substances has as yet been but little investigated,
although there are already several disconnected data concerning
reactions of this kind, and although a great deal is known with
regard to such reactions among inorganic compounds. The essence of
the matter may be expressed thus--two substances, A and B, do not
react on each other of their own accord, but the addition of a
small quantity of a third particularly active substance, C,
produces the reaction of A on B, because A combines with C, forming
AC, and B reacts on this new compound, which has a different store
of chemical energy, forming the compound AB or its products, and
setting C free again or retaining it.

It may here be remarked that all the mineral substances necessary
for plants (those enumerated at the beginning of the note) are the
highest saline compounds of their elements, that they enter into
the plants as salts, that the lower forms of oxidation of the same
elements (for instance, sulphites and phosphites) are harmful to
plants (poisonous), and that strong solutions of the salts
assimilated by plants (their osmotic pressure being great and
contracting the cells, as De Vries showed, (_see_ Chapter I., Note
19)) not only do not enter into the plants but kill them (poison
them).

[8] Besides which, it will be understood from the preceding paragraph
that the salts of potassium may become exhausted from the soil by
long cultivation, and that there may therefore be cases when the
direct fertilisation by salts of potassium may be profitable. But
manure and animal excrements, ashes, and, in general, nearly all
refuse which may serve for fertilising the soil, contain a
considerable quantity of potassium salts, and therefore, as regards
the natural salts of potassium (Stassfurt), and especially
potassium sulphate, if they often improve the crops, it is in all
probability due to their action on the properties of the soil. The
agriculturist cannot therefore be advised to add potassium salts,
without making special experiments showing the advantage of such a
fertiliser on a given kind of soil and plant.

The animal body also contains potassium compounds, which is
natural, since animals consume plants. For example, milk, and
especially human milk, contains a somewhat considerable quantity of
potassium compounds. Cow's milk, however, does not contain much
potassium salt. Sodium compounds generally predominate in the
bodies of animals. The excrement of animals, and especially of
herbivorous animals, on the contrary, often contains a large
proportion of potassium salts. Thus sheep's dung is rich in them,
and in washing sheep's wool salts of potassium pass into the water.

The ash of tree stems, as the already dormant portion of the plant
(Chapter VIII., Note 1), contains little potash. For the extraction
of potash, which was formerly carried on extensively in the east of
Russia (before the discovery of the Stassfurt salt), the ash of
grasses, and the green portions of potatoes, buckwheat, &c., are
taken and treated with water (lixiviated), the solution is
evaporated, and the residue ignited in order to destroy the organic
matter present in the extract. The residue thus obtained is
composed of raw potash. It is refined by a second dissolution in a
small quantity of water, for the potash itself is very soluble in
water, whilst the impurities are sparingly soluble. The solution
thus obtained is again evaporated, and the residue ignited, and
this potash is then called refined potash, or pearlash. This method
of treatment cannot give chemically pure potassium carbonate. A
certain amount of impurities remain. To obtain chemically pure
potassium carbonate, some other salt of potassium is generally
taken and purified by crystallisation. Potassium carbonate
crystallises with difficulty, and it cannot therefore be purified
by this means, whilst other salts, such as the tartrate, acid
carbonate, sulphate, or nitrate, &c., crystallise easily and may
thus be directly purified. The tartrate is most frequently
employed, since it is prepared in large quantities (as a sediment
from wine) for medicinal use under the name of cream of tartar.
When ignited without access of air, it leaves a mixture of charcoal
and potassium carbonate. The charcoal so obtained being in a
finely-divided condition, the mixture (called 'black flux'), is
sometimes used for reducing metals from their oxides with the aid
of heat. A certain quantity of nitre is added to burn the charcoal
formed by heating the cream of tartar. Potassium carbonate thus
prepared is further purified by converting it into the acid salt,
by passing a current of carbonic anhydride through a strong
solution. KHCO_{3} is then formed, which is less soluble than the
normal salt (as is also the case with the corresponding sodium
salts), and therefore crystals of the acid salt separate from the
solution on cooling. When ignited, they part with their water and
carbonic anhydride, and pure potassium carbonate remains behind.
The physical properties of potassium carbonate distinguish it
sufficiently from sodium carbonate; it is obtained from solutions
as a powdery white mass, having an alkaline taste and reaction,
and, as a rule, shows only traces of crystallisation. It also
attracts the moisture of the air with great energy. The crystals do
not contain water, but absorb it from the air, deliquescing into a
saturated solution. It melts at a red heat (1045°), and at a still
higher temperature is even converted into vapour, as has been
observed at glass works where it is employed. It is very soluble.
At the ordinary temperature, water dissolves an equal weight of the
salt. Crystals containing two equivalents of water separate from
such a saturated solution when strongly cooled (Morel obtained
K_{2}CO_{3}3H_{2}O in well-formed crystals at +10°). There is no
necessity to describe its reactions, because they are all analogous
to those of sodium carbonate. When manufactured sodium carbonate
was but little known, the consumption of potassium carbonate was
very considerable, and even now washing soda is frequently replaced
for household purposes by 'lye'--_i.e._ an aqueous solution
obtained from ashes. It contains potassium carbonate, which acts
like the sodium salt in washing tissues, linen, &c.

A mixture of potassium and sodium carbonates fuses with much
greater ease than the separate salts, and a mixture of their
solutions gives well-crystallised salts--for instance (Marguerite's
salt), K_{2}CO_{3},6H_{2}O,2Na_{2}CO_{3},6H_{2}O. Crystallisation
also occurs in other multiple proportions of K and Na (in the above
case 1 : 2, but 1 : 1 and 1 : 3 are known), and always with 6 mol.
H_{2}O. This is evidently a combination _by similarity_, as in
alloys, solutions, &c.

[8 bis] About 25,000 tons of potash annually are now prepared from KCl
by this method at Stassfurt.

[9] _Potassium sulphate_, K_{2}SO_{4}, crystallises from its solutions
in an anhydrous condition, in which respect it differs from the
corresponding sodium salt, just as potassium carbonate differs from
sodium carbonate. In general, it must be observed that the majority
of sodium salts combine more easily with water of crystallisation
than the potassium salts. The solubility of _potassium sulphate_
does not show the same peculiarities as that of sodium sulphate,
because it does not combine with water of crystallisation; at the
ordinary temperature 100 parts of water dissolve about 10 parts of
the salt, at 0° 8·3 parts, and at 100° about 26 parts. _The acid
sulphate_, KHSO_{4}, obtained easily by heating crystals of the
normal salt with sulphuric acid, is frequently employed in chemical
practice. On heating the mixture of acid and salt, fumes of
sulphuric acid are at first given off; when they cease to be
evolved, the acid salt is contained in the residue. At a higher
temperature (of above 600°) the acid salt parts with all the acid
contained in it, the normal salt being re-formed. The definite
composition of this acid salt, and the ease with which it
decomposes, render it exceedingly valuable for certain chemical
transformations accomplished by means of sulphuric acid at a high
temperature, because it is possible to take, in the form of this
salt, a strictly definite quantity of sulphuric acid, and to cause
it to act on a given substance at a high temperature, which it is
often necessary to do, more especially in chemical analysis. In
this case, the acid salt acts in exactly the same manner as
sulphuric acid itself, but the latter is inefficient at
temperatures above 400°, because it all evaporates, while at that
temperature the acid salt still remains in a fused state, and acts
with the elements of sulphuric acid on the substance taken. Hence
by its means the boiling-point of sulphuric acid is raised. Thus
the acid potassium sulphate is employed, where for conversion of
certain oxides, such as those of iron, aluminium, and chromium,
into salts, a high temperature is required.

Weber, by heating potassium sulphate with an excess of sulphuric
acid at 100°, observed the formation of a lower stratum, which was
found to contain a definite compound containing eight equivalents
of SO_{3} per equivalent of K_{2}O. The salts of rubidium, cæsium,
and thallium give a similar result, but those of sodium and lithium
do not. (_See_ Note 1.)

[10] The _bromide_ and _iodide_ of potassium are used, like the
corresponding sodium compounds, in medicine and photography.
Potassium iodide is easily obtained in a pure state by saturating
a solution of hydriodic acid with caustic potash. In practice,
however, this method is rarely had recourse to, other more simple
processes being employed although they do not give so pure a
product. They aim at the direct formation of hydriodic acid in the
liquid in the presence of potassium hydroxide or carbonate. Thus
iodine is thrown into a solution of pure potash, and hydrogen
sulphide passed through the mixture, the iodine being thus
converted into hydriodic acid. Or a solution is prepared from
phosphorus, iodine, and water, containing hydriodic and phosphoric
acid; lime is then added to this solution, when calcium iodide is
obtained in solution, and calcium phosphate as a precipitate. The
solution of calcium iodide gives, with potassium carbonate,
insoluble calcium carbonate and a solution of potassium iodide. If
iodine is added to a slightly-heated solution of caustic potash
(free from carbonate--that is, freshly prepared), so long as the
solution is not coloured from the presence of an excess of iodine,
there is formed (as in the action of chlorine on a solution of
caustic potash) a mixture of potassium iodide and iodate. On
evaporating the solution thus obtained and igniting the residue,
the iodate is destroyed and converted into iodide, the oxygen
being disengaged, and potassium iodide only is left behind. On
dissolving the residue in water and then evaporating, cubical
crystals of the anhydrous salt are obtained, which are soluble in
water and alcohol, and on fusion give an alkaline reaction, owing
to the fact that when ignited a portion of the salt decomposes,
forming potassium oxide. The neutral salt may be obtained by
adding hydriodic acid to this alkaline salt until it gives an acid
reaction. It is best to add some finely-divided charcoal to the
mixture of iodate and iodide before igniting it, as this
facilitates the evolution of the oxygen from the iodate. The
iodate may also be converted into iodide by the action of certain
reducing agents, such as zinc amalgam, which when boiled with a
solution containing an iodate converts it into iodide. Potassium
iodide may also be prepared by mixing a solution of ferrous iodide
(it is best if the solution contain an excess of iodine) and
potassium carbonate, in which case ferrous carbonate FeCO_{3}, is
precipitated (with an excess of iodine the precipitate is
granular, and contains a compound of the suboxide and oxide of
iron), while potassium iodide remains in solution. Ferrous iodide,
FeI_{2}, is obtained by the direct action of iodine on iron in
water. Potassium iodide considerably lowers the temperature (by
24°), when it dissolves in water, 100 parts of the salt dissolve
in 73·5 parts of water at 12·5°, in 70 parts at 18°, whilst the
saturated solution which boils at 120° contains 100 parts of salt
per 45 parts of water. Solutions of potassium iodide dissolve a
considerable amount of iodine; strong solutions even dissolving as
much or more iodine than they contain as potassium iodide (_see_
Note 3 bis and Chapter XI., Note 64).

[11] Caustic potash is not only formed by the action of lime on dilute
solutions of potassium carbonate (as sodium hydroxide is prepared
from sodium carbonate), but by igniting potassium nitrate with
finely-divided copper (_see_ Note 15), and also by mixing
solutions of potassium sulphate (or even of alum, KAlS_{2}O_{8})
and barium hydroxide, BaH_{2}O_{2}. It is sometimes purified by
dissolving it in alcohol (the impurities, for example, potassium
sulphate and carbonate, are not dissolved) and then evaporating
the alcohol.

The specific gravity of potassium hydroxide is 2·04, but that of
its solutions (see Chapter XII., Note 18) at 15° S = 9,992 +
90·4_p_ + 0·28_p_^2 (here _p_^2 is +, and for sodium hydroxide it
is-). Strong solutions, when cooled, yield a crystallo-hydrate,
KHO,4H_{2}O, which dissolves in water, producing cold (like
2NaHO,7H_{2}O), whilst potassium hydroxide in solution develops a
considerable amount of heat.

_Potassium cyanide_, which presents in its chemical relations a certain analogy with the halogen salts of potassium, is not only formed according to the equation, KHO + HCN = H_{2}O + KCN, but also whenever a nitrogenous carbon compound--for instance, animal matter--is heated in the presence of metallic potassium, or of a compound of potassium, and even when a mixture of potash and carbon is heated in a stream of nitrogen. Potassium cyanide is obtained from yellow prussiate, which has been already mentioned in Chapter IX., and whose preparation on a large scale will be described in Chapter XXII. If the yellow prussiate be ground to a powder and dried, so that it loses its water of crystallisation, it then melts at a red heat, and decomposes into carbide of iron, nitrogen, and potassium cyanide, FeK_{4}C_{6}N_{6} = 4KCN + FeC_{2} + N_{2}. After the decomposition it is found that the yellow salt has been converted into a white mass of potassium cyanide. The carbide of iron formed collects at the bottom of the vessel. If the mass thus obtained be treated with water, the potassium cyanide is partially decomposed by the water, but if it be treated with alcohol, then the cyanide is dissolved, and on cooling separates in a crystalline form.[12] A solution of potassium cyanide has a powerfully alkaline reaction, a smell like that of bitter almonds, peculiar to prussic acid, and acts as a most powerful poison. Although exceedingly stable in a fused state, potassium cyanide easily changes when in solution. Prussic acid is so very feebly energetic that even water decomposes potassium cyanide. A solution of the salt, even without access of air, easily turns brown and decomposes, and when heated evolves ammonia and forms potassium formate; this is easily comprehensible from the representation of the cyanogen compounds which was developed in Chapter IX., KCN + 2H_{2}O = CHKO_{2} + NH_{3}. Furthermore, as carbonic anhydride acts on potassium cyanide with evolution of prussic acid, and as potassium cyanate, which is also unstable, is formed by the action of air, it will be easily seen that solutions of potassium cyanide are very unstable. Potassium cyanide, containing as it does carbon and potassium, is a substance which can act in a very vigorously reducing manner, especially when fused; it is therefore used as a powerful reducing agent at a red heat.[13] The property of potassium cyanide of giving double salts with other cyanides is very clearly shown by the fact that many metals dissolve in a solution of potassium cyanide, with the evolution of hydrogen. For example, iron, copper, and zinc act in this manner. Thus--

4KCN + 2H_{2}O + Zn = K_{2}ZnC_{4}N_{4} + 2KHO + H_{2}

[12] When the yellow prussiate is heated to redness, all the cyanogen
which was in combination with the iron is decomposed into
nitrogen, which is evolved as gas, and carbon, which combines with
the iron. In order to avoid this, potassium carbonate is added to
the yellow prussiate while it is being fused. A mixture of 8 parts
of anhydrous yellow prussiate and 3 parts of pure potassium
carbonate is generally taken. Double decomposition then takes
place, resulting in the formation of ferrous carbonate and
potassium cyanide. But by this method, as by the first, a pure
salt is not obtained, because a portion of the potassium cyanide
is oxidised at the expense of the iron carbonate and forms
potassium cyanate, FeCO_{3} + KCN = CO_{2} + Fe + KCNO; and the
potassium cyanide very easily forms oxide, which acts on the sides
of the vessel in which the mixture is heated (to avoid this iron
vessels should be used). By adding one part of charcoal powder to
the mixture of 8 parts of anhydrous yellow prussiate and 3 parts
of potassium carbonate a mass is obtained which is free from
cyanate, because the carbon absorbs the oxygen, but in that case
it is impossible to obtain a colourless potassium cyanide by
simple fusion, although this may be easily done by dissolving it
in alcohol. Cyanide of potassium may also be obtained from
potassium thiocyanate, which is formed from ammonium thiocyanate
obtained by the action of ammonia upon bisulphide of carbon (_see_
works upon Organic Chemistry). Potassium cyanide is now prepared
in large quantities from yellow prussiate for gilding and
silvering. When fused in large quantities the action of the oxygen
of the air is limited, and with great care the operation may be
successfully conducted, and therefore, on a large scale, very pure
salt is sometimes obtained. When slowly cooled, the fused salt
separates in cubical crystals like potassium chloride.

Pure KCN is obtained by passing CNH gas into an alcoholic solution
of KHO. The large amount of potassium cyanide which is now
required for the extraction of gold from its ores, is being
replaced by a mixture (Rossler and Gasslaker, 1892) of KCN and
NaCN, prepared by heating powdered and dried yellow prussiate with
metallic sodium: K_{4}Fe(CN)_{6} + 2Na = 4KCN + 2NaCN + Fe. This
method offers two advantages over the above methods: (1) the whole
of the cyanide is obtained, and does not decompose with the
formation of N_{2}; and (2) no cyanates are formed, as is the case
when carbonate of potash is heated with the prussiate.

[13] A considerable quantity of potassium cyanide is used in the arts,
more particularly for the preparation of metallic solutions which
are decomposed by the action of a galvanic current; thus it is
very frequently employed in electro-silvering and gilding. An
alkaline solution is prepared, which is moderately stable owing to
the fact that potassium cyanide in the form of certain double
salts--that is, combined with other cyanides--is far more stable
than when alone (yellow prussiate, which contains potassium
cyanide in combination with ferrous cyanide, is an example of
this).

Gold and silver are soluble in potassium cyanide in the presence of air, in which case the hydrogen, which would otherwise be evolved in the reaction, combines with the oxygen of the air, forming water (Eissler, MacLaurin, 1893), for example, 4Au + 4KCN + O + H_{2}O = 2AuKC_{2}N_{2} + 2KHO, which is taken advantage of for extracting gold from its ores (Chapter XXIV.).[13 bis] Platinum, mercury, and tin are not dissolved in a solution of potassium cyanide, even with access of air.

[13 bis] A dilute solution of KCN is taken, not containing more than
1 per cent. KCN. MacLaurin explains this by the fact that strong
solutions dissolve gold less rapidly, owing to their dissolving
less air, whose oxygen is necessary for the reaction.

_Potassium nitrate_, or common _nitre_ or _saltpetre_, KNO_{3}, is chiefly used as a component part of gunpowder, in which it cannot be replaced by the sodium salt, because the latter is deliquescent. It is necessary that the nitre in gunpowder should be perfectly pure, as even small traces of sodium, magnesium, and calcium salts, especially chlorides, render the nitre and the gunpowder capable of attracting moisture. Nitre may easily be obtained pure, owing to its great disposition to form crystals both large and small, which aids its separation from other salts. The considerable differences between the solubility of nitre at different temperatures aids this crystallisation. A solution of nitre saturated at its boiling point (116°) contains 335 parts of nitre to 100 parts of water, whilst at the ordinary temperature--for instance, 20°--the solution is only able to retain 32 parts of the salt. Therefore, in the preparation and refining of nitre, its solution, saturated at the boiling point, is cooled, and nearly all the nitre is obtained in the form of crystals. If the solution be quietly and slowly cooled in large quantities then large crystals are formed, but if it be rapidly cooled and agitated then small crystals are obtained. In this manner, if not all, at all events the majority, of the impurities present in small quantities remain in the mother liquor. If an unsaturated solution of nitre be rapidly cooled, so as to prevent the formation of large crystals (in whose crevices the mother liquor, together with the impurities, would remain), the very minute crystals of nitre known as saltpetre flour are obtained.

Common nitre occurs in nature, but only in small quantities in admixture with other nitrates, and especially with sodium, magnesium, and calcium nitrates. Such a mixture of salts of nitric acid is formed in nature in fertile earth, and in those localities where, as in _the soil_, nitrogenous organic remains are decomposed in the presence of alkalis or alkaline bases with free access of air. This method of the formation of nitrates requires moisture, besides the free access of air, and takes place principally during warm weather.[14] In warm countries, and in temperate climates during the summer months, fertile soils produce a small quantity of nitre. In this respect India is especially known as affording a considerable supply of nitre extracted from the soil. The nitre-bearing soil after the rainy season sometimes becomes covered during the summer with crystals of nitre, formed by the evaporation of the water in which it was previously dissolved. This soil is collected, subjected to repeated lixiviations, and treated for nitre as will be presently described. In temperate climates nitrates are obtained from the lime rubbish of demolished buildings which have stood for many years, and especially from those portions which have been in contact with the ground. The conditions there are very favourable for the formation of nitre, because the lime used as a cement in buildings contains the base necessary for the formation of nitrates, while the excrement, urine, and animal refuse are sources of nitrogen. By the methodical lixiviation of this kind of rubbish a solution of nitrogenous salts is formed similar to that obtained by the lixiviation of fertile soil. A similar solution is also obtained by the lixiviation of the so-called _nitre plantations_. They are composed of manure interlaid with brushwood, and strewn over with ashes, lime, and other alkaline rubbish. These nitre plantations are set up in those localities where the manure is not required for the fertilisation of the soil, as, for example, in the south-eastern 'black earth' Governments of Russia. The same process of oxidation of nitrogenous matter freely exposed to air and moisture during the warm season in the presence of alkalis takes place in nitre plantations as in fertile soil and in the walls of buildings. From all these sources there is obtained a solution containing various salts of nitric acid mixed with soluble organic matter. The simplest method of treating this impure solution of nitre is to add a solution of potassium carbonate, or to simply treat it with ashes containing this substance. The potassium carbonate enters into double decomposition with the calcium and magnesium salts, forming insoluble carbonates of these bases and leaving the nitre in solution. Thus, for instance, K_{2}CO_{3} + Ca(NO_{3})_{2} = 2KNO_3 + CaCO_{3}. Both calcium and magnesium carbonates are insoluble, and therefore after treatment with potassium carbonate the solution no longer contains salts of these metals but only the salts of sodium and potassium together with organic matter. The latter partially separates on heating in an insoluble form, and is entirely destroyed by heating the nitre to a low red heat. The nitre thus obtained is easily purified by repeated crystallisation. The greater part of the nitre used for making gunpowder is now obtained from the sodium salt _Chili saltpetre_ or _cubic nitre_, which occurs in nature, as already mentioned. The conversion of this salt into common nitre is also carried on by means of a double decomposition. This is done either by adding potassium carbonate (when, on mixing the strong and hot solutions, sodium carbonate is directly obtained as a precipitate), or, as is now most frequent, potassium chloride. When a mixture of strong solutions of potassium chloride and sodium nitrate is evaporated, sodium chloride first separates, because this salt, which is formed by the double decomposition KCl + NaNO_{3} = KNO_{3} + NaCl, is almost equally soluble in hot and cold water; on cooling, therefore, a large amount of potassium nitrate separates from the saturated solution, while the sodium chloride remains dissolved. The nitre is ultimately purified by recrystallisation and by washing with a saturated solution of nitre, which cannot dissolve a further quantity of nitre but only the impurities.

[14] Besides which Schloesing and Müntz, by employing similar methods
to Pasteur, showed that the formation of nitre in the
decomposition of nitrogenous substances is accomplished by the aid
of peculiar micro-organisms (ferments), without which the
simultaneous action of the other necessary conditions (alkalis,
moisture, a temperature of 37°, air, and nitrogenous substances)
cannot give nitre.

Nitre is a colourless salt having a peculiar cool taste. It crystallises easily in long striated six-sided rhombic prisms terminating in rhombic pyramids. Its crystals (sp. gr. 1·93) do not contain water, but their cavities generally contain a certain quantity of the solution from which they have crystallised. For this reason in refining nitre, the production of large crystals is prevented, _saltpetre flour_ being prepared. At a low red heat (339°) nitre melts to a colourless liquid.[14 bis] Potassium nitrate at the ordinary temperature and in a solid form is inactive and stable, but _at a high temperature_ it acts as a powerful _oxidising agent_, giving up a considerable amount of oxygen to substances mixed with it.[15] When thrown on to incandescent charcoal it brings about its rapid combustion, and a mechanical mixture of powdered charcoal and nitre ignites when brought into contact with a red-hot substance, and continues to burn by itself. In this action, nitrogen is evolved, and the oxygen oxidises the charcoal, in consequence of which potassium carbonate and carbonic anhydride are formed: 4KNO_{3} + 5C = 2K_{2}CO_{3} + 3CO_{2} + 2N_{2}. This phenomenon depends on the fact that oxygen in combining with carbon evolves more heat than it does in combining with nitrogen. Hence, when once the combustion has been started at the expense of the nitre, it is able to go on without requiring the aid of external heat. A similar oxidation or combustion at the expense of the contained oxygen takes place when nitre is heated with different combustible substances. If a mixture of sulphur and nitre be thrown upon a red-hot surface, the sulphur burns, forming potassium sulphate and sulphurous anhydride. In this case, also, the nitrogen of the nitre is evolved as gas: 2KNO_{3} + 2S = K_{2}SO_{4} + N_{2} + SO_{2}. A similar phenomenon occurs when nitre is heated with many metals. The oxidation of those metals which are able to form acid oxides with an excess of oxygen is especially remarkable. In this case they remain in combination with potassium oxide as potassium salts. Manganese, antimony, arsenic, iron, chromium, &c. are instances of this kind. These elements, like carbon and sulphur, displace free nitrogen. The lower oxides of these metals when fused with nitre pass into the higher oxides. Organic substances are also oxidised when heated with nitre--that is, they burn at the expense of the nitre. It will be readily understood from this that nitre is frequently used in practical chemistry and the arts as an oxidising agent at high temperatures. Its application in _gunpowder_ is based on this property; gunpowder consists of a mechanical mixture of finely-ground sulphur, nitre, and charcoal. The relative proportion of these substances varies according to the destination of the powder and to the kind of charcoal employed (a friable, incompletely-burnt charcoal, containing therefore hydrogen and oxygen, is employed). Gases are formed in its combustion, chiefly nitrogen and carbonic anhydride, which create a considerable pressure if their escape be in any way impeded. This action of gunpowder may be expressed by the equation: 2KNO_{3} + 3C + S = K_{2}S + 3CO_{2} + N_{2}.

[14 bis] Before fusing, the crystals of potassium nitrate change their
form, and take the same form as sodium nitrate--that is, they
change into rhombohedra. Nitre crystallises from hot solutions,
and in general under the influence of a rise of temperature, in a
different form from that given at the ordinary or lower
temperatures. Fused nitre solidifies to a radiated crystalline
mass; but it does not exhibit this structure if metallic chlorides
be present, so that this method may be taken advantage of to
determine the degree of purity of nitre.

Carnelley and Thomson (1888) determined the fusing point of
mixtures of potassium and sodium nitrates. The first salt fuses at
339° and the second at 316°, and if _p_ be the percentage amount
of potassium nitrate, then the results obtained were--

_p_ = 10 20 30 40 50 60 70 80 90
298° 283° 268° 242° 231° 231° 242° 284° 306°

which confirms Shaffgotsch's observation (1857) that the lowest
fusing point (about 231°) is given by mixing molecular quantities
(_p_ = 54·3) of the salts--that is, in the formation of the alloy,
KNO_{3},NaNO_{3}.

A somewhat similar result was discovered by the same observers for
the solubility of mixtures of these salts at 20° in 100 parts of
water. Thus, if _p_ be the weight of potassium nitrate mixed with
100-_p_ parts by weight of sodium nitrate taken for solution, and
_c_ be the quantity of the mixed salts which dissolves in 100, the
solubility of sodium nitrate being 85, and of potassium nitrate
34, parts in 100 parts of water, then--

_p_ = 10 20 30 40 50 60 70 80 90
_c_ = 110 136 136 138 106 81 73 54 41

The maximum solubility proved not to correspond with the most
fusible mixture, but to one much richer in sodium nitrate.

Both these phenomena show that in homogeneous liquid mixtures the
chemical forces that act between substances are the same as those
that determine the molecular weights of substances, even when the
mixture consists of such analogous substances as potassium and
sodium nitrates, between which there is no direct chemical
interchange. It is instructive to note also that the maximum
solubility does not correspond with the minimum fusing point,
which naturally depends on the fact that in solution a third
substance, namely water, plays a part, although an attraction
between the salts, like that which exists between sodium and
potassium carbonates (Note 8), also partially acts.

[15] Fused nitre, with a further rise of temperature, disengages oxygen
and then nitrogen. The nitrite KNO_{2} is first formed and then
potassium oxide. The admixture of certain metals--for example, of
finely-divided copper--aids the last decomposition. The oxygen in
this case naturally passes over to the metal.

It is found by this equation that gunpowder should contain thirty-six parts of charcoal (13·3 p.c.), and thirty-two parts (11·9 p.c.) of sulphur, to 202 parts (74·8 p.c.) of nitre, which is very near to its actual composition.[16]

[16] In China, where the manufacture of gunpowder has long been carried
on, 75·7 parts of nitre, 14·4 of charcoal, and 9·9 of sulphur are
used. Ordinary powder for sporting purposes contains 80 parts of
nitre, 12 of charcoal, and 8 of sulphur, whilst the gunpowder used
in heavy ordnance contains 75 of nitre, 15 of charcoal, and 10 of
sulphur. Gunpowder explodes when heated to 300°, when struck, or
by contact with a spark. A compact or finely-divided mass of
gunpowder burns slowly and has but little disruptive action,
because it burns gradually. To act properly the gunpowder must
have a definite rate of combustion, so that the pressure should
increase during the passage of the projectile along the barrel of
the fire-arm. This is done by making the powder in large granules
or in the shape of six-sided prisms with holes through them
(prismatic powder).

The products of combustion are of two kinds: (1) gases which
produce the pressure and are the cause of the dynamical action of
gunpowder, and (2) a solid residue, usually of a black colour
owing to its containing unburnt particles of charcoal. Besides
charcoal, the residue generally contains potassium sulphide,
K_{2}S, and a whole series of other salts--for instance, carbonate
and sulphate. It is apparent from this that the combustion of
gunpowder is not so simple as it appears to be from the above
formula, and hence the weight of the residue is also greater than
indicated by that formula. According to the formula, 270 parts of
gunpowder give 110 parts of residue--that is, 100 parts of powder
give 37·4 parts of residue, K_{2}S, whilst in reality the weight
of the residue varies from 40 p.c. to 70 p.c. (generally 52 p.c.).
This difference depends on the fact that so much oxygen (of the
nitre) remains in the residue, and it is evident that if the
residue varies the composition of the gases evolved by the powder
will vary also, and therefore the entire process will be different
in different cases. The difference in the composition of the gases
and residue depends, as the researches of Gay-Lussac, Shishkoff
and Bunsen, Nobel and Abel, Federoff, Debus, &c., show, on the
conditions under which the combustion of the powder proceeds. When
gunpowder burns in an open space, the gaseous products which are
formed do not remain in contact with the residue, and then a
considerable portion of the charcoal entering into the composition
of the powder remains unburnt, because the charcoal burns after
the sulphur at the expense of the oxygen of the nitre. In this
extreme case the commencement of the combustion of the gunpowder
may be expressed by the equation, 2KNO_{3} + 3C + S = 2C +
K_{2}SO_{4} + CO_{2} + N_{2}. The residue in a blank cartridge
often consists of a mixture of C, K_{2}SO_{4}, K_{2}CO_{3}, and
K_{2}S_{2}O_{3}. If the combustion of the gunpowder be impeded--if
it take place in a cartridge in the barrel of a gun--the quantity
of potassium sulphate will first be diminished, then the amount of
sulphite, whilst the amount of carbonic anhydride in the gases and
the amount of potassium sulphide in the residue will increase. The
quantity of charcoal entering into the action will then be also
increased, and hence the amount in the residue will decrease.
Under these circumstances the weight of the residue will be
less--for example, 4K_{2}CO_{3} + 4S = K_{2}SO_{4} + 3K_{2}S +
4CO_{2}. Besides which, carbonic oxide has been found in the
gases, and potassium bisulphide, K_{2}S_{2}, in the residue of
gunpowder. The amount of potassium sulphide, K_{2}S, increases
with the completeness of the combustion, and is formed in the
residue at the expense of the potassium sulphite. In recent times
the knowledge of the action of gunpowder and other explosives has
made much progress, and has developed into a vast province of
artillery science, which, guided by the discoveries of chemistry,

has worked out a 'smokeless powder' which burns without leaving a
residue, and does not therefore give any 'powder smoke' (to hinder
the rapidity of firing and aiming), and at the same time
disengages a greater volume of gas and consequently gives (under
proper conditions of combustion) the possibility of communicating
to the charge a greater initial velocity, and therefore greater
distance, force, and accuracy of aim. Such 'smokeless powder' is
prepared either from the varieties of nitrocellulose (Chapter VI.,
Note 37) or from a mixture of them with nitro-glycerine (_ibid_).
In burning they give, besides steam and nitrogen, generally a
large amount of oxide of carbon (this is a very serious drawback
in all the present forms of smokeless powder, because carbonic
oxide is poisonous), and also CO_{2}, H_{2}, &c.

_Metallic potassium_ was obtained like sodium; first by the action of a galvanic current, then by reduction of the hydroxide by means of metallic iron, and lastly, by the action of charcoal on the carbonate at a high temperature. The behaviour of metallic potassium differs, however, from that of sodium, because it easily combines with carbonic oxide, forming an explosive and inflammable mass.[17]

[17] The substances obtained in this case are mentioned in Chapter IX.,
Note 31.

Potassium is quite as volatile as sodium, if not more so. At the ordinary temperature potassium is even softer than sodium; its freshly-cut surfaces present a whiter colour than sodium, but, like the latter, and with even greater ease, it oxidises in moist air. It is brittle at low temperatures, but is quite soft at 25°, and melts at 58°. At a low red heat (667°, Perkin) it distils without change, forming a green vapour, whose density,[18] according to A. Scott (1887), is equal to 19 (if that of hydrogen = 1). This shows that the molecule of potassium (like that of sodium, mercury, and zinc) contains but one atom. This is also the case with many other metals, judging by recent researches.[19] The specific gravity of potassium at 15° is 0·87, and is therefore less than that of sodium, as is also the case with all its compounds.[20] Potassium decomposes water with great ease at the ordinary temperature, evolving 45,000 heat units per atomic weight in grams. The heat evolved is sufficient to inflame the hydrogen, the flame being coloured violet from the presence of particles of potassium.[21]

[18] A. Scott (1887) determined the vapour densities of many of the
alkali elements and their compounds in a platinum vessel heated in
a furnace and previously filled with nitrogen. But these, the
first data concerning a subject of great importance, have not yet
been sufficiently fully described, nor have they received as much
attention as could be desired. Taking the density of hydrogen as
unity, Scott found the vapour densities of the following
substances to be--

Na 12·75 (11·5). KI 92 (84).
K 19 (19·5). RbCl 70 (60).
CsCl 89·5 (84·2). CsI 133 (130).
FeCl_{3} 68. AgCl 80 (71·7).

In brackets are given the densities corresponding with the
formulæ, according to Avogadro-Gerhardt's law. This figure is not
given for FeCl_{3}, because in all probability under these
conditions (the temperature at which it was determined) a portion
of the FeCl_{3} was decomposed. If it was not decomposed, then a
density 81 would correspond with the formula FeCl_{3}, and if the
decomposition were Fe_{2}Cl_{6} = 2FeCl_{2} + Cl_{2}, then the
density should be 54. With regard to the silver chloride, there is
reason to think that the platinum decomposed this salt. The
majority of Scott's results so closely correspond with the formulæ
that a better concord cannot be expected in such determinations.
V. Meyer (1887) gives 93 as the density of KI.

[19] The molecules of non-metals are more complex--for instance, H_{2},
O_{3}, Cl_{2}, &c. But arsenic, whose superficial appearance
recalls that of metals, but whose chemical properties approach
more nearly to the non-metals, has a complex molecule containing
As_{4}.

[20] As the atomic weight of potassium is greater than that of sodium,
the volumes of the molecules, or the quotients of the molecular
weight by the specific gravity, for potassium compounds are
greater than those of sodium compounds, because both the
denominator and numerator of the fraction increase. We cite for
comparison the volumes of the corresponding compounds--

Na 24 NaHO 18 NaCl 28 NaNO_{3} 37 Na_{2}SO_{4} 54
K 45 KHO 27 KCl 39 KNO_{3} 48 K_{2}SO_{4} 66

[21] The same precautions must be taken in decomposing water by
potassium as have to be observed with sodium (Chapter II., Note
8).

It must be observed that potassium decomposes carbonic anhydride
and carbonic oxide when heated, the carbon being liberated and the
oxygen taken up by the metal, whilst on the other hand charcoal
takes up oxygen from potassium, as is seen from the preparation of
potassium by heating potash with charcoal, hence the reaction
K_{2}O + C = K_{2} + CO is reversible and the relation is the same
in this case as between hydrogen and zinc.

With regard to the relation of potassium to hydrogen and oxygen, it is closely analogous to sodium in this respect. Thus, with hydrogen it forms potassium hydride, K_{2}H (between 200° and 411°), and with oxygen it gives a suboxide K_{4}O, oxide K_{2}O, and peroxide, only more oxygen enters into the composition of the latter than in sodium peroxide; potassium peroxide contains KO_{2}, but it is probable that in the combustion of potassium an oxide KO is also formed. Potassium, like sodium, is soluble in mercury.[22] In a word, the relation between sodium and potassium is as close as that between chlorine and bromine, or, better still, between fluorine and chlorine, as the atomic weight of sodium, 23, is as much greater than that of fluorine, 19, as that of potassium, 39, is greater than that of chlorine, 35·5.