Chapter XX: Sulphur, Selenium, and Tellurium (1)

The acid character of the higher oxides RO_{3} of the elements of group VI. is still more clearly defined than that of the higher oxides of the preceding groups, whilst feeble basic properties only appear in the oxides RO_{3} of the elements of the even series, and then only for those elements having a high atomic weight--that is, under those two conditions in which, as a rule, the basic characters increase. Even the lower types RO_{2} and R_{2}O_{3}, &c., formed by the elements of group VI., are acid anhydrides in the uneven series, and only those of the elements of the even series have the properties of peroxides or even of bases.

_Sulphur_ is the typical representative of group VI., both on account of the fact that the acid properties of the group are clearly defined in it, and also because it is more widely distributed in nature than any of the other elements belonging to this group. As an element of the uneven series of group VI., sulphur gives H_{2}S, sulphuretted hydrogen, SO_{3}, sulphuric anhydride, and SO_{2}, sulphurous anhydride. And in all of them we find acid properties--SO_{3} and SO_{2} are anhydrides of acids, and H_{2}S is an acid, although a feeble one. As an element sulphur has all the properties of a true non-metal; it has not a metallic lustre, does not conduct electricity, is a bad conductor of heat, is transparent, and combines directly with metals--in short it has all the properties of the non-metals, like oxygen and chlorine. Furthermore, sulphur exhibits a great qualitative and quantitative _resemblance to oxygen_, especially in the fact that, like oxygen, it combines _with two atoms of hydrogen_, and forms compounds resembling oxides with metals and non-metals. From this point of view sulphur is bivalent, if the halogens are univalent.[1] The chemical character of sulphur is expressed by the fact that it forms a very slightly stable and feebly energetic acid with hydrogen. The salts corresponding with this acid are the sulphides, just as the oxides correspond to water and the chlorides to hydrochloric acid. However, as we shall afterwards see more fully, the sulphides are more analogous to the former than to the latter. But although combining with metals, like oxygen, sulphur also forms chemically stable compounds with oxygen, and this fact impresses a peculiar character on all the relations of this element.[2]

[1] The character of sulphur is very clearly defined in the
organo-metallic compounds. Not to dwell on this vast subject, which
belongs to the province of organic chemistry, I think it will be
sufficient for our purpose to compare the physical properties of
the ethyl compounds of mercury, zinc, sulphur and oxygen. The
composition of all of them is expressed by the general formula
(C_{2}H_{5})_{2}R, where R = Hg, Zn, S, or O. They are all
volatile: mercury ethyl, Hg(C_{2}H_{5})_{2}, boils at 159°, its sp.
gr. is 2·444, molecular volume = 106; zinc ethyl boils at 118°, sp.
gr. 1·882, volume 101; ethyl sulphide, S(C_{2}H_{5})_{2}, boils at
90°, sp. gr. 0·825, volume 107; common ether, or ethyl oxide,
O(C_{2}H_{5})_{2}, boils at 35°, sp. gr. 0·736, volume 101, in
addition to which diethyl itself, (C_{2}H_{5})_{2} = C_{4}H_{10},
boils about 0°, sp. gr. about 0·62, volume about 94. Thus the
substitution of Hg, S, and O scarcely changes the volume,
notwithstanding the difference of the weights; the physical
influence, if one may so express oneself, of these elements, which
are so very different in their atomic weights, is almost alike.

[2] Therefore in former times sulphur was known as an amphid element.
Although the analogy between the compounds of sulphur and oxygen
has been recognised from the very birth of modern chemistry (owing,
amongst other things, to the fact that the oxides and sulphides are
the most widely spread metallic ores in nature), still it has only
been clearly expressed by the periodic system, which places both
these elements in group VI. Here, moreover, stands out that
parallelism which exists between SO_{2} and ozone OO_{2}, between
K_{2}SO_{3} and peroxide of potassium K_{2}O_{4} (Volkovitch in
1893 again drew attention to this parallelism).

Sulphur belongs to the number of those elements which _are very widely distributed in nature_, and occurs both free and combined in various forms. The atmosphere, however, is almost entirely free from compounds of sulphur, although a certain amount of them should be present, if only from the fact that sulphurous anhydride is emitted from the earth in volcanic eruptions, and in the air of cities, where much coal is burnt, since this always contains FeS_{2}. Sea and river water generally contain more or less sulphur in the form of sulphates. The beds of gypsum, sodium sulphate, magnesium sulphate, and the like are formations of undoubtedly aqueous origin. The sulphates contained in the soil are the source of the sulphur found in plants, and are indispensable to their growth. Among vegetable substances, the proteïds always contain from one to two per cent. of sulphur. From plants the albuminous substances, together with their sulphur, pass into the animal organism, and therefore the decomposition of animal matter is accompanied by the odour of sulphuretted hydrogen, as the product into which the sulphur passes in the decomposition of the albuminous substances. Thus a rotten egg emits sulphuretted hydrogen. Sulphur occurs largely in nature, as the various insoluble sulphides of the metals. Iron, copper, zinc, lead, antimony, arsenic, &c., occur in nature combined with sulphur. These _sulphides_ frequently have a metallic lustre, and in the majority of cases occur crystallised, and also very often several sulphides occur combined or mixed together in these crystalline compounds. If they are yellow and have a metallic lustre they are called pyrites. Such are, for example, copper pyrites, CuFeS_{2}, and iron pyrites, FeS_{2}, which is the commonest of all. They are all also known as glances or blendes if they are greyish and have a metallic lustre--for example, zinc blende, lead glance, PbS, antimony glance, Sb_{2}S_{3}, &c. And, lastly, sulphur occurs _native_. It occurs in this form in the most recent geological formations in admixture with limestone and gypsum, and most frequently in the vicinity of active or extinct volcanoes. As the gases of volcanoes contain sulphur compounds--namely, sulphuretted hydrogen and sulphurous anhydride, which by reacting on one another may produce sulphur, which also frequently appears in the craters of volcanoes as a sublimate--it might be imagined that the sulphur was of volcanic origin. But on a nearer acquaintance with its mode of occurrence, and more especially considering its relation to gypsum, CaSO_{4}, and limestone, the present general opinion leads to the conclusion that the 'native' sulphur has been formed by the reduction of the gypsum by organic matter and that its occurrence is only indirectly connected with volcanic agencies. Near Tetush, on the Volga, there are beds containing gypsum, sulphur, and asphalt (mineral tar). In Europe the most important deposits of sulphur are in the south of Sicily from Catania to Girgenti.[3] There are very rich deposits of sulphur in Daghestan near Cherkai and Cherkat in Khyut, near Mount Kanabour-bam, near Petrovsk, and in the Kira Koumski steppes in the Trans-Caspian provinces, which are able to supply the whole of Russia with this mineral. Abundant deposits of sulphur have also been found in Kamtchatka in the neighbourhood of the volcanoes. The method of separation of the sulphur from its earthy impurities is based on the fact that sulphur melts when it is heated. The fusion is carried on at the expense of a portion of the sulphur, which is burnt, so that the remainder may melt and run from the mass of the earth. This is carried on in special furnaces called calcaroni, built up of unhewn stone in the neighbourhood of the mines.[4]

[3] When in Sicily, I found, near Caltanisetta, a specimen of sulphur
with mineral tar. In the same neighbourhood there are naphtha
springs and mud volcanoes. It may be that these substances have
reduced the sulphur from gypsum.

The chief proof in favour of the origin of sulphur from gypsum is
that in treating the deposits for the extraction of the sulphur it
is found that the proportion of sulphur to calcium carbonate never
exceeds that which it would be had they both been derived from
calcium sulphate.

[4] Naturally only those ores of sulphur which contain a considerable
amount of sulphur can be treated by this method. With poor ores it
is necessary to have recourse to distillation or mechanical
treatment in order to separate the sulphur, but its price is so low
that this method in most cases is not profitable.

The sulphur obtained by the above-described method still contains
some impurities, but it is frequently made use of in this form for
many purposes, and especially in considerable quantities for the
manufacture of sulphuric acid, and for strewing over grapes. For
other purposes, and especially in the preparation of gunpowder, a
purer sulphur is required. Sulphur may be purified by distillation.
The crude sulphur is called _rough_, and the distilled sulphur
_refined_. The arrangement given in fig. 86 is employed for
refining sulphur. The rough sulphur is melted in the boiler _d_,
and as it melts it is run through the tube F into an iron retort B
heated by the naked flame of the furnace. Here the sulphur is
converted into vapour, which passes through a wide tube into the
chamber G, surrounded by stone walls and furnished with a
safety-valve S.

Sulphur is purified by distillation in special retorts (see fig. 86) by passing the vapour into a chamber G built of stone. The first portions of the vapour entering into the condensing chamber are condensed straightway from the vapour into a solid state, and form a fine powder known as _flowers of sulphur_.[5] But when the temperature of the receiver attains the melting point of sulphur, it passes into a liquid state and is cast into moulds (like sealing wax), and is then known under the name of _roll sulphur_.[6]

[5] Flowers of sulphur always contain a certain amount of the oxides of
sulphur.

[6] Sulphur may be extracted by various other means. It may be
extracted from iron pyrites, FeS_{2}, which is very widely
distributed in nature. From 100 parts of iron pyrites about half
the sulphur contained, namely, about 25 parts, may be extracted by
heating without the access of air, a lower sulphide of iron, which
is more stable under the action of heat, being left behind. Alkali
waste (Chapter XII.), containing calcium sulphide and gypsum,
CaSO_{4}, may be used for the same purpose, but native sulphur is
so cheap that recourse can only be had to these sources when the
calcium sulphide appears as a worthless by-product. The most simple
process for the extraction of sulphur from alkali waste, in a
chemical sense, consists in evolving sulphuretted hydrogen from the
calcium sulphide by the action of hydrochloric acid. The
sulphuretted hydrogen when burnt gives water and sulphurous
anhydride, which reacts on fresh sulphuretted hydrogen with the
separation of sulphur. The combustion of the sulphuretted hydrogen
may be so conducted that a mixture of 2H_{2}S and SO_{2} is
straightway formed, and this mixture will deposit sulphur (Chapter
XII., Note 14). Gossage and Chance treat alkali waste with carbonic
anhydride, and subject the sulphuretted hydrogen evolved to
incomplete combustion (this is best done by passing a mixture of
sulphuretted hydrogen and air, taken in the requisite proportions,
over red-hot ferric oxide), by which means water and the vapour of
sulphur are formed: H_{2}S + O = H_{2}O + S.

In an uncombined state sulphur exists in _several modifications_, and forms a good example of the facility with which an alteration of properties can take place without a change of composition--that is, as regards the material of a substance. Common sulphur has the well-known yellow colour. This colour fades as the temperature falls, and at -50° sulphur is almost colourless. It is very brittle, so that it may be easily converted into a powder, and it presents a crystalline structure, which, by the way, shows itself in the unequal expansion of lumps of sulphur by heat. Hence when a piece of sulphur is heated by the warmth of the hand, it emits sounds and sometimes cracks, which probably also depends on the bad heat-conducting power of this substance. It is easily obtained in a crystalline form by artificial means, because although insoluble in water it dissolves in carbon bisulphide, and in certain oils.[7] Solutions of sulphur in carbon bisulphide when evaporated at the ordinary temperature yield well-formed transparent crystals of sulphur in the form of rhombic octahedra, in which form it occurs native. The specific gravity of these crystals is 2·045. Fused sulphur, cast into moulds and cooled, has, after being kept a long time, a specific gravity 2·066; almost the same as that of the crystalline sulphur of the above form, which shows that common sulphur is the same as that which crystallises in octahedra. The specific heat of octahedral sulphur is 0·17; it melts at 114°, and forms a bright yellow mobile liquid. On further heating, the fused sulphur undergoes an alteration, which we shall presently describe, first observing that the above octahedral state of sulphur is its most stable form. Sulphur may be kept at the ordinary temperature in this form for an indefinite length of time, and many other modifications of sulphur pass into this form after being left for a certain time at ordinary temperature.

[7] One hundred parts of liquid carbon bisulphide, CS_{2}, dissolve
16·5 parts of sulphur at -11°, 24 parts at 0°, 37 parts at 15°, 46
parts at 22°, and 181 parts at 55°. The saturated solution boils at
55°, whilst pure carbon bisulphide boils at 47°. The solution of
sulphur in carbon bisulphide reduces the temperature, just as in
the solution of salts in water. Thus the solution of 20 parts of
sulphur in 50 parts of carbon bisulphide at 22° lowers the
temperature by 5°; 100 parts of benzene, C_{6}H_{6}, dissolves
0·965 part of sulphur at 26°, and 4·377 parts at 71°; chloroform,
CHCl_{3}, dissolves 1·2 part of sulphur at 22°, and 16·35 parts at
174°.

If sulphur be melted and then slightly cooled, so that it forms a crust on the surface and over the sides of the crucible, while the internal mass remains liquid, then the sulphur takes another crystalline form as it solidifies. This may be seen by breaking the crust, and pouring out the remaining molten sulphur.[8] It is then found that the sides of the crucible are covered with _prismatic crystals_ of the monoclinic system; they have a totally different appearance from the above-described crystals of rhombic sulphur. The prismatic crystals are brown, transparent, and less dense than the crystals of rhombic sulphur, their specific gravity being only 1·93, and their melting point higher--about 120°. These crystals of sulphur cannot be kept at the ordinary temperature, which is indeed evident from the fact that in time they turn yellow; the specific gravity also changes, and they pass completely into the ordinary modification. This is accompanied by a considerable development of heat, so that the temperature of the mass may rise 12°. Thus sulphur is _dimorphous_--that is, it exists in two crystalline forms, and in both forms it has independent physical properties. However, no chemical reactions are known which distinguish the two modifications of sulphur, just as there are none distinguishing aragonite from calcspar.[9]

[8] If the experiment be made in a vessel with a narrow capillary tube,
the sulphur fuses at a lower temperature (occurs, as it were, in a
supersaturated state), and solidifying at 90°, appears in a rhombic
form (Schützenberger).

[9] If sulphur be cautiously melted in a U tube immersed in a salt
bath, and then gradually cooled, it is possible for all the sulphur
to remain liquid at 100°. It will now be in a state of superfusion;
thus also by careful refrigeration water may be obtained in a
liquid state at -10°, and a lump of ice then causes such water to
form ice, and the temperature rises to 0°. If a prismatic crystal
of sulphur be thrown into one branch of the U tube containing the
liquid sulphur at 100°, and an octahedral crystal be thrown into
the other branch, then, as Gernez showed, the sulphur in each
branch will crystallise in the corresponding form, and both forms
are obtained at the same temperature; therefore it is not the
influence of temperature only which causes the molecules of sulphur
to distribute themselves in one or another form, but also the
influence of the crystalline parts already formed. This phenomenon
is essentially analogous to the phenomena of supersaturated
solutions.

If molten sulphur be heated to 158° it loses its mobility and becomes thick and very dark-coloured, so that the crucible in which it is heated may be inverted without the sulphur running out. When heated above this temperature the sulphur again becomes liquid, and at 250° it is very mobile, although it does not acquire its original colour, and at 440° it boils. These modifications in the properties of sulphur depend not only on the variations of temperature, but also on a change of structure. If sulphur, heated to about 350°, be poured in a thin stream into cold water, it does not solidify into a solid mass, but retains its brown colour and _remains soft_, may be stretched out into threads, and is elastic, like guttapercha. But in this soft and ductile state, also, it does not remain for a long time. After the lapse of a certain period this soft transparent sulphur hardens, becomes opaque, passes into the ordinary yellow modification of sulphur, and in so doing develops heat, just as in the conversion of the prismatic into the octahedral variety. The soft sulphur is characterised by the fact that a certain portion of it is insoluble in carbon bisulphide. When soft sulphur is immersed in this liquid, only a portion of common sulphur passes into solution, whilst a certain portion is quite insoluble and remains so for a long time. The maximum proportion of insoluble sulphur is obtained by heating slightly above 170°. It melts at 114°. An exactly similar _insoluble amorphous sulphur_ is obtained in certain reactions in the wet way, when sulphur separates out from solutions. Thus sodium thiosulphate, Na_{2}S_{2}O_{3}, when treated with acids, gives a precipitate of sulphur, which is insoluble in carbon bisulphide. The action of water on sulphur chloride also gives a similar modification of sulphur. Certain sulphides, when treated with nitric acid, also yield sulphur in this form.[10]

[10] A certain amount of insoluble sulphur remains for a long time in
the mass of soft sulphur, changing into the ordinary variety.
Freshly-cooled soft sulphur contains about one-third of insoluble
sulphur, and after the lapse of two years it still contains about
15 p.c. Flowers of sulphur, obtained by the rapid condensation of
sulphur from a state of vapour, also contains a certain amount of
insoluble sulphur. _Rapidly distilled and condensed sulphur_ also
contains some insoluble sulphur. Hence a certain amount of
insoluble sulphur is frequently found in roll sulphur. The action
of light on a solution of sulphur converts a certain portion into
the insoluble modification. Insoluble sulphur is of a lighter
colour than the ordinary variety. It is best prepared by
vaporising sulphur in a stream of carbonic anhydride, hydrochloric
acid, &c., and collecting the vapour in cold water. When condensed
in this manner it is nearly all insoluble in carbon bisulphide. It
then has the form of hollow spheroids, and is therefore lighter
than the common variety: sp. gr. 1·82. An idea of the
modifications taking place in sulphur between 110° and 250° may be
formed from the fact that at 150° liquid sulphur has a coefficient
of expansion of about 0·0005, whilst between 150° and 250° it is
less than 0·0003.

Engel (1891), by decomposing a saturated solution of hyposulphite
of sodium (Note 42) with HCl in the cold (the sulphur is not
precipitated directly in this case), obtained, after shaking up
with chloroform and evaporation, crystals of sulphur (sp. gr.
2·135), which, after several hours, passed into the insoluble (in
CS_{2}) state, and in so doing became opaque, and increased in
volume. But if a mixture of solution of Na_{2}S_{2}O_{3} and HCl
be allowed to stand, it deposits sulphur, which, after sufficient
washing, is able to dissolve in water (like the colloid varieties
of the metallic sulphides, alumina, boron, and silver), but this
colloid _solution of sulphur_ soon deposits sulphur insoluble in
CS_{2}.

When a solution of sulphuretted hydrogen in water is decomposed by
an electric current the sulphur is deposited on the positive pole,
and has therefore an electro-negative character, and this sulphur
is soluble in carbon bisulphide. When a solution of sulphurous
acid is decomposed in the same manner, the sulphur is deposited on
the negative pole, and is therefore electro-positive, and the
sulphur so deposited is insoluble in carbon bisulphide. The
sulphur which is combined with metals must have the properties of
the sulphur contained in sulphuretted hydrogen, whilst the sulphur
combined with chlorine is like that which is combined with oxygen
in sulphurous anhydride. Hence Berthelot recognises the presence
of soluble sulphur in metallic sulphides, and of the insoluble
modification of amorphous sulphur in sulphur chloride. Cloez
showed that the sulphur precipitated from solutions is either
soluble or insoluble, according to whether it separates from an
alkaline or acid solution. If sulphur be melted with a small
quantity of iodine or bromine, then on pouring out the molten mass
it forms amorphous sulphur, which keeps so for a very long time,
and is insoluble, or nearly so, in carbon bisulphide. This is
taken advantage of in casting certain articles in sulphur, which
by this means retain their tenacity for a long time; for example,
the discs of electrical machines.

At temperatures of 440° to 700° the vapour density of sulphur is 6·6 referred to air--_i.e._ about 96 referred to hydrogen. Hence, at these temperatures _the molecule of sulphur contains six atoms_, it has the composition S_{6}. The agreement between the observations of Dumas, Mitscherlich, Bineau, and Deville confirms the accuracy of this result. But in this respect the properties of sulphur were found to be variable. When heated to higher temperatures, that is to say, _above_ 800°, the vapour density of sulphur is found to be one-third of this quantity, _i.e._ about 32 referred to hydrogen. At this temperature _the molecule of sulphur_, like that of hydrogen, oxygen, nitrogen, and chlorine, _contains two atoms_; hence the molecular formula is then S_{2}. This variation in the vapour density of sulphur evidently corresponds with a polymeric modification, and may be likened to the transformation of ozone, O_{3}, into oxygen, O_{2}, or better still, of benzene, C_{6}H_{6}, into acetylene, C_{2}H_{2}.[11]

[11] Here, however, it is very important to remark that both benzene
and acetylene can exist at the ordinary temperature, whilst the
sulphur molecule S_{2} only exists at high temperatures; and if
this sulphur be allowed to cool, it passes first into S_{6} and
then into a liquid state. Were it possible to have sulphur at the
ordinary temperature in both the above modifications, then in all
probability the sulphur in the state S_{2} would present totally
different properties from those which it has in the form S_{6},
just as the properties of gaseous acetylene are far from being
similar to those of liquid benzene. Sulphur, in the form of S_{2},
is probably a substance which boils at a much lower temperature
than the variety with which we are now dealing. Paterno and Nasini
(1888), following the method of depression or fall of the
freezing-point in a benzene solution, found that the molecule of
sulphur in solution contains S_{6}.

One must here call attention to the fact that sulphur, with all
its analogy to oxygen (which also shows itself in its faculty to
give the modification S_{2}), is also able to give a series of
compounds containing more atoms of sulphur than the analogous
oxygen compounds do of oxygen. Thus, for instance, compounds of 5
atoms of sulphur with 1 atom of barium, BaS_{5}, are known,
whereas with oxygen only BaO_{2} is known. On every side one
cannot but see in sulphur a faculty for the union of a greater
number of atoms than with oxygen. With oxygen the form of ozone,
O_{3}, is very unstable, the stable form is O_{2}; whilst with
sulphur S_{6} is the stable form, and S_{2} is exceedingly
unstable. Furthermore, it is remarkable that sulphur gives a
higher degree of oxidation, H_{2}SO_{4}, corresponding, as it
were, with its complex composition, if we suppose that in S_{6}
four atoms of sulphur are replaced by oxygen and one by two atoms
of hydrogen. The formulæ of its compounds, K_{2}SO_{4},
K_{2}S_{2}O_{3}, K_{2}S_{5}, BaS_{5}, and many others, have no
analogues among the compounds of oxygen. They all correspond with
the form S_{6} (one portion of the sulphur being replaced by
oxygen and another by metals), which is not attained by oxygen. In
this faculty of sulphur to hold many atoms of other substances the
same forces appear which cause many atoms of sulphur to form one
complex molecule.

_In its faculty for combination_, sulphur most closely resembles oxygen and chlorine; like them, it combines with nearly all elements, with the development of heat and light, forming sulphur compounds, but as a rule this only takes place at a high temperature. At the ordinary temperature it does not enter into reactions, owing, amongst other things, to the fact that it is a solid. In a molten state it acts on most metals and on the halogens. It burns in air at about 300°, and with carbon at a red heat, but it does not combine with nitrogen.

Fine wires, or the powders of the greater number of metals, burn in the vapour of sulphur. The direct combination of hydrogen with sulphur is restricted by a limit--that is, at a given temperature and under other given conditions it does not proceed unrestrictedly; there is no explosion or recalescence. Sulphuretted hydrogen, H_{2}S, decomposes at its temperature of combination--that is, it is easily dissociated.[12] The same phenomenon is repeated here as with water, except that the temperatures at which the attraction of hydrogen for sulphur begins and ceases are much lower than in the case of oxygen and hydrogen. The temperature at which combination takes place is here, as in many other instances, nearly the same as that at which dissociation begins. Hence _sulphuretted hydrogen_ is formed in a small quantity by the direct ignition of a mixture of the vapour of sulphur and hydrogen. However, the temperature must not be high, because otherwise the whole of the sulphuretted hydrogen is decomposed; but at lower temperatures a small amount of sulphuretted hydrogen is formed by direct combination.[13] Sulphuretted hydrogen however, like all other hydrogen compounds, may be easily obtained by the double decomposition of its corresponding metallic compounds, the replacement of the metal by hydrogen being effected by the action of acids on the sulphides. The metallic sulphides are, as a rule, easily formed. A sulphide, when mixed with a non-volatile acid, may give, by double decomposition, a salt of the acid taken and sulphuretted hydrogen, M_{2}S + H_{2}SO_{4} = H_{2}S + M_{2}SO_{4}. However, it is not all sulphides nor solutions of all acids that will evolve sulphuretted hydrogen, which fact is exceedingly characteristic, because, for example, all carbonates evolve carbonic anhydride when treated with any acid. Sulphuric acid will only evolve sulphuretted hydrogen from those sulphides which contain a metal capable of decomposing the acid with the evolution of hydrogen. Thus zinc, iron, calcium, magnesium, manganese, potassium, sodium, &c., form sulphides which evolve sulphuretted hydrogen when treated with sulphuric acid, and the metals themselves evolve hydrogen with acids.[14] The sulphides of those metals which do not liberate hydrogen from acids do not generally act on acids--that is, do not form sulphuretted hydrogen with them; such are, for example, the sulphides of lead, silver, copper, mercury, tin, &c. Therefore, the _modus operandi_ of the formation of sulphuretted hydrogen by the action of acids on metallic sulphides may be looked on as a phenomenon of the combination of hydrogen, at the moment of its evolution, with the sulphur, which is combined with the metal. Such a representation is all the more simple as all the circumstances under which sulphuretted hydrogen is formed are exactly similar to the conditions of the formation of hydrogen itself. Thus the usual mode of preparing sulphuretted hydrogen is by the action of _sulphuric acid on ferrous sulphide_, in which the same apparatus and method are employed as in the preparation of hydrogen, only replacing the metallic iron or zinc by ferrous sulphide or zinc sulphide. The reaction between sulphide of iron and sulphuric acid takes place at the ordinary temperature, and is accompanied by just as small a development of heat as in the liberation of hydrogen itself, FeS + H_{2}SO_{4} = FeSO_{4} + H_{2}S.[15]

[12] In the formation of potassium sulphide, K_{2}S (that is, in the
combination of 32 parts of sulphur with 78 parts of potassium),
about 100 thousand heat units are developed. Nearly as much heat
is developed in the combination of an equivalent quantity of
sodium; about 90 thousand heat units in the formation of calcium
or strontium sulphide; about 40 thousand for zinc or cadmium
sulphide, and about 20 thousand for iron, cobalt, or nickel
sulphide. Less heat is evolved in the combination of sulphur with
copper, lead, and silver. According to Thomsen, sulphur develops
heat with hydrogen in solutions. The reaction I_{2},Aq,H_{2}S =
21,830 calories. But, as the reaction I_{2} + H_{2} + Aq develops
26,842 calories, it follows that the reaction H_{2} + S develops
4,512 calories.

[13] If sulphur be melted in a flask and heated nearly to its boiling
point, as Lidoff showed, the addition, drop by drop (from a funnel
with a stopcock) of heavy (0·9) naphtha oil (of lubricating
oleonaphtha), &c., is followed by a regular evolution of
sulphuretted hydrogen. This is analogous to the action of bromine
or iodine on paraffin and other oils, because hydrobromic or
hydriodic acid is then formed (Chapter XI.) A certain amount of
hydrogen sulphide is even formed when sulphur is boiled with
water.

[14] However, the matter is really much more complicated. Thus zinc
sulphide evolves sulphuretted hydrogen with sulphuric or
hydrochloric acids, but does not react with acetic acid and is
oxidised by nitric acid. Ferrous sulphide evolves sulphuretted
hydrogen with acids, whilst the bisulphide, FeS_{2}, does not
react with acids of ordinary strength. This absence of action
depends, among other things, on the form in which the native iron
pyrites occurs; it is a crystalline, compact, and very dense
substance; and acids in general react with great difficulty on
such metallic sulphides. This is seen very clearly in the case of
zinc sulphide; if this substance is obtained by double
decomposition, it separates as a white precipitate, which evolves
sulphuretted hydrogen with great ease when treated with acids.
Zinc sulphide is obtained in the same form when zinc is fused with
sulphur, but native zinc sulphide--which occurs in compact masses
of zinc blende, and has a metallic lustre--is not decomposed or
scarcely decomposed by sulphuric acid.

Another source of complication in the behaviour of the metallic
sulphides towards acids depends on the action of water, and is
shown in the fact that the action varies with different degrees of
dilution or proportion of water present. The best known example of
this is antimonious sulphide, Sb_{2}S_{3}, for strong hydrochloric
acid, containing not more water than corresponds with HCl,6H_{2}O,
even decomposes native antimony glance, with evolution of
sulphuretted hydrogen, whilst dilute acid has no action, and in
the presence of an excess of water the reaction 2SbCl_{3} +
3H_{2}S = Sb_{2}S_{3} + 6HCl occurs, whilst in the presence of a
small amount of water the reaction proceeds in exactly the
opposite direction. Here the participation of water in the
reaction and its affinity are evident.

The facts that lead sulphide is insoluble in acids, that zinc
sulphide is soluble in hydrochloric acid but insoluble in acetic
acid, that calcium sulphide is even decomposed by carbonic acid,
&c.--all these peculiarities of the sulphides are in correlation
with the amount of heat evolved in the reaction of the oxides with
hydrogen sulphide and with acids, as is seen from the observations
of Favre and Silberman, and from the comparisons made by Berthelot
in the Proceedings of the Paris Academy of Sciences, 1870, to
which we refer the reader for further details.

[15] _Ferrous sulphide_ is formed by heating a piece of iron to an
incipient white heat, and then removing it from the furnace and
bringing it into contact with a piece of sulphur. Combination then
proceeds, accompanied by the development of heat, and the ferrous
sulphide formed fuses. The sulphide of iron thus formed is a
black, easily-fusible substance, insoluble in water. When damp it
attracts oxygen from the air, and is converted into green vitriol,
FeSO_{4}. If all the iron does not combine with the sulphur in the
method described above, the action of sulphuric acid will evolve
hydrogen as well as hydrogen sulphide.

We will not describe the details of the preparation of
sulphuretted hydrogen employed as a reagent in the laboratory,
because, in the first place, the methods are essentially the same
as in the preparation of hydrogen, and, in the second place,
because the apparatus and methods employed are always described in
text-books of analytical chemistry. Ferrous sulphide may be
advantageously replaced by calcium sulphide or a mixture of
calcium and magnesium sulphides. A solution of magnesium
hydrosulphide, MgS,H_{2}S, is very convenient, as at 60° it
evolves a stream of pure hydrogen sulphide. A paste, consisting of
CuS with crystals of MgCl_{2} and water, may also be employed,
since it only evolves H_{2}S when heated (Habermann).

_In nature_ sulphuretted hydrogen is formed in many ways. The most usual mode of its formation is by the decomposition of albuminous substances containing sulphur, as mentioned above. Another method is by the reducing action of organic matter on sulphates, and by the action of water and carbonic acid on the sulphides formed by this reduction. Volcanic eruptions are a third source of sulphuretted hydrogen in nature. Although sulphuretted hydrogen is formed in small quantities everywhere, it nevertheless soon disappears from the atmosphere, owing to its being easily decomposed by oxidising agencies. Many mineral waters contain sulphuretted hydrogen, and smell of it; they are called 'sulphur waters.'

Sulphuretted hydrogen, at the ordinary temperature, is a colourless gas, having a very unpleasant odour. It has, as its composition H_{2}S shows, a specific gravity seventeen times greater than hydrogen, and therefore it is somewhat heavier than air. Sulphuretted hydrogen _liquefies_ at about -74°, and at the ordinary temperature when subjected to a pressure of 10 to 15 atmospheres; at -85° it is converted into a solid crystalline mass.[15 bis] The easy liquefaction of sulphuretted hydrogen is evidently allied to its solubility. One volume of water at 0° dissolves 4·37 volumes of sulphuretted hydrogen, at 10° 3·58 volumes, and at 20° 2·9 volumes.[16] The solutions impart a very feeble red coloration to litmus paper. This gas is poisonous. One part in fifteen hundred parts of air will kill birds. Mammalia die in an atmosphere containing 1/200 of this gas.

[15 bis] Liquid sulphuretted hydrogen is most easily obtained by the
decomposition of hydrogen polysulphide, which we shall presently
describe, by the action of heat, and in the presence of a small
amount of water. If poured into a bent tube, like that described
for the liquefaction of ammonia (Chapter VI.), the hydrogen
polysulphide is decomposed by heat, in the presence of water, into
sulphur and sulphuretted hydrogen, which condenses in the cold end
of the tube into a colourless liquid.

[16] Sulphuretted hydrogen is still more soluble in alcohol than in
water; one volume at the ordinary temperature dissolves as much as
eight volumes of the gas. The solutions in water and alcohol
undergo change, especially in open vessels, owing to the fact that
the water and alcohol dissolve oxygen from the atmosphere, which,
acting on the sulphuretted hydrogen, forms water and sulphur. The
solution may be so altered in this manner that every trace of
sulphuretted hydrogen disappears. Solutions of sulphuretted
hydrogen in glycerine change much more slowly, and may therefore
be kept for a long time as reagents. De Forcrand obtained a
hydrate, H_{2}S,16H_{2}O, resembling the hydrates given by many
gases.

Sulphuretted hydrogen is very easily _decomposed_ into its component parts by the action of heat or a series of electric sparks. Hence it is not surprising that sulphuretted hydrogen undergoes change under the action of many substances having a considerable affinity for hydrogen and oxygen. Very many metals[17] evolve hydrogen with sulphuretted hydrogen, so that in this respect it presents the property of an acid; for instance, 2H_{2}S + Sn = 2H_{2} + SnS_{2}. This may be taken advantage of for determining the composition of sulphuretted hydrogen, because a given volume then leaves the same volume of hydrogen. On the other hand, oxygen,[18] chlorine,[19] and even iodine decompose sulphuretted hydrogen, removing the hydrogen from it and leaving free sulphur, so that in this reaction the sulphur is replaced by the above-named elements; for example, H_{2}S + Br_{2} = 2HBr + S. In no other hydrogen compound is it so easy to show the _substitution_, both of hydrogen and of the element combined with it, as in hydrogen sulphide. This clearly proves the feeble union between the elements forming this gas. Compounds containing a considerable amount of oxygen, with which they easily part, can accomplish the separation of the sulphur very easily. Such are, for instance, nitrous acid, chromic acid, and even ferric oxide and the higher oxides like it. Thus, if sulphuretted hydrogen be passed into a solution of chromic acid or an acid solution of ferric oxide, water is formed, _and the sulphur is separated in a free state_. Thus, sulphuretted hydrogen acts as a _reducing agent_, in virtue of the hydrogen it contains. Salts of iodic, chlorous, chloric, and other acids are reduced by sulphuretted hydrogen, their oxygen acting mainly on its hydrogen; but in the presence of an excess of a powerful oxidising agent a portion of the sulphur may also be oxidised to sulphurous anhydride. The reducing action of sulphuretted hydrogen is frequently applied in chemical manipulations for the preparation of lower oxides, and for the conversion of certain oxygen compounds into hydrogen compounds: thus, the higher oxides of nitrogen are converted into ammonia by it, and in the presence of alkalis the nitro-compounds are converted into ammonia derivatives. The reaction of sulphuretted hydrogen on sulphurous anhydride belongs to this class of phenomena, the chief products of which are sulphur and water, 2H_{2}S + SO_{2} = 2H_{2}O + S_{3}.

[17] Some metals evolve hydrogen from sulphuretted hydrogen at the
ordinary temperature. For example, the light metals, and copper
and silver (especially with the access of air?) among the heavy
metals. Hence articles made of silver turn black in the presence
of vapours containing sulphuretted hydrogen, because silver
sulphide is black. Zinc and cadmium act at a red heat, but not
completely.

[18] If sulphuretted hydrogen escapes from a fine orifice into the air,
it will burn when lighted, and be transformed into sulphurous
anhydride and water. But if it burns in a limited supply of
air--for instance, when a cylinder is filled with it and
lighted--then only the hydrogen burns, which has, judging from the
amount of heat developed in its combustion and from all its
properties, a greater affinity for oxygen than sulphur. In this
respect the combustion of sulphuretted hydrogen resembles that of
hydrocarbons.

[19] Hence bleaching powder and chlorine destroy the disagreeable smell
of sulphuretted hydrogen. (For the reaction of hydrogen sulphide
and iodine, _see_ Chapter XI. p. 504.)

The acid character of sulphuretted hydrogen is clearly seen in its action on alkalis and salts.[19 bis] Thus lead oxide and its salts in the presence of sulphuretted hydrogen form water or an acid, and sulphide of lead: PbX_{2} + H_{2}S = PbS + 2HX. This reaction takes place even in the presence of powerful acids, because lead sulphide is one of those sulphides which are unacted on by acids, and in solutions the reaction is a complete one. This reaction is taken advantage of for the preparation of many acids, by first converting into a lead salt, and then submitting this salt to the action of sulphuretted hydrogen. For example, lead formate with sulphuretted hydrogen gives formic acid. Sulphuretted hydrogen in acting on a number of metallic acid substances in solution or in an anhydrous state also forms corresponding sulphates: (1) if it does not reduce the acid; (2) if the sulphur compound corresponding with the anhydride of the acid be insoluble in water, the reaction proceeds in solutions; (3) if the sulphuretted hydrogen and the acid taken do not come in contact with an alkali, on which they would be able to act first; and (4) if the sulphur compound be not decomposed by water. Thus solutions of arsenious acid give a precipitate of arsenious sulphide, As_{2}S_{3}, with sulphuretted hydrogen. This reaction proceeds not only in the presence of water, but also of acids, because the latter do not decompose the resultant sulphur compounds. The type of the decomposition is the same as with bases--that is, the sulphur and oxygen change places: RO_{_n_} + _n_H_{2}S = RS_{_n_} + _n_H_{2}O. Some sulphides corresponding with acid anhydrides are decomposed by water, and therefore are not formed in the presence of water. Such, for example, are the sulphides of phosphorus.[20]

[19 bis] Perfectly dry H_{2}S (Hughes 1892) has no action upon
perfectly dry salts, just as dry HCl does not react with dry
NH_{3} or metals (Chapter IX., Note 29).

[20] The sulphide P_{4}S is obtained by cautiously fusing the requisite
proportions of common phosphorus and sulphur under water; it is a
liquid which solidifies at 0°, and may be distilled without
undergoing change, but it fumes in air and easily takes fire. The
higher sulphide, P_{2}S, has similar properties. But little heat
is evolved in the formation of these compounds, and it may be
supposed that they are formed by the direct conjunction of whole
molecules of phosphorus and sulphur; but if the proportion of
sulphur be increased, the reaction is accompanied by so
considerable a rise of temperature that an explosion takes place,
and for the sake of safety red phosphorus must be used, mixed as
intimately as possible with powdered sulphur and heated in an
atmosphere of carbonic anhydride. The higher compounds are
decomposed by water. By increasing the proportion of sulphur, the
following compounds have been obtained: P_{4}S_{3} as prisms
(fuses at 165°, Rebs), soluble in carbon bisulphide, and unaltered
by air and water; _phosphorus trisulphide_, P_{2}S_{3}, is the
analogue of P_{2}O_{3}; it is a light yellow crystalline compound
only slightly soluble in carbon bisulphide, fusible and volatile,
decomposed into hydrogen sulphide and phosphorous acid by water,
and, like the highest compound of sulphur and phosphorus,
P_{2}S_{5}, it forms thio-salts with potassium sulphide, &c. This
_phosphorus pentasulphide_ corresponds with phosphoric anhydride;
like the trisulphide it gives hydrogen sulphide and phosphoric
acid with an excess of water. It reacts in many respects like
phosphoric chloride. The sulphide PS_{2} is also known; the vapour
density of this compound seems to indicate a molecule P_{3}S_{6}.

_Phosphorus sulphochloride_, PSCl_{3}, corresponds with phosphorus
oxychloride. It is a colourless, pleasant-smelling liquid, boiling
at 124°, and of sp. gr. 1·63; it fumes in air and is decomposed by
water: PSCl_{3} + 4H_{2}O = PH_{3}O_{4} + H_{2}S + 3HCl. It is
obtained when phosphoric chloride is treated with hydrogen
sulphide, hydrochloric acid being also formed; it is also produced
by the action of phosphoric chloride on certain sulphides--for
example, on antimonious sulphide, also by the (cautious) action of
phosphorus on sulphur chloride: 2P + 3S_{2}Cl_{2} = 2PSCl_{3} +
4S, by the action of PCl_{5} upon certain sulphides, for example,
Sb_{2}S_{3}, by the reaction: 3MCl + P_{2}S_{5} = PSCl_{3} +
M_{3}PS_{4} (Glatzel, 1893), and in the reaction 3PCl_{3} +
SOCl_{2} = PCl_{5} + POCl_{3} + PSCl_{3}, showing the reducing
action of phosphorus trichloride, which is especially clear in the
reaction SO_{3} + PCl_{3} = SO_{2} + POCl_{3}. Thorpe and Rodger
(1889), by heating 3PbF_{2} or BiF_{3} with phosphorus
pentasulphide (and also by heating AsF_{3} and PSCl_{3} to 150°),
obtained thiophosphoryl fluoride as a colourless, spontaneously
inflammable gas (see further on, Note 74 bis, and Chapter XIX.,
Note 25). The action of PSCl_{3} upon NaHO gives a salt of
monothiophosphoric acid (Würtz, Kubierschky), H_{3}PSO_{3}, which
gives soluble salts of the alkalis.

The metallic sulphides corresponding with the metallic oxides have either a feeble alkaline or a feeble acid character, according to the character of the corresponding oxide, and therefore by combining together they are able to form saline substances--that is, salts in which the oxygen is replaced by sulphur. Thus sulphuretted hydrogen having the properties of a feeble acid[21] has, at the same time, the properties of water, and forms the type of the sulphur derivatives, which may also be formed by means of sulphuretted hydrogen, just as the oxides may be formed by the aid of water. But as sulphuretted hydrogen has acid properties, it combines more easily with the basic metallic sulphides. Hence, for instance, there exists a compound of sulphuretted hydrogen with potassium sulphide, potassium hydrosulphide, 2KHS = K_{2}S + H_{2}S, just as there are potassium hydroxides; but there are scarcely any compounds of sulphuretted hydrogen with the sulphides corresponding with acids. Thus the sulphides of the metals may be regarded either as salts of sulphuretted hydrogen or as oxides of the metals in which the oxygen is replaced by sulphur. In general terms the sulphides exhibit the same degrees of difference with respect to their solubility in water as do the oxides. Thus the oxides of the alkali metals, and of some of the metals of the alkaline earths, are soluble in water, whilst those of nearly all the other metals are insoluble. The same may be said as to the sulphides; the sulphides of the metals of the alkalis and certain of the alkaline earths are soluble in water, whilst those of the other metals are insoluble. Those metals, like aluminium, whose oxides--for example, Al_{2}O_{3}--have intermediate properties and do not form compounds with feeble acids, at least in a wet way, also do not form sulphides by this method, although these may be obtained indirectly. And in general the sulphides of the metals are easily formed in a wet way, and with particular ease if they are insoluble in water. In this case their salts enter into double decomposition with sulphuretted hydrogen, or with soluble sulphides, and give an insoluble sulphide--for instance, a salt of lead gives lead sulphide with sulphuretted hydrogen. By the action of sulphuretted hydrogen on a salt of a metal, a free acid must be formed besides the metallic sulphide. Thus if a metal M be in a state of combination MX_{2}, then by the action of sulphuretted hydrogen there will be formed, besides MS,[22] an acid 2HX. It is evident that sulphuretted hydrogen will not precipitate an insoluble sulphide from the salts of those metals whose sulphides react with free acid, such as zinc, iron, manganese, &c. The reaction FeCl_{2} + H_{2}S = FeS + 2HCl, and the like, do not take place because the acid acts on the ferrous sulphide. Antimonious sulphide is not acted on by dilute hydrochloric acid, but it is decomposed by strong acid, and therefore in presence of an excess of hydrochloric acid antimonious chloride does not entirely react with hydrogen sulphide, whilst the reaction 2SbCl_{3} + 3H_{2}S = Sb_{2}S_{3} + 6HCl is a complete one in a dilute solution and with a small quantity of acid. Those metallic sulphides which are decomposed by acids may be obtained in a wet way by the double decomposition of the salts of the metals, not with hydrogen sulphide, but with soluble metallic sulphides, such as sulphide of ammonium or of potassium, because then no free acid is formed, but a salt of the metal (potassium or ammonium) which was taken as a soluble sulphide. So, for example, FeCl_{2} + K_{2}S = FeS + 2KCl.[23]

[21] Sulphuretted hydrogen does not saturate the alkaline properties of
alkali hydroxides, so that a solution of potassium hydroxide will
not under any circumstances give a neutral liquid with
sulphuretted hydrogen. In this case the sulphuretted hydrogen
forms in solution only an acid salt with the potassium: KHO +
H_{2}S = KHS + H_{2}O. It must be supposed that the normal salt is
not formed in the solution--that is, that the reaction 2KHO +
H_{2}S = K_{2}S + 2H_{2}O does not take place. This is seen from
the fact that a development of heat, depending on the formation of
potassium hydrosulphide, KHS, is remarked when as much hydrogen
sulphide is passed into a solution of potassium hydroxide as it
will absorb. But if a further quantity of potassium hydroxide be
added to the resultant solution, heat is not developed, whilst if
alkali be added to potassium acid sulphate or sodium acid
carbonate, heat is developed. It must not be concluded from this
that H_{2}S is a monobasic acid, for here there is a question of
the decomposing action of water upon K_{2}S; K_{2}S and H_{2}O in
reacting on each other should absorb heat if the reaction of KHS
upon KHO evolves heat. Furthermore, it must be taken into account
that potassium oxide, K_{2}O, and the anhydrous oxides like it,
also do not exist in solutions, for whenever they are formed they
immediately react with the water, forming caustic potash, KHO, &c.
In the same way, directly potassium sulphide, K_{2}S, is formed in
water it is decomposed into potassium hydroxide and hydrosulphide:
K_{2}S + H_{2}O = KHO + KHS. Potassium sulphide, K_{2}S, in a
solid state corresponds with K_{2}O, although neither can exist in
solution.

[22] During recent years (beginning with Schulze, 1882) it has been
found that many metallic sulphides which were considered totally
insoluble do, under certain circumstances, form very unstable
solutions in water, as already mentioned in Chapter I., Note 57.
Arsenic sulphide is very easily obtained in the form of a solution
(hydrosol). Solutions of copper and cadmium sulphides may also be
easily obtained by precipitating their salts CuX_{2}, or CdX_{2},
with ammonium sulphide, and washing the precipitate; but they are
re-precipitated by the addition of foreign salts.

[23] In reality the preceding reaction should be expressed thus:
FeCl_{2} + 2KHS = FeS + 2KCl + H_{2}S (Note 21), because in the
presence of water not K_{2}S but KHS reacts. But as the
sulphuretted hydrogen takes no part in the reaction, it is usual
to express the formation of such sulphides without taking the
hydrogen sulphide proceeding from the potassium or ammonium
hydrosulphides into account. It is not usual to employ potassium
sulphide but ammonium sulphide--or, to speak more accurately,
ammonium hydrosulphide--in order to avoid the formation of a
non-volatile salt of potassium and to have, together with the
formation of the sulphide, a salt of ammonium which can always be
driven off by evaporating the solution and igniting the
residue--for instance: FeCl_{2} + (NH_{4})_{2}S = FeS + 2NH_{4}Cl.
Thus the metallic sulphides may be divided into three chief
classes: (1) _those soluble in water_, (2) _those insoluble in
water but reacting with acids_, and (3) _those insoluble both in
water and acids_. The third class may be easily subdivided into
two groups; to the first group belong those sulphides which
correspond with bases or basic oxides, and are therefore unable to
play the part of an acid with the sulphides of the alkalis, and
are insoluble in NH_{4}HS, whilst the sulphides of the second
group are of an acid character, and give soluble thio-salts with
the sulphides of the alkaline metals, in which they play the part
of an acid. To this group belong those metals whose corresponding
oxides have acid properties. It must be observed, however, that
not all metallic acids have corresponding sulphides, partly owing
to the fact that certain acids are reducible by sulphuretted
hydrogen, especially when their lower degrees of oxidation are of
a basic character. Such are, for instance, the acids of chromium,
manganese, &c. Sulphuretted hydrogen converts them into lower
oxides, having the properties of bases. Those bases which do not
combine with feeble acids, such as carbonic acid and hydrogen
sulphide, give a precipitate of hydroxide with ammonium
sulphide--for example, aluminium salts react in this manner. This
difference of the metals in their behaviour towards sulphuretted
hydrogen gives a very valuable means of separating them from each
other, and _is taken advantage of in analytical chemistry_. If,
for instance, the metals of the first and third groups occur
together, it is only necessary to convert them into soluble salts,
and to act on the solution of the salts with sulphuretted
hydrogen; this will precipitate the metals of the third group in
the form of sulphides, whilst the metals of the first group will
not be in the least acted on. Such a method of separating the
metals is considered more fully in analytical chemistry, and we
will therefore limit ourselves here to pointing out to which
groups the most common metals belong, and the colour which is
proper to the sulphide precipitated.

_Metals which are precipitated by sulphuretted hydrogen_, as
sulphides from a solution of their salts, even in the presence of
free acid:

The precipitate is soluble in ammonium sulphide:

_Platinum_ (dark brown) | _Antimony_ (orange)
_Gold_ (dark brown) | _Arsenic_ (yellow)
_Tin_ (yellow and brown) |

The precipitate is insoluble in ammonium sulphide: