Metallic Sulphides

Various methods are available for the preparation of the sulphides of the metals; only a few of the more general are here mentioned, details of other methods being obtainable under the headings of the respective metals in other volumes of this series.

1. Direct combination of the metal and sulphur can generally be effected, sometimes even at the ordinary temperature. This process is of especial advantage for the preparation of sulphides which are decomposed by water, e.g. aluminium sulphide. Finely divided mixtures of zinc and sulphur may be exploded by shock, heat or friction.
2. The metal may be heated in a current of hydrogen sulphide gas. As a rule this method will have the disadvantage relative to method (a) that the necessity to effect the displacement of hydrogen will render the reaction more difficult, especially with the less electropositive metals.
3. The metallic oxide, either alone or mixed with charcoal, may be heated in a stream of carbon disulphide vapour or merely with carbon disulphide under pressure, e.g. at 250° C. This process shares with method (a) the advantage of being conveniently applicable to the production of sulphides which are decomposed by water.
4. Salts of the metal with one of the sulphur oxyacids, e.g. a sulphate, sulphite or thiosulphate, may be reduced to the corresponding sulphide by heating with some suitable reducing agent such as charcoal, hydrogen or sulphur. This method is commonly employed as a stage in the conversion of the mineral barium sulphate into other barium salts, reduction with charcoal first yielding the more reactive sulphide.
5. In the case of many metals the sulphides can be formed by precipitation from aqueous salt solutions. The extensive use of hydrogen sulphide as a reagent in qualitative analysis depends, of course, on the different conditions necessary for the formation of the various sulphides. Some sulphides, e.g. antimony sulphide, can be quantitatively precipitated from solutions which are fairly strongly acidic, whereas in other cases the presence of acid leads to the setting up of an equilibrium; thus, lead sulphide is precipitated only partially in a solution containing 5 per cent, of hydrogen chloride, whilst ferrous sulphate solution yields ferrous sulphide only in the presence of sodium acetate or ammonium acetate, which removes the sulphuric acid as fast as it is liberated:
   
   PbCl₂ + H₂S ⇔ PbS + 2HCl.
   FeSO₄ + H₂S ⇔ FeS + H₂SO₄,
   H₂SO₄ + 2NaC₂H₃O₂ = 2C₂H₃O₂.H + Na₂SO₄.
   
   The presence of alkali and alkaline earth chlorides may also hinder precipitation of the sulphide; thus, from a 0.001 molar solution of lead chloride in water at 20° C., precipitation is completely inhibited by hydrogen chloride alone if in a concentration of 1.4N, and by decreasing concentrations of the acid in the presence of increasing quantities of calcium, ammonium or potassium chloride. Cadmium sulphide, precipitated from hydrochloric acid solution, contains adsorbed chlorine, the amount depending on the conditions of the precipitation; the precipitation is incomplete at 80° C.
   
   By using alcoholic benzene solutions of the alkali alcoholates it has been found possible to precipitate the corresponding alkali hydrogen sulphides by hydrogen sulphide.
6. Electrolysis of a solution of an alkali sulphide with an anode consisting of a metal such as copper, cadmium or silver will lead to the conversion of the anodic metal into sulphide.

In aqueous solution the normal sulphides of the alkali metals are very largely hydrolysed into the corresponding hydro- sulphides, so that the solutions react strongly alkaline on account of the liberated alkali hydroxide:

Na₂S + H₂O ⇔ NaSH + NaOH.

Electrolysis of such solutions with platinum electrodes and with low current densities yields polysulphides; with higher current densities sulphates and dithionates are formed.

The solutions undergo oxidation on exposure to air, sulphur first being liberated and then polysulphides formed, which in turn are oxidised to thiosulphates. The rate of this oxidation is greatly accelerated by the presence of small quantities of certain of the heavy metals, particularly manganese and nickel.

Solutions of the alkali sulphides give a deep violet to purple coloration with a solution of sodium nitroprusside, and this may be used as a test for sulphides. By the interaction of the nitroprusside with the sulphides of lithium, sodium, potassium and rubidium, stable crystalline compounds of the type M₄[Fe(CN)₅NOS] have been obtained. From electrotitrimetric evidence the reaction appears to proceed in two stages:

[Fe^III(CN)₅N:O]'' + HS' = ,
+ OH' = + H₂O

All the other normal sulphides are insoluble in water or are decomposed by it. The sulphides of the alkaline earth metals are sparingly soluble but, like the alkali sulphides, undergo hydrolysis to the corresponding hydrosulphides:

2CaS + 2H₂O = Ca(SH)₂ + Ca(OH)₂.

These hydrosulphides of the alkaline earth metals are soluble in water and therefore, by treating an emulsion of the normal sulphides with a current of hydrogen sulphide, it is possible to obtain solutions of the hydrosulphides Ca(SH)₂, Sr(SH)₂, Ba(SH)₂ and Mg(SH)₂.

A remarkable property of the sulphides of the alkaline earth metals and' of beryllium and zinc is their power, when certain impurities are present, to exhibit phosphorescence after exposure to bright light. The phenomenon is not due to slow oxidation and is still observable in samples which have been kept hermetically sealed for years; it is obvious, therefore, that the effect is a physical one and not analogous to the phosphorescence observable with sulphur. The nature and amount of impurity present considerably affect the phosphorescence, chlorides for example causing an increase; some impurities inhibit the action.

The sulphides of aluminium and silicon are decomposed immediately by water at the ordinary temperature, whilst even the sulphides of the heavier metals such as copper, lead and iron are decomposed by steam at a red heat:

Al₂S₃ + 6H₂O = 2Al(OH)₃ + 3H₂S.

Some sulphides, such as lead sulphide or iron pyrites, which are not decomposed by hydrochloric acid alone, yield their sulphur as hydrogen sulphide when metallic zinc is also present, the nascent hydrogen effecting the desired result. Hydrogen sulphide is also formed when iron pyrites is heated with coal, or in a stream of dry or moist hydrogen or moist carbon dioxide.

Cobaltous, nickelous and zinc sulphides when spread on a lead plate and subjected to cathodic polarisation in aqueous sulphuric acid, suffer reduction, some of the metal passing into solution and hydrogen sulphide being liberated.

The sulphides of the heavier metals are characterised by their pronounced colours, insolubility and ability under certain conditions to carry down, when precipitated, normally soluble sulphides of other metals. The freshly precipitated sulphides differ somewhat in solubility from those which have "aged" for some time; they are also capable of forming double salts, such as 2HgS.HgCl₂, FeS.NiS, Tl₂S.2CuS. Cadmium and manganese sulphides, precipitated together by means of ammonium sulphide, yield the compound 2MnS.3CdS. Mercuric sulphide, precipitated from acid solutions containing zinc or cadmium, always carries down considerable amounts of the .latter metals, and the ordinary use of hydrogen sulphide in qualitative analysis fails to effect a complete separation of mercury from zinc and cadmium, or, for similar reasons, of tin from cobalt. Examples of solubility changes are as follows: nickel sulphide is insoluble in 10 per cent, hydrogen chloride solution, but when co-precipitated with lead sulphide it is appreciably soluble; manganese sulphide is readily soluble in acetic acid, but digestion with acetic acid of a co-precipitated mixture of manganese and zinc sulphides leaves a residue of zinc sulphide which may contain up to 24 per cent, of manganese. The addition of a little hydrogen peroxide aids solution of mercuric sulphide in dilute hydrochloric acid, and of nickel and cobalt sulphides in dilute acetic acid.

The stability of the sulphides other than those of the alkali metals, the alkaline earth metals and aluminium, ranges from that of manganese sulphide, which is easily decomposed by dilute acids and slowly by ordinary steam, to that of mercuric sulphide or molybdenum sulphide, which resist the action of concentrated hydrochloric acid solution.

It will be noticed that, as a general rule, the more basic the character of a metal the more stable is its sulphide towards oxidation; an analogy therefore appears to exist between the function of sulphur in a sulphide and oxygen in an oxide. This analogy extends, in a less marked manner, to the behaviour of the corresponding sulphides and oxides towards alkalis. Just as carbon dioxide, arsenious oxide and the antimony oxides lend themselves to salt formation with alkalis, so carbon disulphide, arsenious sulphide and the antimony sulphides can combine with the alkali sulphides, giving rise to sulphur compounds (thio-salts) of analogous composition. Indeed, on treatment with an alkali, such a sulphide generally produces a mixture of the corresponding oxy-salt and the thio-salt. These sulphides can therefore be regarded as "thio-anhydrides." From the similarity in behaviour, it is probable that the sulphides are structurally analogous to the corresponding oxides.

The sulphides of the heavy metals may readily be desulphurised by ignition with aluminium; even zinc sulphide, to which is attributed a higher heat of formation than to the equivalent amount of aluminium' sulphide, may be thus reduced.

The crystal structures of certain metallic sulphides have been investigated.