[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, H2S,16H2O, resembling the hydrates given by many gases.
[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.])
[19 bis] Perfectly dry H2S (Hughes 1892) has no action upon perfectly dry salts, just as dry HCl does not react with dry NH3 or metals (Chapter IX., Note [29]).
[20] The sulphide P4S 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, P2S, 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: P4S3 as prisms (fuses at 165°, Rebs), soluble in carbon bisulphide, and unaltered by air and water; phosphorus trisulphide, P2S3, is the analogue of P2O3; 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, P2S5, 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 PS2 is also known; the vapour density of this compound seems to indicate a molecule P3S6.
Phosphorus sulphochloride, PSCl3, 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: PSCl3 + 4H2O = PH3O4 + H2S + 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 + 3S2Cl2 = 2PSCl3 + 4S, by the action of PCl5 upon certain sulphides, for example, Sb2S3, by the reaction: 3MCl + P2S5 = PSCl3 + M3PS4 (Glatzel, 1893), and in the reaction 3PCl3 + SOCl2 = PCl5 + POCl3 + PSCl3, showing the reducing action of phosphorus trichloride, which is especially clear in the reaction SO3 + PCl3 = SO2 + POCl3. Thorpe and Rodger (1889), by heating 3PbF2 or BiF3 with phosphorus pentasulphide (and also by heating AsF3 and PSCl3 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 PSCl3 upon NaHO gives a salt of monothiophosphoric acid (Würtz, Kubierschky), H3PSO3, which gives soluble salts of the alkalis.
[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 + H2S = KHS + H2O. It must be supposed that the normal salt is not formed in the solution—that is, that the reaction 2KHO + H2S = K2S + 2H2O 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 H2S is a monobasic acid, for here there is a question of the decomposing action of water upon K2S; K2S and H2O 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, K2O, 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, K2S, is formed in water it is decomposed into potassium hydroxide and hydrosulphide: K2S + H2O = KHO + KHS. Potassium sulphide, K2S, in a solid state corresponds with K2O, 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]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 CuX2, or CdX2, 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: FeCl2 + 2KHS = FeS + 2KCl + H2S (Note [21]), because in the presence of water not K2S 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: FeCl2 + (NH4)2S = FeS + 2NH4Cl. 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 NH4HS, 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.