Exp.—A solution of potassium hydrosulphide is saturated with hydrogen sulphide, in order to prevent hydrolysis and the formation of potassium hydroxide[496] (p. [180]), as far as possible, and the solution is added to an acid (hydrochloric) solution of methyl orange;[497] the acid color is changed to orange, as a result of the almost complete neutralization of the acid. The potassium hydrosulphide (the hydrosulphide-ion HS−) neutralizes the hydrogen-ion (of hydrochloric acid), that converts methyl orange into its pink salt, and hydrogen sulphide is formed, which is too weak an acid to affect the color of the indicator (p. [79]).
The objection that potassium sulphide and hydrosulphide are salts, the salts of hydrogen sulphide, might be raised against the conception of their possessing a certain measure of basic functions; but the common oxygen base, potassium hydroxide, is also a salt, the salt of a still weaker acid, water. Indeed, the characteristic properties of ordinary bases are due essentially to the fact, that they are the more or less readily ionizable salts of an extremely weak acid, water, and these properties may well be duplicated by salts of other weak acids, duplicated in a very much weaker way, in proportion as the acids are stronger than water. The difference is, then, really one of degree and not of kind.[498]
Owing to the fact that hydrogen sulphide is a much stronger acid than water, the action of potassium hydrosulphide on an acid sulphide, like carbon disulphide (equation (2), p. [243]), is reversed to a correspondingly greater degree than the action of potassium hydroxide on carbon dioxide[499] (equation (1), p. [243]). The dissociation constant for the secondary ionization of hydrogen sulphide (HS− ⇄ H+ + S2−) is very much smaller than the constant for the primary ionization (HS− is a much weaker acid than HSH), and so we find that a sulphide like K2S exhibits very much stronger basic functions than do the hydrosulphides, as, for instance, in forming salts with acid-forming sulphides [p246] (equation (3), p. [243]) and in neutralizing acids. There can be no question that, if we could have an aqueous solution of potassium oxide, K2O, it would show, similarly, the characteristic actions of strong bases even more powerfully than the hydroxide, KOH; for instance, in acting on acid-forming oxides (equation (1), p. [243]), in neutralizing acids, in saponifying esters (p. [81]), and so forth. It is, in fact, on account of this property, that potassium oxide is decomposed by water. It is a salt involving the secondary ionization of water, (HO− ⇄ H+ + O2−), which has a much smaller dissociation constant even than the primary ionization (H2O ⇄ H+ + HO−). The oxide, K2O, is decomposed by neutralizing hydrogen ions formed by the primary ionization of water. We have 2 K+ + O2− + H+ + HO− ⥂ 2 K+ + 2 HO−, which is entirely analogous, in principle, to K+ + HO− + H+ + Cl− ⥂ K+ + Cl− + HOH.
Sulphoxy-Salts.
Complex Sulphide Ions.
The ammonium-ion, appearing with the same coefficient on both sides of the last equation, evidently takes no direct part in the action and we have more simply: Sn4+ + 3 S2− ⇄ SnS32−.
For the condition of equilibrium between the complex and its components we have:[501],[502]
[Sn4+] × [S2−]3 / [SnS32−] = K.