It has already been mentioned on p. 14 that the rate of chemical change depends on the amount of each of the reacting substances present in unit volume. This last is generally termed the " concentration " of these substances, for the more concentrated the solution the greater the mass present in unit volume. Now, if two kinds of ions, such as Na and -CI, are present in solution, necessarily in equal numbers, the solution will also contain a certain number of molecules of non-ionised NaCl, formed by their union, and the relative number of ions and molecules will depend on the concentration; the number of ions in proportion to the number of non-ionised molecules will be greater, the greater the dilution. For each dilution (and for each temperature) a state of balance will result; the position of this equilibrium will depend on the relative rate at which ionisation and union of ions to form molecules go on; if ionisation takes place twice as quickly as combinations of ions to form molecules, then two-thirds of the dissolved substance will exist as ions, the remaining third being non-ionised molecules. If the solution is made more concentrated by evaporation, the conditions are changed, and the rate of ionisation is reduced compared with the rate of union of ions with each other. Suppose that concentration be pushed so far that solid salt separates out; the limit of concentration will be reached, since it is now impossible to alter the number of ions and of molecules in unit volume of the solution. The ratio will now remain constant, and if c and c be the concentrations of the ions (and they are, of course, equal), and if C be that of the non-ionised molecules, then c.c -k.C, k being a factor expressing the relative proportions of the non-ionised molecules. If k is very small, then there are many molecules and few ions present; if, on the contrary, k is large, the ions are numerous and the molecules few. The expression k.C is termed the "solubility-product."
To take a specific case:-A solution of ammonia in water consists partly of the ions NH4 and OH, and partly of non-ionised molecules of NH4OH ; it is a weak base-that is, the number of non-ionised molecules is much greater than that of the ions. In a solution containing 1.7 grams of ammonia per litre (one-tenth normal solution), only 1.5 per cent, of the total number of molecules exist in the ionic state. Hence a solution of ammonia, unlike one of caustic soda or potash, gives no precipitate of hydroxide when added to a solution of salts of the relatively strong bases, such as calcium, strontium, or barium chlorides. With salts of the weaker base magnesia, however, ammonia produces a precipitate of magnesium hydroxide. It is possible still further to reduce the ionisation of ammonia solution; this can be done by the addition of an ammonium salt, such as the chloride, which, like most such salts, is highly ionised. The reason is, that while (concentration of NHJ x (concentration of -OH) = k x (concentration of NH4OH), if more ammonium ions be added, the number of hydroxyl ions will diminish by union with NH4 ions, forming non-ionised ammonium hydroxide, because the increase of the number of ammonium ions will increase the value of the product on the left-hand side of the equation, and in order that it may balance that of the right, the relative number of molecules of NH4OH must be increased ; and we may see that if ammonium chloride is added to a solution of magnesium chloride, ammonia solution will no longer produce a precipitate of magnesium hydroxide; the ammonia is too weak a base, that is, it contains too few hydroxyl ions, which are the reason of its basic nature.
Let us now return to the consideration of the insolubility of sulphides of the copper group in acids and the solubility of such sulphides as that of zinc. No substance, as has been before remarked, is wholly insoluble in water; zinc sulphide, however, belongs to the very sparingly soluble compounds. Hence the product c(Zn)xc'=(S) has a very small value, for it is equal to ^.C(ZnS), which must necessarily be very small, seeing that the compound is so sparingly soluble. Now, the ions of H2S are H, H, and =S ; but though the ionisation is very small, hydrogen sulphide being a very weak acid, they are yet sufficient to reach the value of the very small solubility-product i.C(ZnS). If, however, the concentration of the =S-ions is still further diminished by addition of some compound rich in H-ions, such as H|-Cl.Aq, then the product <:(Zn) x <r'=(S) will be less than /'.C(ZnS), and there will be no precipitate; or if hydrochloric acid be added to precipitated zinc sulphide, it will be dissolved. On the other hand, the addition of acetic acid, a weak acid, and poor in hydrogen ions, does not bring about solution of zinc sulphide; indeed, the precipitation of zinc from a solution of its acetate by hydrogen sulphide is almost complete.
The solubility-product of copper sulphide, and of the other sulphides which are not soluble in dilute acids, is still less; hence hydrogen sulphide precipitates them from acid solution, for the concentration of the =S-ions of the hydrogen sulphide may be very much diminished without the product f(Cu) x c'=S becoming less than i.C(CuS), for CuS is still less soluble in water than ZnS.