190 EQUILIBRIUM BETWEEN ELECTROLYTES, chap.
acetate ions, we have, according to the law of mass action, before the addition —
(7o(Ag) X 6VCH3COO) = Kx qAgCHaCOO),
and after the addition —
(7,(Ag) X (7i(CHsC00) = X X qAgCHaCOO).
^(AgCHaCOO) is the same in both cases. On the other hand, ^(CHsCOO) is greater than Cf^GRnCOO) on account of the addition of CH3COO ions from sodium acetate or other compound. Consequently Ca(Ag) must be just as much smaller. The amount of dissolved silver is there- fore smaller in the second case than in the first. This agrees with the long-known fact that the solubility of many diflScultly soluble salts is decreased by the addition of neutral salts with a common ion. Apparent exceptions to this rule, e.g. increase of solubility of silver cyanide by the addition of potassium cyanide, are due to the formation of double salts (such as KAg(CN)2). In order to effectively precipitate difficultly soluble salts, e.g, barium sulphate, it is usually recommended in analytical descriptions to add excess of the precipitant, in this case barium chloride or sulphuric acid.
Van't Hoff (1) first suggested that the product of the ionic concentrations of a difficultly soluble electrolyte is constant.
As already mentioned (p. 164), salts deviate from the law of mass action so that their dissociation constants, JT, must be replaced in this relationship by a function of the quantity of the ions pi-esent, therefore the equations given cannot claim an absolute exactitude.
Another circumstance aids the deviation of the equations from exactness. The solubility of these difficultly soluble substances (in water) is frequently considerably influenced by the presence of even quite small quantities of foreign substances, such as alcohols, cane sugar, glycerol, etc. Euler (?) and Rothmund (-7) have shown from their own and previous experiments that ions possess in a marked
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