president's address — SECTION A. 51 



took respecting the functions of the solvent, and ascribe to it a purely 

 passive part, viz., that of furnishing a space in which the ions could 

 play their various parts unhindered. Such a view necessarily involved 

 the assumption that the solution pressure of a metal is independent of 

 the chemical nature of the solvent — in other words, that it is a specific 

 property of the metal. For this reason, Nernst held that the solution 

 pressures may be taken as a numerical measure of the relative chemical 

 affinities of the metals ; not, of course, in quite the same sense as the 

 electromotive force; the latter measures the chemical energy, M'hile 

 the former measures the chemical force, so that the two definitions 

 of affinity may be regarded as collateral. 



The application of Nernst's complete formula to cells of ordinary 

 type would have led to very complicated expressions for the electro- 

 motive force. On the ground of sheer expediency, therefore, Nernst 

 suggested a simplified expression obtained by neglecting, as relatively 

 unimportant, all the potential differences other than those at the metal- 

 liquid surfaces. The electromotive force is then given by 



(s) E = — ] — log. — -^ — — los:. 



Pi «2 ° Pi 



The enunciation of this formula lent a new interest to the determina- 

 tion of single potential differences, as, according to it, the difference of 

 two such quantities is, very approximately, the electromotive force of 

 any ordinary cell. Measurements of this kind soon showed that the 

 formula gave consistent results. The solution pressures of any one 

 metal deduced from the potential at its point of contact with aqueous 

 solutions of its sulphate and of its chloride came out identical within 

 the Hmits of experimental error. Nernst's conclusion, that the part of 

 the electromotive force dependent on the negative ion is negligibly 

 small, was also confirmed in a large number of cases. The formula was, 

 however, found to break down wherever the negative ion was of such 

 a nature as greatly to reduce the solubility of either electrolyte. Great 

 was the joy of the opponents of the theory, while, so far as I know, none 

 of its supporters ever succeeded in eliminating the difficulty. Yet the 

 explanation is not far to seek — at least, I think not. To my mind these 

 are the very cases in which we might expect the electromotive forces, 

 neglected in deducing the approximate formula, to be large enough to 

 upset the basis of the approximation. The reasonableness of such an 

 explanation is best seen by considering an extreme case. If one of the 

 electrolytes is practically insoluble, while the solubility of the other is 

 finite, it is not very difficult to show (from Planck's complete formula (?) 

 for the electromotive force of electrolytes in contact) that the potential 

 difference at the boundary of the two electrolytes tends to become 

 indefinitely large. It is not the theory that Isreaks down in these 

 cases : it is only the approximation. So far, then, Nernst's theory 

 of solution pressures stands on a pretty safe basis, while the enormous 

 utihty it has had for investigators is beyond question. 



(s) The usual notation of the exponents of Nernst's theory is followed. 



(i) Wied. Ann., XL. I do not give the mathematics of my argument, as they 



are too straightforward to be worth printing. 



