630 



8GIENGE 



[N. S. Vol. XXXIV. No. i 



first sight, this apparent obedience of ab- 

 normal electrolytes to the mass-action law 

 seems strange, but a little consideration 

 shows that if it is only the non-ionized por- 

 tion of a salt that is truly abnormal, the 

 theoretical result is to be expected. Sup- 

 pose that the ions do behave normally in 

 the ionization, then they must also act with 

 normal active mass with reference to the 

 solid, with which they may be regarded as 

 in direct equilibrium according to the closed 

 scheme referred to above. A change, then, 

 in the concentration of any one of the ions, 

 brought about by the addition of a foreign 

 salt with that ion, will necessarily bring 

 about the change in solubility of the salt 

 calculated from the mass-action law, so far 

 at least as experiment can teU us, for any 

 variation from theory is caused by the 

 change in the nature of the solvent due to 

 the addition of the foreign substance. We 

 ought, then, on the assumption that the 

 ions behave normally, to expect that the 

 principle of the constant solubility product 

 would yield results of the same degree of 

 accuracy in dilute solutions whether the 

 electrolytes considered were normal or ab- 

 normal. This, as I have said, is actually 

 the case. 



To put the whole matter briefly, in the 

 equilibrium between electrolytes agreement 

 will be obtained between theory and experi- 

 ment whether we use the mass-action law, 

 or an empirical law such as van't Hoff's 

 dilution formula, provided only that we 

 attribute the abnormality to the non-ionized 

 portion of the electrolyte. Thus we can 

 deduce the ordinary formula for hydroly- 

 sis or for isohydric solutions as readily for 

 abnormal as for normal electrolytes, and 

 find the most satisfactory agreement with 

 experiment in both cases. 



By this one simple assumption, then, for 

 which I have offered some direct justifica- 

 tion, it is possible to find a basis for calcu- 



lation with abnormal electrolytes. The 

 problem of luhy certain electrolytes should 

 be normal and others abnormal is, of course, 

 in no way touched by this assumption. 

 That is a matter for further investigation 

 and research. 



Another great desideratum of the theory 

 of solutions is to find a general basis for 

 the calculation of hydrates. The present 

 position of the theory of hydrates in solu- 

 tion may perhaps most aptly be compared 

 to the theory of electrolytic dissociation for 

 solvents other than water. That hydrates 

 exist in some aqueous solutions is un- 

 doubted, but no general rule or method ex- 

 ists for determining what the hydrates are 

 and in what proportions they exist. Simi- 

 larly the theory of electrolytic dissociation 

 applied to other than aqueous solutions af- 

 fords no general means of determining 

 what the ions are and how great is the 

 degree of ionization. It is only for aqueous 

 solutions that Arrhenius was able to give 

 a practically realizable definition of degree 

 of ionization, and it is on this definition 

 that the whole effective work on aqueous 

 electrolj^tes is based ; and until some general 

 practically applicable principle of a similar 

 character is attained for hydrates, the work 

 done on that subject, however interesting 

 and important it may be in itself, must 

 necessarily be of an isolated character. 



Arrhenius did not originate the doctrine 

 of electrolytic dissociation or free ions: 

 that was enunciated in 1857 by Clausius, 

 and remained relatively barren. What he 

 did was to introduce measurable quantities 

 into the doctrine, and to show its simple 

 quantitative applicability to aqueous solu- 

 tions; immediately it became fertile. And 

 as soon as a simple quantitative principle 

 is developed for hydrates in solution, that 

 doctrine will become fertile also. 



It is surely now time that aU the irrele- 

 vant and intemperate things that have 



