PRESIDENTIAL ADDRESS. 355 



i. if. proportional to c m , but is rather proportional to a power of c„ other than the 

 first, namely c. 1 *. 



This is, to my mind, a very strong piece of evidence that in the case of the 

 abnormal electrolyte, ammonium cyanate, the abnormality of the ionisation 

 equilibrium is to be attributed entirely to the non-ionised portion. But am- 

 monium cyanate differs in no respect, with regard to its electrolytic conductivity, 

 from the hundreds of other abnormal binary electrolytes with univalent ions ; 

 and I am therefore disposed to conclude that it is to the non-ionised portion in 

 general of these electrolytes that the abnormality is to be attributed. 



As I have already indicated, this conclusion is not altogether novel, but in my 

 opinion it has not been sufficiently emphasised. Even in discussions where it is 

 formally admitted that the divergence from the dilution law may be due to the 

 non-ionised portion, yet the argument is almost invariably conducted so as to 

 throw the whole responsibility on the ions. The point which ought to be made 

 clear is whether the constant k of the equation 



or the constant h' of the reverse equation 



dx ,, 

 -dt =kc " 



is really constant. If the former, then the ions are truly normal, and primary 

 explanations of" the abnormality of the strong electrolytes can scarcely be sought 

 in high total ionic concentrations and the like, though a connection between 

 the two no doubt exists, both being determined by the same cause. 



In my illustration I have assumed that there holds good a dilution law of 

 the kind given by Storch, of which van't Hoff's dilution law is a particular case. 

 Here the active mass is represented as a power of the concentration other than 

 the first power. The argument I have used is altogether independent of this 

 special assumption; the active mass of the abnormal substance may be any 

 function of its concentration, and the same conclusion will be reached. 



Nernst's principle of the constant ionic solubility product affords additional 

 evidence that the ions act normally in solution. In deducing this principle it is 

 generally assumed that it is the constant solubility of the non-ionised salt that 

 determines the final equilibrium. This assumption, though convenient, is not 

 necessary. The equilibrium is a closed one, thus : — 



Ions „ 



\ 



Non-ionised Salt 



The solid is not only in equilibrium with the non-ionised salt but also with the 

 ions. Now, in the deduction of the change of solubility caused by the addition 

 of a substance having one ion in common with the original electrolyte the mass- 

 action law for ionisation is assumed. This is of course justified when we deal 

 with feeble electrolytes, but in the case of salts and strong acids which do not 

 follow the mass-action law the experiments are found- still to be in harmony with 

 the theoretical deductions. This is not only so when the two substances in 

 solution are both abnormal, but also when one is abnormal and the other normal, 

 no matter which is used to produce the saturated solution. In fact, the principle 

 of the constant ionic solubility product may be employed with equal success to 

 calculate the effect on the solubility of one electrolyte of the addition of another 

 electrolyte with a common ion, whether both electrolytes are normal, both 

 abnormal, or whether one is normal and the other abnormal. At first sight, this 

 apparent obedience of abnormal electrolytes to the mass-action law seems strange, 

 but a little consideration shows that if it is only the non-ionised portion of a 



A A 2 



