November 10, 1911] 



SCIENCE 



629 



proportional to c,„ but rather proportional 

 to a power of d, other than the first, 

 namely, Cu"^*. 



This is, to my mind, a very strong piece 

 of evidence that in the ease of the abnormal 

 electrolyte, ammoniiim eyanate, the abnor- 

 mality of the ionization equilibrium is to be 

 attributed entirely to the non-ionized por- 

 tion. But ammonium eyanate differs in no 

 respect, with regard to its electrolytic con- 

 ductivity, from the hundreds of other ab- 

 normal binary electrolytes with univalent 

 ions; and I am therefore disposed to con- 

 clude that it is to the non-ionized portion 

 in general of these electrolytes that the ab- 

 normality is to be attributed. 



As I have already indicated, this con- 

 clusion is not altogether novel, but in my 

 opinion it has not been sufficiently empha- 

 sized. Even in discussions where it is 

 formally admitted that the divergence from 

 the dilution law may be due to the non- 

 ionized portion, yet the argument is almost 

 invariably condiicted 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 



di -'''•' 



or the constant k' of the reverse equation 



is really constant. If the former, then the 

 Ions are truly normal, and primary expla- 

 nations 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 Storeh, of which van't Hoff's dilu- 

 tion 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 evi- 

 dence that the ions act normally in solution. 

 In deducing this principle it is generally 

 assumed that it is the constant solubility 

 of the non-ionized salt that determines the 

 final equilibrium. This assumption, though 

 convenient, is not necessary. The equi- 

 librium is a closed one, thus: 



Solid Salt fc? Non-ionized Salt 



The solid is not only in equilibrium with 

 the non-ionized salt but also with the ions. 

 Now, in the deduction of the change of 

 solubility caused by the addition of a sub- 

 stance having one ion in common with the 

 original electrolyte the mass-action law for 

 ionization is assumed. This is of course 

 justified when we deal with feeble electro- 

 lytes, 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 M'hich 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 electro- 

 lyte of the addition of another electrolyte 

 with a common ion, whether both electro- 

 lytes are normal, both abnormal, or whether 

 one is normal and the other abnormal. At 



