ACIDITY 297 



the proportion between the excess number of osmotically active 

 particles and the total number of molecules known to be present 

 in the solution. This factor i took care of the divergence of 

 electrolytic solutions from the perfect gas laws, but it told 

 nothing about the cause of the divergence. 



The young Swedish chemist Arrhenius, who later was a student 

 in van't Hoff's laboratory, interpreted the constant i of van't 

 Hoff in another way. He assumed that salt molecules break 

 down, or dissociate, when in solution, into their respective ions. 

 (It is now believed that there is no breaking down of the salt 

 molecules but merely a separation of the atoms or ions, because 

 there are no molecules as such in the salt crystal.) But there is a 

 slight discrepancy. The osmotic pressure of salt solutions is not 

 quite proportional to the number of ions that should be present 

 if dissociation is complete. The osmotic pressure of sodium 

 chloride is not quite twice, and of calcium chloride not thrice, 

 as great as is an equimolecular concentration of sugar. This fact 

 led Arrhenius to make a further postulate, a corollary to his 

 theory of dissociation. He proposed that some of the molecules 

 of a salt do not break down into ions and that therefore the num- 

 ber of particles (molecules and ions) present in the solution is 

 not quite twice (in a monovalent salt) that of the total number 

 of original molecules. This is the theory of incomplete dissocia- 

 tion. As a result, we have so-called dissociation constants of 

 electrolytes. 



The dissociation constant K expresses the ratio between the 

 product of the concentrations of the dissociated ions and the 

 concentration of the undissociated molecules. It is a special 

 case of the law of mass action and may be expressed as follows 

 for a hypothetical electrolyte AB: 



[A+] X [B-] 



K 



AB 



[AB] 



Dissociation constants are not constant for strong electrolytes, 

 nor do they indicate what it was originally thought that they 

 indicated in strong electrolytes (as we shall see). 



The dissociation constant K of picric acid is 1.4 X 10-^ (0.14), 

 which is a moderately high value. The constant of the mild 

 acetic acid is 1.86 X 10-^ (0.0000186), a low value; of the very 

 weak boric acid, K is 6.4 X 10""^°. Expressed in percentage, 



