DISSOCIATION OF STRONG ELECTROLYTES 127 



because of this that the preceding theoretical observations were 

 made. The result is the following. 



35. The significance of the activity theory for the pH of buffers 



In a mixture of a free weak acid and of its sodium salt the relation 

 (page 44) holds: 



= k ,f5^ (1) 



[Na — salt] 



where k is the dissociation constant of the acid, and assuming that: 

 (1) the salt is completely dissociated, and (2) as it must be again 

 emphasized, that the total concentration of electrolytes (i.e., ions) 

 in the solution is sufficiently small, so that concentration and activity 

 may be taken as being equal to each other. 



As was demonstrated in the last chapters, the first assumption is 

 quite valid for all practical purposes. The basis on which the in- 

 complete dissociation of Na-salts had been previous^ assumed has 

 been shown to be untenable, and at least all Na-salts are much more 

 strongly dissociated than it had been assumed; and even at that, in 

 0.1 N solutions for example, a 90 per cent dissociation had been 

 assumed for these salts. And now we must provisionally extend it 

 to 100 per cent. 



But the second assumption is valid only for very dilute solutions 

 and besides for buffer solutions free of neutral salts; and these are 

 physiologically unimportant. In fluids which are of physiological 

 interest we find much more frequently buffer solutions of the follow- 

 ing type: 



A small amount of free weak acid + a small amount of its Na-salt 

 + a relatively large amount of neutral salt (NaCl) ; for example, in 

 the blood: 



CO2 of the order of magnitude 0.002 molar 



NaHCOa of the order of magnitude 0.02 molar 



NaCl of the order of magnitude 0. 12 molar 



"a'^ 



or in sea-water : 



CO2 of the order of magnitude 0.0002 molar 



Bicarbonates of the order of magnitude 0.002 molar 



NaCl of the order of magnitude 0.6 molar 



