54 PHYSICAL CHEMISTRY 



i. Solutions made up to 200 cc and containing 50 cc 0.2 m KH-phtha- 

 late 4- 



0.2 n HC1 I 46.70 39.60 32.95 26.42 20.32 14.70 9.90 5.97 2.63 

 PH at 20 [ 2.2 2.4 2.6 2.8 3.0 3.2 3.4 3.6 3.8 



2. Solutions made up to 200 cc and containing 50 cc 0.2 m KH-tphtha- 

 late + 



"0.2 n NaOH | 0.40 3.70 7.50 12.15 17.70 23.85 29.95 3545 39-85 43-QO 45.45 47 -° 

 PH lat 20° I 4.0 4.2 4.4 4.6 4.8 5.0 5.2 5.4 5.6 5.8 6.0 6.2 



3. Solutions made up to 200 cc and containing 50 cc 0.2 m KH 2 P0 4 + 

 50 cc 0.2 n KCI 4- 



PH at 20 ] 5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2 7.4 7.6 7.8 8.0 

 0.2 n NaOH I 3.72 5.70 8.6012.6017.8023.6529.6335.0039.5042.8045.2046.80 



4. Solutions made up to 200 cc and containing 50 cc o>f 0.2 m H 3 B0 3 + 

 0.2 n NaOH | 2.61 3.97 5.90 8.50 12.00 16.30 21.30 26.70 32.00 36.85 40.80 43.9 

 PH at 20 I 7.8 8.0 8.2 8.4 8.6 8.8 9.0 9.2 9.4 9.6 9.8 10.0 



These solutions should be made from the purest reagents, that 

 have been recrystallized three times. The acid potassium phthalate, 

 acid potassium phosphate, and potassium chloride may be dried 

 at no° but the boric acid must be dried at room temperature 

 in a desiccator. 



The more acid of the mixtures will deposit crystals of phthalic 

 acid if not kept considerably above 20 . 



It may be inferred that the buffer value of these solutions is 

 less than those of Sorensen, but their salt action on indicators 

 is also less. 



Dissociation Constants of Acids and Bases 



Acids and bases differ from neutral salts in that among their 



dissociation products are the ions of water, H" or OH' as the 



case may be. Since the strength of acids and bases lies in the 



number of H and OH ions they dissociate, it is determined by 



the dissociation constant, c. 



a 2 



From the law of mass action we have: c = — , where 



v(i— a) 



v is the number of liters in which 1 mol of the acid (for in- 

 stance) is dissolved, i- — a represents the proportion of the un- 

 dissociated molecules and a represents the proportion of anions 

 = cations. If we increase the dilution, v, then a 2 increases and 

 therefore (a = H°) increases. (H* per liter decreases.) 



Strong acids appear to disobey this law when dissociation is 

 determined by electric conductivity. If c is determined from 

 the dissociation of a more concentrated solution, a more dilute 

 solution is found to be dissociated less than is calculated from 

 the formula. In other words, the dissociation in concentrated 



