By Gould R.F. (ed.)
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Of pure water and of a buffer solution on addition of an acid or base in the same amounts is illustrated in figure 2. The magnitude of the pH-stabilizing effect of buffer solutions, the buffer capacity, is the concentration of a strong acid or base, in gequivalent/dm 3 required to cause a unit change in ρΆ. Buffer capacit y is unity if 1 g-equivalent of an acid or base changes the ρΉ. of 1 dm 3 of buffer solution by one unit. If the acid-base ratio is too small or too 35 CHEMICAL EQUILIBRIA IN SOLUTION 14 12 / 1 10 8 6 4 X \2 \ 2 v»_ 005 vj O-IO 0-15 0-20 mol/dm : i Figure 2 C h a n g e s i n t h e pH.
From conditions (i), (ii) and (iii), the hydrogen ion concentration of a strong acid (perchloric acid) is: [H+] = [C104-] + [ 0 H - ] = Ga + - ^ - . After rearrangement, this gives: and [ H + ] « - O e [ H + ] - Z w = 0, [H+] ca + (Ci + Œvy» thus PK 2 = - lg °° + «* + ^ " . (52) The concentration of the hydrogen ions produced by the autoprotolysis of water, and the hydroxide ion concentration equal to it, are negligible if they are less than 5 % of the analytical concentration of the strong acid.
Buffer solutions are widely used in laboratory practice to keep the ρΉ. of various solutions at a constant value. They also play an important role, for instance, in physiology, because living organisms can function properly only within certain narrow pK ranges. The constant pH of blood, for example, is maintained by a carbonate-hydrogen carbonate buffer system. The pH of polybasic acids and multiacidic bases Weak polybasic acids are dissociated in simple, consecutive steps. The dissociation constants for each subsequent step in general tend to decrease, t h a t is, Kal > Ka2 > Kan.
Applied Chemistry at Protein Interfaces by Gould R.F. (ed.)