Page 77 - Vogel's TEXTBOOK OF QUANTITATIVE CHEMICAL ANALYSIS

COMPLEX IONS 2.21

Table 2.2 pH of acetic acid-sodium acetate buffer
mixtures
10 mL mixtures of x mL of 0.2M acetic acid and y mL of
0.2M sodium acetate
Acetic acid (x mL) Sodium acetate (y mL) pH

The activity coefficient y, of the undissociated acid is approximately unity in
dilute aqueous solution. Expression (24) thus becomes:
[ Acid]
aH+ = x Ko
[Salt] x y,-
or pH = pK, + log[Salt]/[Acid] +log y,- (26)
This is known as the Henderson-Hasselbalch equation.
If a buffer solution is diluted, the ionic concentrations are decreased and so,
as shown in Section 2.5, the ionic activity coefficients are increased. It follows
from equation (26) that the pH is increased.
Buffer mixtures are not confined to mixtures of monoprotic acids or monoacid
bases and their salts. We may employ a mixture of salts of a polyprotic acid,
e.g. NaH2P04 and Na2HP04. The Salt NaH2P04 is completely dissociated:
NaH,PO,e Na+ + H2P04
The ion H2P04 acts as a monoprotic acid:
H2P04 = H+ + HPO2-
for which K ( = K, for phosphoric acid) is 6.2 x IO-' mol L-'. The addition
of the Salt Na2HP04 is analogous to the addition of, Say, acetate ions to a
solution of acetic acid, since the tertiary ionisation of phosphoric acid
(HP0;-=H+ + PO:-) is small (K, = 5 x IO-', molL-'). The mixture of
NaH2P04 and Na2HP04 is therefore an effective buffer over the range
pH 7.2 + 1.0 (=pK + 1). It will be noted that this is a mixture of a
Br~nsted-Lowry acid and its conjugate base.
Buffer solutions find many applications in quantitative analysis, e.g. many
precipitations are quantitative only under carefully controlled conditions of pH,
as are also many compleximetric titrations: numerous examples of their use will
be found throughout the book.
2.21 COMPLEX IONS

The increase in solubility of a precipitate upon the addition of excess of the
precipitating agent is frequently due to the formation of a complex ion. A