Page 110 - Basic physical chemistry for the atmospheric sciences

96 Basic physical chemistry

Compare this with the change in pH when 0.020 mole of hydrochloric
acid is added to l L of pure water.
Solution. Before the addition of the hydrochloric acid, HCl , the
concentrations of acetic acid, HC 2 H30 2 , and acetate, C 2 H30 2 ( aq), in
the solution are each 0. 1 0 M . Because HCl is a strong acid, it reacts
completely with the acetate to form acetic acid. Therefore, after the
HCl is added, the concentrations of acetic acid and acetate are
(0. 1 0 + 0.020) M and (0. 10 - 0.020) M, that is, 0. 1 2 M and 0.080 M ,
respectively. Therefore, from Eq. (5 . 2 8), the new p H o f the solution
will be

[base]
pH = pK a + log
[acid]
0.08
= - l og( l . 7 4 x 1 0 - 5 ) + log
O. l 2
= 4 . 7 6 - 0. 1 8 = 4 .58
Therefore, the change in the pH of the buffer solution is (4.76 - 4 . 5 8)
or a decrease of 0 . 1 8 pH units. [Note that the Na+ (aq) released from
-
the sodium acetate combines with the Cl ( aq) to form NaCl, but this
is neither acidic nor basic.] We can assume that when 0.020 mole of
HCl dissolves in l L of pure water, 0 . 0 2 mole of H + (aq) is formed,
and this is not neutralized. Therefore, the concentration of the H + ( aq)
is 0.020 M and the pH of the solution is

pH = - log[H + (aq)] = - l og(0.020) = 1 . 7

Since the pH of pure water is 7 , the lowering of the pH in this case is
5 . 3 pH units. Comparing this with the decrease of 0. 1 8 pH units for
the buffer solution, we see the remarkable ability of a buffer to stabi­
lize the pH.

5 . 1 0 Complex ions
Metal ions can act as Lewis acids (i. e . , electron acceptors) toward
water molecules that serve as Lewis bases (i.e . , electron donors) and
toward other Lewis bases. This can have a profound effect on the
solubility of a metal salt. For example, AgCl(s) dissolves in aqueous
ammonia because of the Lewis acid-base interaction between
Ag + (aq) and NH 3 (aq)