7.11.8 Common Ion Effect in the Ionization of Acids and Bases

Consider an example of acetic acid dissociation equilibrium represented as:

CH3COOH(aq) ⇔  H+(aq) + CH3COO(aq)

Or HAc(aq) ⇔   H+(aq) + AC(aq)

Ka = [H+][AC]/[HAc]


Addition of acetate ions to an acetic acid solution results in decreasing the concentration of hydrogen ions, [H+].

Also, if H+ ions are added from an external source then the equilibrium moves in the direction of undissociated acetic acid i.e., in a direction of reducing the concentration of hydrogen ions, [H+]. This

phenomenon is an example of common ion effect. It can be defined as a shift in equilibrium on adding a substance that provides more of an ionic species already present in the dissociation equilibrium.

Thus, we can say that common ion effect is a phenomenon based on the Le Chatelier’s principle discussed in section 7.8.


In order to evaluate the pH of the solution resulting on addition of 0.05M acetate ion to 0.05M acetic acid solution, we shall consider the acetic acid dissociation equilibrium once again,

HAc(aq) ⇔   H+(aq) + AC(aq)

Initial concentration (M)

0.05                0               0.05


Let x be the extent of ionization of acetic acid.

Change in concentration (M)

-x                     +x              +x

Equilibrium concentration (M)

0.05-x               x               0.05+x



Ka = [H+][AC]/[HAc] = {(0.05+x)(x)}/(0.05-x)

As Ka is small for a very weak acid, x<<0.05.

Hence (0.05+x) is similar to (0.05-x) is similar to 0.05

Thus, 1.8 X 10-5 = (x)(0.05+x)/(0.05-x) = x(0.05)/0.05 = x = [H+] = 1.8 X 10-5 M

pH = -log(1.8 X 10-5) = 4.74

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