# The Common Ion Effect On Solubility Of Ionic Salts

The reduction in solubility of an ionic precipitate when a soluble compound that has one of the ions of the precipitate is added to the solution in equilibrium with the precipitate is called common ion effect. The common ion effect states that if the concentration of any ion is increased, then according to Le Chatelier’s principle the ions which are present in the excess should combine with the oppositely charged ions.Some amount of the salt will be precipitated out until the ionic product is equal to the solubility product. Similarly if the concentration of any one ion is decreased then more salt will be dissolved in order to increase the concentration of both the ions.

In simple words we can say that common ion effect is the suppression of disassociation of the weak electrolyte containing the common ion. Let us take an example to understand this, when HCl gas is passed through a saturated solution of sodium chloride, sodium chloride is precipitated out due to excess concentration of chloride ions. There is an excess of chloride ions due to disassociation of HCl. Sodium chloride obtained from the process is highly pure.

We use common ion effect for the complete precipitation of any particular ion as its sparingly soluble salt which has a low value of solubility product for gravimetric estimation. Therefore silver ion can be precipitated as silver chloride and ferric ions as hydroxide.

At lower value of the pH, the solubility of weak acid salts like phosphates increases. This happens because the concentration of anions at the lower value of pH decreases due to protonation. This in turn increases the solubility of the salt in order to maintain Ksp = Qsp . In this case we have to satisfy two conditions of equilibria simultaneously,

Ksp = [M+][X]

HX (aq) ⇔ H+(aq) + X (aq);

Ka = $\frac {[H^+ (aq)][X^- (aq)]}{[HX (aq)]}$

[X] / [HX] = $\frac {Ka}{[H^+]}$

Taking the inverse of the above equation and adding 1 on both the sides,

$\frac{[HX]}{X^-}$ + 1 =$\frac{[H^+]}{K_a}$  + 1

$\frac {[HX]+[X^-]}{X^-}$=$\frac {[H^+]+K_a}{K_a}$

Now again we take the inverse of the above equation,

$\frac {[X^-]}{[X^-]+[HX]}$ = f = $\frac {K_a}{(K_a + [H^+])}$

We can see that the value of f decreases as pH decreases. If S is the solubility of a salt at a given value of pH, then

Ksp = [S][f S] = S $\frac {K_a}{(K_a + [H^+])}$

S = $\left( \frac {K_{sp} ([H^+] + K_a)}{K_a}\right)^{0.5}$

Therefore solubility increases with increases in [H+] or decrease in the pH value.

We have so far read about the common ion effect, its properties, application and the dependence of solubility on pH. For any further details on this topic install Byju’s the learning app and enjoy the experience of learning in the most innovative ways.’

#### Practise This Question

The initial rate of hydrolysis of methyl acetate (1M) by a weak acid (HA, 1M) is 1/100th of  that of a strong acid (HX, 1M), at 25oC. The Ka of HA is