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Common Ion Effect

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The common ion effect is responsible for the reduction in solubility of an ionic precipitate when a soluble compound combining one of the ions of the precipitate is added to the solution in equilibrium with the precipitate. It states that if the concentration of any one of the ions is increased, then according to Le Chateliers principle, the ions in excess should combine with the oppositely charged ions.

The solubility product is useful for evaluating the influence of other species on the solubility of salts of low aqueous solubility. One of the most important influences on the solubility of poorly soluble salts occurs on addition of another compound that has an identical ion to one of those of the salt.

Definition

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"The decrease in ionization of weak electrolyte by the addition of a strong electrolyte having ion common with the weak electrolyte, is termed as common ion effect." It is the supression of the dissociation of a weak electrolyte by the addition of a strong electrolyte having common ion.

The common ion effect is defined as the reduction of the solubility of a sparingly soluble salt if the solution already contains one of the ions. For example, the addition of NH4Cl (strong electrolyte) to a solution of NH4OH (weak electrolyte) results in the decrease in dissociation of NH4OH.

Common Ion Effect

Application

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Applications of common ion effect has given below:
  • Precipitation of soap: Soap is a sodium or potassium salt of higher fatty acid. To precipitate soap, add saturated NaCl solution in the soap solution.
  • Formation of common salt: When HCl gas is passed through sea water, crystallization of NaCl takes place due to the common ion effect.
  • Second group sulphides and third group hydroxides: The common ion effect is also applicable to the precipitation of second group sulphides and third group hydroxides.

Common Ion Effect Solubility

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From Le Chatelier principle, it is expected that increase in the concentration of any of the ion, it should combine with ion of opposite charge and some salt is precipitated till once again Ksp = Qsp.

If concentrated solution of NaCl and HCl is passed through it, the NaCl is precipitated due to increased concentration of Cl- ions from HCl.

$NaCl$ $\rightleftharpoons$ $Na^+ + Cl^-$

Due to common ion effect the equilibrium will shift to left side and a high purity NaCl can be obtained by common ion effect. The common ion effect is also used for almost complete precipitation of a particular ion as its sparingly soluble salt with very low value of solubility product for gravimetric estimation.

Explanation

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The concept of solubility product plays the important role in predicting the precipitation of salts but for adjusting the conditions for precipitation another concept called common ion effect, plays a significant role. Common ion effect can also be explained bu applying Le Chatelier principle. The common ion effect is an application of Le Chatelier principle. The common ion effect is quite general. For example, solid NH4Cl added to a 1.0M NH3 solution produces additional ammonium ions.

NH4Cl(s) $\overset{H_{2}O}{\rightarrow}$ NH4+(aq) + Cl-(aq)

and this causes the position of the ammonia water equilibrium to shift to the left.

NH3(aq) + H2O(l) $\rightleftharpoons$ NH4+(aq) + OH-(aq)

This reduces the equilibrium concentration of OH- ions.

The common ion effect is also important in solution of polyprotic acids. The production of protons by the first dissociation step greatly inhibits the succeeding dissociation steps, which of course also produce protons, the common ion.

Examples

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The suppression of the dissociation of a weak electrolyte by the addition of a strong electrolyte containing a "common ion" is known as common ion effect. For example, if sodium acetate is added to acetic acid, ionization of acetic acid decreases.

CH3COOH $\rightleftharpoons$ H+ + CH3COO-
CH3COONa $\rightarrow$ Na+ + CH3COO-

Salting of soap (RCOONa) on adding NaCl or purification of NaCl by passing HCl involve the principle of common ion effect. The solubility of a precipitate is lower in a solution that contains an ion in common with the substance. In a precipitation reaction, the common ion effect predicts decreased solubility of a precipitate.

The degree of ionization of NH4OH (a weak base) is suppressed by the addition of NH4Cl (a strong electrolyte). The ionization of NH4OH and NH4Cl in solution is represented as follows.

NH4OH(aq) $\rightleftharpoons$ NH4+(aq) + OH-(Aq) ..... weakly ionised
NH4Cl $\rightarrow$ NH4+(aq) + Cl-(aq) ..... strongly ionised

Due to the addition of NH4Cl, which is strongly ionized in the solution, concentration of NH4+ ion increases in the solution. Therefore according to Le Chatelier's principle, equilibrium in equation shifts in the backward direction in the favor of unionised NH4OH. The addition of NH4Cl suppresses the degree of ionisation of NH4OH. Thus the concentration of OH-ions in the solution is considerably reduced and the weak base NH4OH becomes still weaker base.

Problems

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Problem based on common ion effect is given below.

Solved Example

Question: The Ksp of AgI in aqueous solution is 1$\times$10-16. If a 1$\times$10-5M solution of AgNO3 is saturated with AgI, what will be the final concentration of the iodide ion?
Solution:
The concentration of Ag+ in the original AgNO3 solution will be 1$\times$10-5M because AgNO3 will fully dissociate. Some amount of AgI will dissociate into the solution. If this amount is called x, the silver concentration will become 1$\times$10-5M+x. Because no iodide was present in solution until the AgI began dissociating, the concentration of iodide will be x. Thus the Ksp expression is

Ksp = [Ag+][I-]

1 $\times$ 10-16 = [1 $\times$ 10-5 M+x][x]

The value of x is negligible. Thus the equation can be simplified as 

1 $\times$ 10-16 = [1 $\times$ 10-5M][x]

1 $\times$ 10-11 = x

Thus the concentration of [I-] = 1 $\times$ 10-11M.