What is the primary principle behind the Common Ion Effect?

The primary principle behind the Common Ion Effect is Le Chatelier's Principle. This principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change. In the context of solubility, adding a common ion increases the concentration of one of the product ions, which is a disturbance. To relieve this stress, the equilibrium shifts towards the reactants (the undissolved solid), thereby reducing the solubility of the sparingly soluble salt.

Does the solubility product constant ($K_{sp}$) change due to the Common Ion Effect?

No, the solubility product constant ($K_{sp}$) does not change due to the Common Ion Effect. $K_{sp}$ is an equilibrium constant, and its value is dependent only on temperature. The Common Ion Effect alters the *equilibrium concentrations* of the ions in solution, causing the system to adjust its position of equilibrium, but the inherent ratio of ion concentrations at saturation (represented by $K_{sp}$) remains constant at a given temperature.

How does the Common Ion Effect differ from the 'salt effect' or 'ionic strength effect'?

The Common Ion Effect describes a *decrease* in solubility when an ion common to the sparingly soluble salt is added. In contrast, the 'salt effect' (or 'ionic strength effect') describes an *increase* in solubility of a sparingly soluble salt when a *non-common* inert salt is added. This increase is due to the added inert ions increasing the ionic strength of the solution, which reduces the activity coefficients of the sparingly soluble salt's ions, effectively allowing more of them to dissolve.

Can the Common Ion Effect be used for purification?

Yes, the Common Ion Effect is a very practical method for purifying salts. For example, to purify crude sodium chloride (NaCl), concentrated hydrochloric acid (HCl) gas is bubbled through a saturated solution of impure NaCl. The added $Cl^-$ ions from HCl act as a common ion, causing pure NaCl to precipitate out of the solution, leaving most of the impurities, which typically do not contain $Cl^-$ ions, dissolved in the mother liquor. This technique is widely used in industrial processes.

Why is the approximation often made that the contribution of the sparingly soluble salt to the common ion concentration is negligible?

The approximation is made because sparingly soluble salts, by definition, dissolve to a very small extent. When a soluble salt providing a common ion is added, its concentration is typically much higher (e.g., $0.1\,M$ or more) compared to the concentration of the common ion contributed by the sparingly soluble salt (often $10^{-5}\,M$ or less). Therefore, adding the small amount from the sparingly soluble salt to the already large concentration from the soluble salt has a negligible impact on the total common ion concentration, simplifying calculations without significant loss of accuracy.

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The Common Ion Effect describes the decrease in the solubility of a sparingly soluble ionic compound when a soluble salt containing a common ion is added to the solution. This phenomenon is a direct consequence of Le Chatelier's Principle, which states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change. In the context …

Quick Summary

The Common Ion Effect is a fundamental principle in chemistry that explains why the solubility of a sparingly soluble ionic compound decreases when a soluble salt containing a common ion is added to its solution.

This phenomenon is a direct application of Le Chatelier's Principle. When a common ion is introduced, it shifts the dissolution equilibrium of the sparingly soluble salt towards the undissolved solid, causing more of it to precipitate out of solution.

For example, adding sodium chloride (NaCl) to a saturated solution of silver chloride (AgCl) will decrease the solubility of AgCl because both salts provide chloride ions ( $Cl^-$ ). The solubility product constant ( $K_{sp}$ ) remains unchanged, but the equilibrium concentrations of the ions adjust to maintain this constant value.

This effect is crucial for understanding precipitation reactions, selective separation in qualitative analysis, and purification processes.

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Key Concepts

Impact on Solubility Product (

K_{sp}

)

It's a common misconception that the Common Ion Effect changes the $K_{sp}$ value. This is incorrect. The…

Quantitative Calculation of Solubility with Common Ion

Calculating solubility in the presence of a common ion involves setting up the $K_{sp}$ expression,…

Effect of pH on Solubility (Related to Common Ion Effect)

For sparingly soluble salts where one of the ions is a conjugate acid or base of a weak acid or base, the…

Definition: — Decrease in solubility of a sparingly soluble salt by adding a soluble salt with a common ion.

Principle: — Le Chatelier's Principle.

Equilibrium: —

M_aX_b(s) \rightleftharpoons aM^{b+}(aq) + bX^{a-}(aq)

$K_{sp}$: — Remains constant at a given temperature.

Effect: — Shifts equilibrium left, reducing solubility.

Calculation: — For

MX(s)

K_{sp} = [M^+][X^-]

. If common ion

[X^-]_C

is added, then

[X^-]_{\text{total}} \approx [X^-]_C

. Solubility

s = [M^+] = K_{sp}/[X^-]_C

Approximation: —

s \ll [Common\,Ion]_{\text{initial}}

is often valid.

pH Effect: — For salts with basic anions (e.g.,

OH^-

CO_3^{2-}

), adding acid increases solubility (removes anion). Adding base (common

OH^-

) decreases solubility of hydroxides.

Common Ion Effect: Causes Ions to Exit (precipitate)!

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