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The common ion effect occurs when a common ion shared by two compounds is added to an equilibrium solution, disturbing the balance. When sodium acetate is added to acetic acid solution, the increased acetate ion concentration shifts equilibrium leftward, decreasing acid dissociation. This phenomenon results from Le Châtelier's principle: the system counterbalances the concentration change by shifting to reduce the added ion's effect.

Adding a common ion decreases the solubility of an ionic compound. For example, adding silver ions to a silver iodide solution shifts the dissolution equilibrium leftward, causing AgI precipitation and lowering dissolved ion concentrations. The solubility product remains constant, so increased concentration of one ion requires proportional decrease in the other, reducing overall solubility.

Adding ammonium chloride to ammonia solution introduces ammonium ions, a common ion that shifts ammonia's dissociation equilibrium leftward. This decreased dissociation produces fewer hydroxide ions, lowering the solution's pH. For example, adding 0.040 M ammonium chloride to 0.050 M ammonia reduces pH from 10.97 to 9.34 due to suppressed ammonia ionization.

Set up an ICE table listing initial concentrations of the weak acid or base and common ion, changes in concentration denoted by x, and equilibrium values. Substitute equilibrium expressions into the dissociation constant equation. Verify the approximation using the 5% rule: if x is less than 5% of the common ion concentration, the approximation is valid, allowing simplified calculations.

Le Châtelier's principle states that equilibrium systems shift to counterbalance concentration changes. When a common ion is added, the system perceives stress and shifts equilibrium to consume the added ion, restoring balance. This principle explains why adding acetate ions to acetic acid shifts equilibrium leftward, decreasing dissociation and reestablishing the equilibrium constant.

Adding magnesium chloride to a saturated Mg(OH)2 solution increases Mg2+ concentration, making the reaction quotient exceed Ksp. The system shifts left, precipitating Mg(OH)2 until quotient equals Ksp again. Similarly, adding potassium hydroxide increases OH– concentration, causing the same leftward shift and precipitation until equilibrium is restored.

The common ion effect enables selective precipitation and separation of ions in qualitative analysis. By controlling common ion concentrations, chemists can precipitate specific ions while keeping others in solution. This principle is essential for identifying cations and anions, purifying compounds, and understanding solubility behavior in complex solutions containing multiple ionic species.