common ion effect
Summary
The common ion effect is the decrease in solubility of a sparingly soluble salt, or in ionization of a weak acid or base, when the solution already contains an ion the species produces. Le Châtelier's principle explains it: the added ion shifts equilibrium toward the solid or undissociated form.
Detail
Meaning: A sparingly soluble salt such as AgCl dissolves by AgCl(s) ⇌ Ag⁺ + Cl⁻. If Cl⁻ is already present from NaCl, the ion product Q exceeds Ksp, so the equilibrium shifts left and AgCl precipitates until Q = Ksp. The salt is less soluble than in pure water. The same logic suppresses ionization of a weak acid when its conjugate base is added.
Equation: Ksp = [Ag⁺][Cl⁻] = s(c + s) ≈ sc, so s ≈ Ksp/c. Here s is the molar solubility of AgCl (mol/L), c is the Cl⁻ concentration already present (mol/L), and Ksp is the solubility product (usually written without units). The approximation c + s ≈ c works because s is small.
On the MCAT: You may be asked to rank a salt's solubility in water, 0.1 M NaCl and 0.1 M CaCl₂ (more Cl⁻ gives lower solubility), or to calculate s using the approximation. Buffer passages use the same idea for weak acids.
Common confusion: Ksp does not change; only the solubility does, because Ksp depends on temperature alone. The ion must truly be shared: NaNO₃ has no common ion with AgCl. Adding acid to dissolve CaF₂ is a different effect, since H⁺ removes F⁻ and raises solubility.
Sources
- Chemistry: The Central Science
- General Chemistry: Principles and Modern Applications
Written for the MCAT and checked before it was published. For study reference only. Spotted an error? Let us know.