Common Ion Effect
Definition and meaning of Common Ion Effect in chemistry.
The common ion effect is the decrease in the degree of ionization of a weak electrolyte, or the decrease in the solubility of a slightly soluble ionic compound, caused by adding a soluble source of an ion that the equilibrium already produces.
In more detail
Introducing extra common ion raises its concentration on the product side of the equilibrium, so by Le Chatelier's principle the system shifts back toward the reactants, suppressing further ionization or dissolution. For a sparingly soluble salt, this means the ion product briefly exceeds Ksp, driving more solid to precipitate until equilibrium is restored at a lower solubility. The effect is central to how acid-base buffers resist pH change and how selective precipitation is used to separate ions in qualitative analysis.
Key facts
| Field | General Chemistry |
|---|---|
| Governing Principle | Le Chatelier's Principle |
| Effect on Weak Acids/Bases | Decreases degree of ionization |
| Effect on Sparingly Soluble Salts | Decreases solubility (system re-equilibrates to Q = Ksp) |
Adding solid sodium acetate (CH3COONa) to a solution of acetic acid (CH3COOH) supplies extra acetate ions, shifting the equilibrium CH3COOH ⇌ H+ + CH3COO- to the left and giving a lower H+ concentration (higher pH) than acetic acid alone at the same concentration.
Frequently asked questions
Why does the common ion effect lower the solubility of a salt like AgCl?
Adding a soluble chloride source (e.g., NaCl) increases [Cl-] so the ion product [Ag+][Cl-] momentarily exceeds Ksp; extra AgCl precipitates until the product again equals Ksp, leaving less AgCl dissolved than in pure water.
How does the common ion effect create a buffer?
Mixing a weak acid with a salt of its conjugate base (a common ion source) suppresses the acid's ionization, giving a solution with a stable, predictable H+ concentration that resists changes on adding small amounts of acid or base.