Common-Ion Effect

What the common-ion effect is

Drop a pinch of table salt into a beaker that is already a saturated solution of silver chloride and something quietly dramatic happens: the liquid goes faintly cloudy as fresh white solid drifts out of solution. Nothing was added that contains silver, yet more silver chloride precipitates. That counterintuitive result is the common-ion effect — the suppression of a substance's solubility (or of a weak electrolyte's ionization) when a second, fully soluble compound supplies an ion that the first substance already contributes to the solution.

The key word is common: the added salt and the dissolving salt share an ion. Sodium chloride and silver chloride both put chloride into the water. Once you flood the solution with chloride from the freely soluble NaCl, the sparingly soluble AgCl can no longer hold as much silver and chloride in solution simultaneously, so the excess crashes out. The effect is not a new law of nature — it is Le Chatelier's principle applied to an ionic equilibrium, named for the special case where the stress you apply is an ion the system already owns.

The mechanism: Q, Ksp, and the equilibrium shift

A slightly soluble ionic solid sits in dynamic equilibrium with its dissolved ions:

AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq)

The position of that equilibrium is fixed by the solubility product constant, Ksp = [Ag⁺][Cl⁻] = 1.8×10⁻¹⁰ at 25 °C. In pure water the solid dissolves until both ion concentrations equal the molar solubility s, so s² = Ksp and s = √(1.8×10⁻¹⁰) ≈ 1.34×10⁻⁵ M. Roughly 13 micromoles of silver per liter — already not much.

Now add 0.10 M NaCl. Sodium chloride is fully soluble, so it dumps 0.10 M chloride into the beaker instantly. The reaction quotient Q = [Ag⁺][Cl⁻] momentarily becomes (1.34×10⁻⁵)(0.10) ≈ 1.3×10⁻⁶, which is about 7000 times larger than Ksp. The solution is now supersaturated: Q > Ksp. Le Chatelier's principle says a system stressed by extra product will shift to consume it, so silver and chloride combine and deposit as solid until Q falls back to 1.8×10⁻¹⁰. Solving for the new solubility:

s(0.10 + s) = 1.8×10⁻¹⁰ → s ≈ 1.8×10⁻¹⁰ / 0.10 = 1.8×10⁻⁹ M

(The approximation 0.10 + s ≈ 0.10 is excellent because s is nanomolar.) Solubility has collapsed from ~13 µM to under 2 nM — a 7000-fold reduction — purely by adding an ion that contains no silver at all. Crucially, Ksp itself never changed. Equilibrium constants depend only on temperature; what moved was the equilibrium position, redistributing matter between solid and solution to keep the product of ion concentrations pinned at 1.8×10⁻¹⁰.

Working the numbers across salts

The size of the effect scales with Ksp and with how many ions the formula produces. A 1:1 salt like AgCl loses solubility in direct proportion to the common-ion concentration. A 1:2 salt like lead(II) iodide, PbI₂ ⇌ Pb²⁺ + 2 I⁻ with Ksp = 7.1×10⁻⁹, is even more sensitive to a common iodide because iodide appears squared in the expression: s = Ksp/[I⁻]². The table below makes the leverage concrete.

System	Ksp (25 °C)	Solubility in pure water	Solubility with common ion	Fold decrease
AgCl in 0.10 M NaCl	1.8×10⁻¹⁰	1.3×10⁻⁵ M	1.8×10⁻⁹ M	~7,400×
BaSO₄ in 0.10 M Na₂SO₄	1.1×10⁻¹⁰	1.0×10⁻⁵ M	1.1×10⁻⁹ M	~9,500×
PbI₂ in 0.10 M KI	7.1×10⁻⁹	1.2×10⁻³ M	7.1×10⁻⁷ M	~1,700×
CaF₂ in 0.10 M NaF	3.9×10⁻¹¹	2.1×10⁻⁴ M	3.9×10⁻⁹ M	~54,000×
Mg(OH)₂ in 0.10 M NaOH	5.6×10⁻¹²	1.1×10⁻⁴ M	5.6×10⁻¹⁰ M	~200,000×

Notice the pattern: salts whose anion (or cation) is squared in the Ksp expression — fluoride in CaF₂, hydroxide in Mg(OH)₂ — show the most violent collapse, because driving up a squared term in the denominator crushes solubility quadratically. This is why a modest 0.10 M excess of hydroxide can lower Mg(OH)₂ solubility by five orders of magnitude.

The same effect in weak acids and bases

The common-ion effect is not limited to dissolving solids. A weak acid ionizes only partly, and that ionization is an equilibrium too:

CH₃COOH ⇌ H⁺ + CH₃COO⁻, Ka = 1.8×10⁻⁵ (pKa = 4.74)

Pure 0.10 M acetic acid gives [H⁺] = √(Ka × C) ≈ √(1.8×10⁻⁶) ≈ 1.3×10⁻³ M, a pH of about 2.87, with roughly 1.3 % of the acid ionized. Now add 0.10 M sodium acetate. Acetate is the common ion shared with the acid's ionization, so it stresses the equilibrium from the product side. Dissociation is suppressed, ionization falls toward ~0.018 %, and the pH climbs. Plugging into Ka:

[H⁺] = Ka × ([acid]/[base]) = 1.8×10⁻⁵ × (0.10/0.10) = 1.8×10⁻⁵ M → pH = 4.74

That jump from pH 2.87 to 4.74, with the acid's dissociation throttled by its own conjugate base, is exactly the chemistry of a buffer. The common ion is what makes a buffer resist change: there is a large reservoir of un-ionized acid held back by the added acetate, ready to neutralize any base, and a large reservoir of acetate ready to mop up acid. Ammonia behaves symmetrically — adding ammonium chloride supplies the common NH₄⁺ ion and suppresses the ionization of NH₃, lowering [OH⁻] and giving a basic buffer.

When the effect bends or breaks

Push a common ion to high concentration and two competing phenomena spoil the simple picture. First, complex-ion formation: silver chloride that has precipitated in modest chloride will re-dissolve in a large excess of chloride because Ag⁺ binds extra Cl⁻ to form soluble AgCl₂⁻ and AgCl₃²⁻ complexes. Plot AgCl solubility against [Cl⁻] and you see a minimum near 10⁻³–10⁻² M chloride, then a rise — the common-ion effect wins at first and loses later. Second, the salt effect (or "uncommon-ion effect"): adding any inert electrolyte, even one with no shared ion, raises ionic strength, lowers the activity coefficients of the dissolved ions, and lets more salt dissolve to satisfy the activity-based Ksp. Rigorous work therefore writes Ksp in terms of activities, a = γ·[concentration], and uses Debye–Hückel theory to estimate γ. For the introductory case these corrections are small, but they explain why measured solubilities never fall quite as far as the naive calculation predicts, and why very concentrated solutions can reverse the trend.

Where it matters

• Gravimetric analysis. To weigh an analyte accurately, you precipitate it as completely as possible. Adding a slight excess of the precipitating reagent — a common ion — drives residual analyte out of solution, often leaving under 0.1 % behind, well within analytical tolerance.
• Qualitative cation separation. Classic schemes drop metal sulfides and hydroxides in groups by controlling a common ion (S²⁻ via H₂S at tuned pH, or OH⁻), so cations precipitate selectively rather than all at once.
• Buffers in blood and the lab. The bicarbonate buffer keeping blood at pH 7.4 and benchtop acetate or phosphate buffers all rely on a conjugate common ion suppressing ionization.
• Salting-out. Soap is separated from glycerol by adding NaCl: the common Na⁺ (and the ionic-strength change) forces the sodium carboxylate out of solution. Dyes and proteins are precipitated the same way.
• Industrial precipitation. The Solvay process and many metallurgical recoveries push product ions to crystallize useful solids; controlled common-ion addition maximizes yield.
• Biology and medicine. Kidney-stone formation (calcium oxalate, calcium phosphate) and tooth-enamel dissolution are common-ion problems — raising or lowering shared Ca²⁺, phosphate, or fluoride shifts whether the mineral dissolves or deposits. Fluoride toothpaste exploits the effect to stabilize enamel.

Common ion vs. uncommon ion at a glance

Aspect	Common-ion effect	Salt (uncommon-ion) effect
Added species	Shares an ion with the salt	Inert electrolyte, no shared ion
Direction on solubility	Decreases it	Slightly increases it
Mechanism	Le Chatelier shift, Q > Ksp	Lower activity coefficients (γ)
Magnitude	Large (orders of magnitude)	Small (tens of percent)
Constant affected	None — position moves	None — activities adjust