8.11 pH and Solubility

pH and Solubility

Sweep the pH slider from 0 to 14 and watch a saturated beaker's salt pile grow or shrink. Compare 5 salts to see when a weak-base or hydroxide anion makes solubility pH-sensitive.

Le Châtelier5 SaltspH 0–14Qualitative
Topic 8.11

pH and Solubility

Identify the qualitative effect of changes in pH on the solubility of a salt.

The solubility of a salt is pH sensitive when one of its constituent ions is a weak acid, a weak base, or the hydroxide ion. These effects can be understood qualitatively using Le Châtelier's principle.

The mechanism. Consider a salt whose anion is a weak base:

CaF₂(s) ⇌ Ca²⁺(aq) + 2 F⁻(aq)
F⁻(aq) + H₃O⁺(aq) ⇌ HF(aq) + H₂O(l)

Adding acid consumes F⁻ by converting it to HF. That removes a product from the first equilibrium, which shifts right, so more CaF₂ dissolves. In Q terms: consuming F⁻ drops Qsp below Ksp, and the system responds by dissolving more solid.

Which salts are pH sensitive? Those whose anion is the conjugate base of a weak acid, so it can actually grab a proton:

  • Sensitive: F⁻, CH₃COO⁻, CO₃²⁻, PO₄³⁻, S²⁻, OH⁻, CN⁻ — solubility increases in acid.
  • Insensitive: Cl⁻, Br⁻, I⁻, NO₃⁻, ClO₄⁻ — these are conjugate bases of strong acids, far too weak to be protonated, so pH has essentially no effect.

Hydroxide salts are the most dramatic case. Mg(OH)₂ dissolves readily in acid because H₃O⁺ consumes OH⁻ directly. Conversely, raising the pH adds OH⁻, which is a common ion, and suppresses the solubility of any metal hydroxide.

Carbonates are the everyday example: CaCO₃ in limestone, marble, and seashells dissolves in acid because CO₃²⁻ is protonated to HCO₃⁻ and then to H₂CO₃, which decomposes to CO₂ gas. Acid rain damage to marble monuments is exactly this reaction.

Key points

  • Adding acid increases the solubility of salts whose anion is a weak base.
  • Salts of strong-acid anions (Cl⁻, NO₃⁻) show essentially no pH dependence.
  • Metal hydroxides dissolve in acid and are suppressed by added base (common-ion effect).
  • The reasoning is Le Châtelier’s principle applied to two coupled equilibria.

Common mistakes

  • Claiming acid dissolves everything. AgCl is essentially unaffected by pH.
  • Thinking Ksp changes with pH. Only temperature changes Ksp; what changes is the ion concentration.
  • Forgetting that raising pH suppresses hydroxide solubility — that is a common-ion effect, not a new principle.
  • Reasoning from only one equilibrium. The explanation requires both the dissolution and the acid–base equilibrium.

Worked example

Predict and explain the effect of lowering the pH on the solubility of (a) CaCO₃, (b) AgCl, and (c) Mg(OH)₂.

(a) CaCO₃ — solubility increases substantially.
CaCO₃(s) ⇌ Ca²⁺(aq) + CO₃²⁻(aq)
Carbonate is the conjugate base of the weak acid HCO₃⁻, so added H₃O⁺ protonates it:
CO₃²⁻ + H₃O⁺ → HCO₃⁻ + H₂O, and further HCO₃⁻ + H₃O⁺ → H₂CO₃ → CO₂(g) + H₂O
Removing carbonate from solution drives the dissolution equilibrium to the right, so more CaCO₃ dissolves. The escaping CO₂ gas removes product permanently, pushing the reaction essentially to completion — this is why limestone visibly fizzes in acid.

(b) AgCl — essentially no effect.
AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq)
Chloride is the conjugate base of HCl, a strong acid, which means Cl⁻ is an extremely weak base and is not measurably protonated at any accessible pH. Since neither ion is removed from solution, the dissolution equilibrium is undisturbed and solubility is unchanged.

(c) Mg(OH)₂ — solubility increases greatly.
Mg(OH)₂(s) ⇌ Mg²⁺(aq) + 2 OH⁻(aq)
Added H₃O⁺ neutralizes hydroxide directly: H₃O⁺ + OH⁻ → 2 H₂O. Removing OH⁻ shifts the dissolution equilibrium sharply to the right. This is why milk of magnesia, a suspension of Mg(OH)₂, dissolves in stomach acid and neutralizes it.

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