Solubility is not always a fixed property: for many ionic solids, changing pH changes the concentrations of ions that participate in equilibrium. Using Le Châtelier’s principle, you can predict whether lowering or raising pH increases or decreases solubility.

Core idea: link solubility to equilibrium shifts

Le Châtelier’s principle in words

For sparingly soluble salts, the dissolution equilibrium can shift if one of the ions is removed from solution by an acid–base reaction (or increased by it).

Removing a product drives more dissolution; increasing a product drives precipitation.

Solubility equilibrium language

Even without calculating, the key prediction tool is: anything that lowers an ion concentration appearing in $K_{sp}$ tends to increase solubility, and anything that raises it tends to decrease solubility.

When pH changes solubility: three common patterns

1) Salts containing a weak base anion (conjugate base of a weak acid)

These include anions like F⁻, CO₃²⁻, SO₃²⁻, S²⁻, PO₄³⁻, and many organic anions.

• Lowering pH (adding H₃O⁺) protonates the anion:

  • $\text{A}^- + \text{H}_3\text{O}^+ \rightarrow \text{HA} + \text{H}_2\text{O}$

• This removes $\text{A}^-$ from solution (a “product” of dissolution).

• By Le Châtelier, the solid dissolves more to replace $\text{A}^-$, so solubility increases in acidic solution.

Raising pH does the opposite (less protonation), leaving more $\text{A}^-$ in solution, so dissolution is less favoured and solubility decreases (or precipitation becomes more likely).

2) Metal hydroxides and salts where $\text{OH}^-$ is an ion

Many metal hydroxides are sparingly soluble: $\text{M(OH)}_n(s) \rightleftharpoons \text{M}^{n+} + n\text{OH}^-$

• Lowering pH consumes $\text{OH}^-$:

  • $\text{H}_3\text{O}^+ + \text{OH}^- \rightarrow 2\text{H}_2\text{O}$

• Removing $\text{OH}^-$ shifts dissolution right, so hydroxides become more soluble in acid.

• Raising pH directly increases $\text{OH}^-$ (common ion effect), shifting left, so solubility decreases in base.

This is a prime case of the syllabus phrase “equilibria involving … OH−.”

A log(concentration) vs pH speciation-style diagram for an iron(III)–hydroxide system, showing which aqueous hydroxo species dominate at different pH values. As pH changes, equilibria involving $\mathrm{H^+}$ and $\mathrm{OH^-}$ redistribute dissolved forms, which can strongly affect whether a metal hydroxide tends to remain dissolved or precipitate. This connects the “pH consumes $\mathrm{OH^-}$” idea to a visual map of pH-dependent equilibria. Source

3) Salts containing a weak acid cation (conjugate acid of a weak base)

Some salts dissolve to give a cation that can donate a proton (often written as BH⁺), typical for ammonium or protonated amines.

• Raising pH (adding base) deprotonates the cation:

  • $\text{BH}^+ + \text{OH}^- \rightarrow \text{B} + \text{H}_2\text{O}$

• This removes $\text{BH}^+$ from solution, so dissolution shifts right and solubility increases in basic solution.

• Lowering pH suppresses deprotonation (keeps more $\text{BH}^+$), so solubility decreases relative to basic conditions.

A quick decision checklist (qualitative, no maths)

Identify whether pH can “consume” an ion from the solid

• If an ion from the solid is a weak base, acid will consume it → acid increases solubility.

• If the solid produces $\text{OH}^-$, acid consumes it → acid increases solubility and base decreases it.

• If an ion from the solid is a weak acid (or behaves as BH⁺), base will consume it → base increases solubility.

Keep the equilibrium framing explicit

Use phrasing such as:

• “Adding $\text{H}_3\text{O}^+$ reduces $[\text{A}^-]$, so the dissolution equilibrium shifts right.”

• “Adding $\text{OH}^-$ increases a product ion, so the equilibrium shifts left and precipitation is favoured.”