Lesson 2 · 25 min

The common-ion effect and pH

Add a common ion and solubility plunges; protonate the anion and it soars. Limestone caves, tooth decay and ocean acidification are all stories about these two effects.

In this lesson

Quantify the common-ion effect with Le Chatelier’s principle and Ksp
Explain why salts of weak-acid anions — hydroxides, carbonates, fluorides — dissolve more in acid
Combine Ksp with acid–base fractions to calculate solubility at a given pH
Use both effects to explain limestone caves, tooth decay and ocean acidification

Limestone caves are eaten out by rain: water carrying dissolved CO₂ is slightly acidic and dissolves calcium carbonate bit by bit. Tooth enamel slowly dissolves in the acid bacteria make from sugar, and fluoride toothpaste makes it more acid-resistant. Conversely, analysts routinely add a large excess of one ion to drive a precipitation to completion. Solubility is not a fixed number — it shifts with the other ions present and with pH.

Predict first

AgCl dissolves to 1.3 × 10⁻⁵ mol/L in pure water. Roughly what is its solubility in 0.10 M NaCl?

Pick an answer and lock it in before you look. A wrong guess teaches more than no guess.

Unchanged, still 1.3 × 10⁻⁵ M About 1.3 × 10⁻⁶ M (ten times less) About 1.8 × 10⁻⁹ M (about 7000 times less) Larger: there are more ions in solution

The common-ion effect

Adding a soluble electrolyte that shares an ion with a sparingly soluble salt shifts the dissolution equilibrium towards the solid and lowers the solubility — the common-ion effect. When the added ion far exceeds the solubility, the arithmetic is easy: treat its concentration as a constant in Ksp.

This is exactly how analysts make a precipitation complete: a modest excess of precipitant leaves less of the target ion behind.

s \approx \frac{K_\mathrm{sp}}{[\mathrm{B}]_0^{\,y}} \quad (\text{for } \mathrm{AB}_y \text{ with added } \mathrm{B})

The approximation when the added common ion [B]₀ far exceeds the solubility; the bench curves are exact.

$[\mathrm{B}]_0$: concentration of the added common ion, mol/L
$y$: number of that ion per formula unit

pH and solubility

If the anion of a sparingly soluble salt is the conjugate base of a weak acid ( $\ce{OH-}$ , $\ce{F-}$ , $\ce{CO3^2-}$ , $\ce{C2O4^2-}$ , $\ce{PO4^3-}$ ), acid picks it up: $\ce{F-} + \ce{H+} \mathrel{\ce{->}} \ce{HF}$ . The free anion drops, Q falls below Ksp and more solid dissolves. With $\alpha$ the fraction of the dissolved anion still free (set by the acid–base equilibrium), $K_\mathrm{sp} = [\ce{Ca^2+}](\alpha\,c_{\ce{F}})^2$ .

$\ce{Cl-}$ , $\ce{Br-}$ and $\ce{I-}$ are conjugate bases of strong acids and take up no protons, so AgCl does not become more soluble in acid.

Solubility coupled to acid–base equilibria

The common-ion effect and pH-dependent solubility

Solution pH (held by a buffer) · 2.0

Mg(OH)₂
CaCO₃
CaF₂
CaC₂O₄
BaSO₄
AgCl
In pure water (sets its own pH)

$\ce{CaF2(s) <=> Ca^2+ + 2F-}$ The more acidic, the more the anion is protonated: $\ce{HF <- F-}$

Solubility at this pH: 0.00134 mol/L
Mass dissolved per litre: 105 mg
Fraction of anion left free, α: 0.0633
In pure water (sets its own pH): 2.13 × 10⁻⁴ mol/L · pH 7.11
Compared with pure water: 6.29 ×

Each curve solves Ksp together with the acid–base equilibria (a buffer holds the pH). CaF₂ becomes much more soluble below pH 3; CaCO₃ gains more than fifty-fold between pH 9 and pH 6; AgCl is a flat line. The black dot is the salt dissolved in pure water, setting its own pH.

Worked example

Calcium fluoride in acid

Find the solubility of CaF₂ ( $K_\mathrm{sp} = 3.9\times10^{-11}$ ) in a buffer at pH 2.00 (HF has $\mathrm{p}K_\mathrm{a} = 3.17$ ).

Concept check

Which salt is far more soluble in dilute hydrochloric acid than in pure water?

\ce{AgCl}

\ce{AgBr}

\ce{CaCO3}

\ce{PbCl2}

Try the calculation

Find the solubility (mol/L) of AgCl ( $K_\mathrm{sp} = 1.82\times10^{-10}$ ) in 0.050 M NaCl.

Key takeaways

Common-ion effect: adding a shared ion shifts the equilibrium left; s ≈ Ksp/[B]₀ʸ.
Salts of weak-acid anions dissolve more in acid because the anion is protonated; salts of strong-acid anions (Cl⁻, Br⁻, I⁻) ignore pH.
Quantitatively, combine Ksp with the acid–base fraction α: $K_\mathrm{sp} = [\mathrm{M}](\alpha\,c_\text{A})^y$ .
If the common ion also forms complexes, solubility climbs again at high concentration.

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Data and references

D. C. Harris, Quantitative Chemical Analysis, 9th ed. (2016), Ch. 6 (common-ion effect; complex formation with PbI₂) and Ch. 13 (coupled equilibria).
A. E. Martell & R. M. Smith, Critical Stability Constants (Plenum, 1974–1989), via D. Harvey, Analytical Chemistry 2.1, Appendices 10–12.
W. Stumm & J. J. Morgan, Aquatic Chemistry, 3rd ed. (Wiley, 1996), Ch. 7 (carbonate solubility).
IPCC, Climate Change 2021: The Physical Science Basis, Ch. 5 (ocean acidification).