Buffer solutionsAQA A-Level Chemistry: Revision notes

A buffer solution is one that maintains an approximately constant pH when small amounts of acid or base are added, or when the solution is diluted. It does not stop the pH changing altogether; the change is small compared with the change in an unbuffered solution.

A buffer works because it contains both a species that can remove added H⁺ and a species that can remove added OH⁻, each in large amounts, so a small addition barely changes the ratio of the two.

An acidic buffer contains a weak acid and the salt of that weak acid, for example ethanoic acid, CH₃COOH, and sodium ethanoate, CH₃COONa.

• The weak acid is only slightly ionised: CH₃COOH ⇌ H⁺ + CH₃COO⁻. It supplies a reservoir of undissociated acid.
• The salt is fully dissociated: CH₃COONa → Na⁺ + CH₃COO⁻. It supplies a large reservoir of conjugate base (ethanoate ions).

The salt's ethanoate ions also push the acid equilibrium to the left, so there is little ionisation of the acid.

Adding a small amount of acid: extra H⁺ reacts with the ethanoate ions, CH₃COO⁻ + H⁺ → CH₃COOH. The equilibrium shifts left and [H⁺] hardly changes.

Adding a small amount of alkali: OH⁻ reacts with H⁺, H⁺ + OH⁻ → H₂O. The equilibrium CH₃COOH ⇌ H⁺ + CH₃COO⁻ shifts right to replace H⁺, using up some CH₃COOH and making more CH₃COO⁻. [H⁺] hardly changes.

A basic buffer contains a weak base and the salt of that weak base, for example ammonia, NH₃, and ammonium chloride, NH₄Cl.

• NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (weak base, only slightly ionised)
• NH₄Cl → NH₄⁺ + Cl⁻ (fully dissociated, large reservoir of NH₄⁺)

Adding a small amount of acid: H⁺ reacts with NH₃, NH₃ + H⁺ → NH₄⁺, or reacts with OH⁻ and the equilibrium shifts right to replace it. [OH⁻] hardly changes.

Adding a small amount of alkali: OH⁻ reacts with NH₄⁺, NH₄⁺ + OH⁻ → NH₃ + H₂O. The equilibrium shifts left and [OH⁻] hardly changes.

For a weak acid HA: $K_a = \frac{[H^+][A^-]}{[HA]}$, so $[H^+] = K_a \times \frac{[HA]}{[A^-]}$.

Two assumptions are made. The salt is fully dissociated, so $[A^-]$ equals the salt concentration. The acid is only slightly ionised, and the salt suppresses this further, so $[HA]$ equals the acid concentration put in.

Worked example: a buffer contains 0.250 mol dm⁻³ ethanoic acid and 0.100 mol dm⁻³ sodium ethanoate; $K_a = 1.74 \times 10^{-5}$ mol dm⁻³.

$[H^+] = 1.74 \times 10^{-5} \times \frac{0.250}{0.100} = 4.35 \times 10^{-5}$ mol dm⁻³

pH = $-\log_{10}(4.35 \times 10^{-5})$ = 4.36

If the two solutions are mixed, the ratio of concentrations equals the ratio of moles, because both are in the same final volume. After adding a small amount of OH⁻, subtract its moles from HA and add them to A⁻, then recalculate the ratio.

Buffers keep pH steady where a change would cause damage.

• Blood: the H₂CO₃/HCO₃⁻ system keeps blood at about pH 7.4. Enzymes and proteins only work in a narrow pH range.
• Foods and drinks: buffers (such as ethanoates and citrates) control acidity, and so affect taste and preservation.
• Shampoos and skin products: buffers keep the pH close to that of skin.
• Laboratory and industry: buffers are used to calibrate pH meters and to control pH in reactions and biological work.