Explain how buffer solutions resist changes in pH using Le Chatelier's principle and conjugate acid-base equilibria

What this dot point is asking

A buffer solution resists changes in pH when small quantities of strong acid or strong base are added, and also when the solution is diluted. It is built from a conjugate acid-base pair present in comparable amounts: most often a weak acid together with a salt of its conjugate base (for example ethanoic acid, $\text{CH}_3\text{COOH}$, with sodium ethanoate, $\text{CH}_3\text{COONa}$), or a weak base with its conjugate acid (ammonia with ammonium chloride).

How a buffer works

Take an ethanoic acid / ethanoate buffer. The relevant equilibrium is

The buffer contains a large reservoir of both the weak acid molecule ($\text{CH}_3\text{COOH}$) and its conjugate base ($\text{CH}_3\text{COO}^-$).

When acid ($\text{H}^+$) is added, the extra $\text{H}^+$ reacts with the conjugate base: $\text{CH}_3\text{COO}^- + \text{H}^+ \rightarrow \text{CH}_3\text{COOH}$. By Le Chatelier's principle, the added $\text{H}^+$ pushes the equilibrium to the left, consuming most of the added protons. The conjugate base mops up the acid, so $[\text{H}^+]$ and the pH barely change.

When base ($\text{OH}^-$) is added, the hydroxide reacts with the weak acid: $\text{CH}_3\text{COOH} + \text{OH}^- \rightarrow \text{CH}_3\text{COO}^- + \text{H}_2\text{O}$. Removing $\text{H}^+$ (as the $\text{OH}^-$ consumes it) shifts the equilibrium to the right to replace it, so again the pH changes only slightly.

The buffer works because it holds a large store of both partners. Each can react with an incoming disturbance, and the equilibrium re-establishes a hydrogen ion concentration close to the original.

Buffer capacity and limits

A buffer only works for small additions. If you add enough strong acid or base to use up one of the reservoirs, the buffering collapses and pH changes rapidly. The amount of acid or base a buffer can absorb before this happens is its buffer capacity, which is greatest when the weak acid and its conjugate base are present in roughly equal concentrations.

Where buffers matter

Blood is buffered close to pH 7.4, largely by the carbonic acid / hydrogen carbonate system: $\text{H}_2\text{CO}_3 \rightleftharpoons \text{HCO}_3^- + \text{H}^+$. This keeps the pH within the narrow range enzymes need. Buffers are also used in fermentation, shampoos, and many laboratory and industrial processes where a stable pH is required.

When answering buffer questions in the WACE examination, always name both species and write the equation showing how each one reacts, then link the response to Le Chatelier's principle to earn the reasoning marks.