Identify conjugate acid-base pairs in proton-transfer reactions and explain amphiprotic behaviour

What this dot point is asking

The Bronsted-Lowry theory defines an acid as a proton donor and a base as a proton acceptor. Because a proton transfer involves giving from one species and accepting by another, acids and bases always come in linked conjugate pairs.

Identifying the pairs

Consider ethanoic acid reacting with water:

Ethanoic acid donates a proton to become ethanoate, so $\text{CH}_3\text{COOH}$ and $\text{CH}_3\text{COO}^-$ are one conjugate pair (acid and its conjugate base). Water accepts a proton to become hydronium, so $\text{H}_2\text{O}$ and $\text{H}_3\text{O}^+$ are the second pair (base and its conjugate acid). Every Bronsted-Lowry equation contains exactly two such pairs.

The inverse strength relationship

For a conjugate pair, acid strength and conjugate base strength are inversely related: $K_a \times K_b = K_w$. A strong acid such as HCl is essentially fully ionised, so its conjugate base $\text{Cl}^-$ is negligibly basic and does not affect pH. Conversely, the conjugate base of a weak acid (such as ethanoate from ethanoic acid) is itself a meaningful weak base. This is exactly why a salt of a weak acid, such as sodium ethanoate, gives a slightly basic solution.

Amphiprotic species

Some species can either donate or accept a proton. These are called amphiprotic (a subset of amphoteric behaviour involving proton transfer).

The classic example is the hydrogen carbonate ion, $\text{HCO}_3^-$:

• As an acid (donating a proton): $\text{HCO}_3^- \rightleftharpoons \text{CO}_3^{2-} + \text{H}^+$
• As a base (accepting a proton): $\text{HCO}_3^- + \text{H}^+ \rightleftharpoons \text{H}_2\text{CO}_3$

Water itself is amphiprotic, which is why it self-ionises. The dihydrogen phosphate ion $\text{H}_2\text{PO}_4^-$ is another common example. Amphiprotic ions are central to buffer systems such as the carbonate buffer in blood.

Why this matters

Recognising conjugate pairs lets you predict whether a salt solution is acidic, basic or neutral, explain buffer action, and write balanced proton-transfer equations confidently. Amphiprotic species are the working components of the buffers that keep biological and environmental systems stable.