WAChemistrySyllabus dot point

How are acids and bases linked in conjugate pairs, and how can one species act as both an acid and a base?

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

A focused answer to the WACE Year 12 Chemistry dot point on conjugate acid-base pairs, how they differ by one proton, the inverse strength relationship, and amphiprotic species, with a worked example and common exam mistakes.

Generated by Claude Opus 4.76 min answerUpdated 2026-05-29

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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:

\text{CH}_3\text{COOH} + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}_3\text{O}^+

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.

Identifying conjugate pairs

Question

In the reaction $\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$ , identify the two conjugate acid-base pairs.

Solution

Ammonia accepts a proton from water to become the ammonium ion. So $\text{NH}_3$ is a base and $\text{NH}_4^+$ is its conjugate acid: that is pair one ( $\text{NH}_4^+ / \text{NH}_3$ , differing by one proton).

Water donates the proton to become hydroxide, so $\text{H}_2\text{O}$ is the acid and $\text{OH}^-$ is its conjugate base: that is pair two ( $\text{H}_2\text{O} / \text{OH}^-$ ). Here water behaves as an acid, whereas with ethanoic acid it behaved as a base, which shows water is amphiprotic.

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.

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