What is the acid ionisation constant Ka?

For a weak acid HA the ionisation equilibrium is HA H^+ + A^-, and its equilibrium constant is the acid ionisation constant:

What is the base ionisation constant Kb?

For a weak base B that reacts with water, B + H_2O BH^+ + OH^-, the base ionisation constant is

WAChemistrySyllabus dot point

What is the difference between a strong and a weak acid, and how do Ka, Kb and pKa measure acid and base strength?

Distinguish strong from weak acids and bases by degree of ionisation, and define and use Ka, Kb and pKa to compare strengths

A focused answer to the WACE Year 12 Chemistry dot point on strong versus weak acids and bases, the acid and base ionisation constants Ka and Kb, the meaning of pKa, and how they quantify strength, with a worked example and common exam mistakes.

Generated by Claude Opus 4.77 min answerUpdated 2026-05-29

Reviewed by: AI editorial process; not yet individually human-reviewed

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What this dot point is asking

Acids and bases differ in strength, which describes how completely they ionise in water, and in concentration, which describes how much is dissolved. These are independent ideas and confusing them is the single biggest error in this topic.

Strong versus weak

A strong acid such as hydrochloric acid ionises essentially completely:

\text{HCl}(aq) \rightarrow \text{H}^+(aq) + \text{Cl}^-(aq)

We write a single arrow because effectively all the HCl molecules donate their proton. A weak acid such as ethanoic acid only partially ionises, so we write an equilibrium:

\text{CH}_3\text{COOH}(aq) \rightleftharpoons \text{CH}_3\text{COO}^-(aq) + \text{H}^+(aq)

At any moment most of the ethanoic acid remains as un-ionised molecules. The same distinction applies to bases: sodium hydroxide is a strong base (fully dissociated), while ammonia is a weak base that establishes an equilibrium with water.

The acid ionisation constant Ka

For a weak acid HA the ionisation equilibrium is $\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^-$ , and its equilibrium constant is the acid ionisation constant:

K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}

A larger $K_a$ means the equilibrium lies further to the right, so the acid is stronger. Strong acids have very large $K_a$ values, so we do not usually quote them.

The base ionisation constant Kb

For a weak base B that reacts with water, $\text{B} + \text{H}_2\text{O} \rightleftharpoons \text{BH}^+ + \text{OH}^-$ , the base ionisation constant is

K_b = \frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}

A larger $K_b$ means a stronger base. For a conjugate acid-base pair, $K_a \times K_b = K_w$ , so a stronger acid has a weaker conjugate base and vice versa.

Calculating pH of a weak acid

Because a weak acid is only partly ionised, you cannot assume $[\text{H}^+]$ equals the acid concentration. You set up an ICE table and use $K_a$ . For a weak acid of concentration $c$ where the ionised amount $x$ is small compared with $c$ , the approximation $K_a \approx x^2 / c$ gives $[\text{H}^+] = x \approx \sqrt{K_a \cdot c}$ .

pH of a weak acid

Question

Calculate the pH of $0.10$ mol L $^{-1}$ ethanoic acid. $K_a = 1.8 \times 10^{-5}$ .

Solution

Let $x = [\text{H}^+]$ at equilibrium. With the approximation $[\text{CH}_3\text{COOH}] \approx 0.10$ :

K_a = \frac{x^2}{0.10} = 1.8 \times 10^{-5}

x^2 = 1.8 \times 10^{-6}, \quad x = 1.34 \times 10^{-3} \text{ mol L}^{-1}

\text{pH} = -\log_{10}(1.34 \times 10^{-3}) = 2.87

Compare this with $0.10$ mol L $^{-1}$ HCl (a strong acid), which would give pH $= 1.0$ . The weak acid is much less acidic at the same concentration because it is only partly ionised.

Why this matters

$K_a$ and $K_b$ let you rank acids and bases, calculate the pH of weak acid and base solutions, and explain buffer behaviour and the position of titration end points. They turn the qualitative idea of strength into a number you can compute with.

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