The equilibrium constant K and its expression, Le Chatelier's principle, the dissociation of weak acids in terms of Ka and pKa, the calculation of pH for weak acids, the action of buffer solutions, and the selection of indicators for titrations.

What this key area is asking

The SQA wants you to write and use the equilibrium constant $K$, apply Le Chatelier's principle, treat weak acids using $K_a$ and $pK_a$, calculate the pH of a weak acid, explain how buffer solutions resist pH change, and select a suitable indicator for a titration. The weak-acid pH calculation and the buffer explanation are reliable exam earners.

The equilibrium constant

A large value of $K$ (much greater than 1) means the equilibrium lies to the right, favouring products; a small value means it lies to the left. $K$ depends only on temperature: changing concentration or pressure shifts the position but does not change $K$, whereas changing temperature changes $K$ itself.

Le Chatelier's principle

Weak acids: Ka and pKa

A strong acid is fully dissociated in water, whereas a weak acid only partly dissociates, setting up an equilibrium:

The acid dissociation constant is $K_a = \dfrac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}$, and $pK_a = -\log_{10} K_a$. A smaller $K_a$ (larger $pK_a$) means a weaker acid. For a weak acid, the approximation $[\text{H}^+] = \sqrt{K_a c}$ holds because dissociation is slight, so the undissociated concentration is taken as the original concentration $c$.

Buffer solutions

The pH of a buffer is set by the ratio of acid to conjugate base and by the acid's $pK_a$:

When the concentrations of acid and conjugate base are equal, the pH equals the $pK_a$ of the weak acid.

Choosing an indicator

An indicator is itself a weak acid (written $\text{HIn}$) whose dissociated and undissociated forms have different colours. It changes colour over a range roughly $pK_a \pm 1$. For a titration, a suitable indicator is one whose colour-change range falls on the steep part of the pH curve at the equivalence point. For a strong acid with a weak base, methyl orange is suitable; for a weak acid with a strong base, phenolphthalein is suitable.

Examples in context

Equilibrium control runs through industry and biology. The Haber process for ammonia and the Contact process for sulfuric acid both balance temperature, pressure and catalyst to maximise yield and rate, a direct application of Le Chatelier and equilibrium. Buffers keep the pH of blood near $7.4$ using a carbonic-acid-hydrogencarbonate system, and laboratory buffers calibrate pH meters and run enzyme assays. Weak-acid behaviour explains why ethanoic acid (vinegar) and carbonic acid (in fizzy drinks) are far less corrosive than strong acids of the same concentration, and indicator choice is exactly the skill needed when designing the titrations in the Researching Chemistry area.

Try this

Q1. Write the equilibrium constant expression for $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$. [1 mark]

• Cue. $K = \dfrac{[\text{NH}_3]^2}{[\text{N}_2][\text{H}_2]^3}$.

Q2. State the effect of adding a catalyst on the position of an equilibrium. [1 mark]

• Cue. No effect on the position; it only reaches equilibrium faster.

Q3. A weak acid has $K_a = 1.0 \times 10^{-4}$. Calculate the pH of a $0.10 \text{ mol l}^{-1}$ solution. [2 marks]

• Cue. $[\text{H}^+] = \sqrt{1.0 \times 10^{-4} \times 0.10} = 3.16 \times 10^{-3}$, so $\text{pH} = 2.50$.