Indicators and Weak/Strong Acid-Base Reactions Revision Notes | A Level AQA Chemistry

1. Why Indicator Choice Is Not Arbitrary

Picture this: you add exactly the right volume of sodium hydroxide to ethanoic acid, but your indicator has not changed colour yet. You keep adding base, overshoot the equivalence point, and your result is wrong. The culprit is not your technique — it is a mismatched indicator.

Every indicator changes colour over a characteristic pH range, roughly spanning 2 pH units. For the indicator to signal the endpoint reliably, that entire colour-change range must fall within the steep region of the pH curve. If any part of the indicator's range sits outside the steep region, the colour change will be gradual and the endpoint will be ambiguous.

This is where pH curves become your most powerful tool. By reading the steep region from the curve, you can immediately judge which indicators are acceptable and which will mislead you.

2. The Four Types of Acid-Base Titration

There are four possible combinations of strong and weak acids and bases. Each produces a pH curve with a distinctive shape, and crucially, a different steep region — or in one case, no steep region at all.

The weak acid–weak base combination is the critical exception. Without a steep region, no indicator can give a sharp, reliable endpoint, so this titration cannot be performed accurately using a simple indicator. AQA expects you to know this and to justify it in terms of the absence of a steep region.

3. Reading pH Curves to Select an Indicator

A pH curve plots pH on the vertical axis against the volume of base added on the horizontal axis. The steep region is the near-vertical section of the S-shaped curve. To select an indicator, you identify the pH range of the steep region and then check whether the indicator's colour-change interval sits entirely within it.

Notice how the strong acid–strong base curve has the widest steep region, spanning roughly pH 3 to pH 11. Both methyl orange (colour change approximately pH 3.1–4.4) and phenolphthalein (colour change approximately pH 8.2–10.0) fall entirely within this range, so either is acceptable.

For the weak acid–strong base curve, the steep region begins at around pH 7 and extends to approximately pH 11. Phenolphthalein's range (pH 8.2–10.0) sits comfortably within this region. Methyl orange changes colour between pH 3.1 and 4.4 — well below the steep region — so it would change colour gradually during the buffer region of the curve, giving a diffuse, unreliable endpoint.

For the strong acid–weak base curve, the steep region runs from roughly pH 3 to pH 7. Methyl orange (pH 3.1–4.4) falls within this range. Phenolphthalein (pH 8.2–10.0) lies entirely above the steep region, so it would never give a sharp colour change at the equivalence point.

4. Weak Acid Reacting with a Strong Base — pH Changes in Detail

Take ethanoic acid ($\text{CH}_3\text{COOH}$) being titrated with sodium hydroxide ($\text{NaOH}$). Because ethanoic acid is a weak acid, it only partially ionises in solution. This has a profound effect on the shape of the pH curve compared with a strong acid.

At the start, before any base is added, the pH is higher than you would expect for the same concentration of a strong acid. For a 0.1 mol dm$^{-3}$ solution of ethanoic acid, the pH is approximately 2.9 rather than 1.0, because only a small fraction of the molecules have ionised to release $\text{H}^+$ ions.

As sodium hydroxide is added, it reacts with the ethanoic acid, converting it to ethanoate ions. A buffer region develops because the solution now contains both the weak acid and its conjugate base (ethanoate ions), which resists pH change. This buffer region is a distinctive feature of weak acid–strong base curves and is absent from strong acid–strong base curves.

At the half-equivalence point — when exactly half the acid has been neutralised — the pH equals the $\text{p}K_a$ of the weak acid. For ethanoic acid, $\text{p}K_a = 4.76$, so the pH at the half-equivalence point is 4.76. This is a useful relationship that AQA may ask you to apply.

At the equivalence point, all the ethanoic acid has been converted to sodium ethanoate ($\text{CH}_3\text{COONa}$). The ethanoate ion is the conjugate base of a weak acid, so it hydrolyses water slightly: $\text{CH}_3\text{COO}^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{CH}_3\text{COOH}(\text{aq}) + \text{OH}^-(\text{aq})$. This produces a small excess of hydroxide ions, making the solution alkaline at the equivalence point — typically around pH 8.7 for 0.1 mol dm$^{-3}$ solutions.

Beyond the equivalence point, excess sodium hydroxide dominates and the pH rises steeply, then levels off as the solution becomes strongly alkaline. The steep region for this titration spans roughly pH 7–11, which is why phenolphthalein is the correct choice.

5. Strong Acid Reacting with a Weak Base — pH Changes in Detail

Now consider the reverse situation: hydrochloric acid ($\text{HCl}$) being titrated with ammonia solution ($\text{NH}_3$). This time the base is weak, and the pattern of the pH curve mirrors — but does not reverse — the weak acid case.

At the start, before any ammonia is added, the pH is low and determined entirely by the strong acid. For 0.1 mol dm$^{-3}$ $\text{HCl}$, the pH is 1.0, because hydrochloric acid is fully ionised. This is identical to the start of a strong acid–strong base titration.

As ammonia is added, it reacts with the hydrochloric acid: $\text{NH}_3(\text{aq}) + \text{HCl}(\text{aq}) \rightarrow \text{NH}_4\text{Cl}(\text{aq})$. A buffer region develops here too, because the solution contains a mixture of ammonia (weak base) and ammonium ions (its conjugate acid). The pH rises slowly through this region.

At the equivalence point, all the hydrochloric acid has been converted to ammonium chloride ($\text{NH}_4\text{Cl}$). The ammonium ion is the conjugate acid of a weak base, so it donates a proton to water: $\text{NH}_4^+(\text{aq}) \rightleftharpoons \text{NH}_3(\text{aq}) + \text{H}^+(\text{aq})$. This produces a small excess of hydrogen ions, making the solution acidic at the equivalence point — typically around pH 5.3 for 0.1 mol dm$^{-3}$ solutions.

Beyond the equivalence point, excess ammonia raises the pH steeply and then the curve levels off. The steep region spans roughly pH 3–7, which is why methyl orange is the correct indicator for this titration. Phenolphthalein would not change colour within the steep region and would give no reliable endpoint.

6. Comparing the Four pH Curves — Key Differences

Now that you have seen each curve individually, it is worth placing them side by side. The differences are systematic and examiners test them directly with 'compare' and 'describe' command words.

The starting pH is the first clue when identifying a curve: a weak acid starts higher than a strong acid at the same concentration, because partial ionisation means fewer $\text{H}^+$ ions are present initially. Similarly, the end pH (excess base region) is lower for a weak base than for a strong base, because ammonia is less alkaline than sodium hydroxide at the same concentration.

7. Investigating pH Changes — The Required Practical Context

AQA requires you to have investigated how pH changes when a weak acid reacts with a strong base and when a strong acid reacts with a weak base. The practical uses a pH metre (or a calibrated pH probe) rather than an indicator, because a metre gives continuous, quantitative data across the full titration — not just a colour change at one point.

A pH metre is preferred over a single indicator in this investigation because it gives a continuous record of pH throughout the titration, allowing the full shape of the curve — including the buffer region and the precise location of the equivalence point — to be determined. An indicator only signals one point on the curve.

8. Exam Technique — Answering pH Curve and Indicator Questions

pH curve questions appear regularly on AQA papers and reward students who structure their answers precisely. The command words 'describe', 'explain', 'compare', and 'evaluate' each demand a different response structure.

Worked example: How to Answer: Select and Justify an Indicator from a pH Curve

A student titrates 25.0 cm³ of 0.1 mol dm⁻³ ethanoic acid with 0.1 mol dm⁻³ sodium hydroxide. The pH curve shows a steep region from pH 7.5 to pH 11.0. The student has access to methyl orange (colour change pH 3.1–4.4) and phenolphthalein (colour change pH 8.2–10.0). Select the appropriate indicator and justify your choice. [2 marks]

1. Steep region: pH 7.5 to pH 11.0 — Read this directly from the question or the graph. This is the range over which a suitable indicator must change colour.
2. Methyl orange: pH 3.1–4.4. Does 3.1–4.4 fall within 7.5–11.0? No. — The entire colour-change range of methyl orange lies below the steep region. It would change colour gradually in the buffer region, giving an unreliable endpoint.
3. Phenolphthalein: pH 8.2–10.0. Does 8.2–10.0 fall within 7.5–11.0? Yes. — The entire colour-change range of phenolphthalein lies within the steep region. A sharp, reliable colour change will occur at the endpoint.
4. Phenolphthalein is the appropriate indicator. Its colour-change range (pH 8.2–10.0) falls entirely within the steep region of the pH curve (pH 7.5–11.0), so it will give a sharp colour change at the equivalence point. — [1 mark] for naming phenolphthalein; [1 mark] for stating that its colour-change range falls within the steep region of the curve.

Answer: Phenolphthalein; its colour-change range (pH 8.2–10.0) falls entirely within the steep region of the pH curve (pH 7.5–11.0).

When the command word is 'compare', you must address both similarities and differences explicitly. A common error is to describe only one curve and ignore the other. Structure your answer as: 'Both curves... However, the weak acid curve... whereas the strong acid curve...'

When the command word is 'explain', link every observation to a chemical reason. Do not state that the pH is higher at the equivalence point without explaining that the salt formed is the conjugate base of a weak acid and hydrolyses water to produce $\text{OH}^-$ ions.

9. Consolidation and Retrieval Check

You have now covered both specification requirements: using pH curves to select indicators, and describing how pH changes for weak acid–strong base and strong acid–weak base reactions. The table below brings together the key indicator data you need to recall for AQA examinations.

Mastery of this topic rests on two linked ideas: the shape of the pH curve determines the steep region, and the steep region determines which indicator is suitable. Every question on indicators and weak/strong acid-base reactions traces back to these two principles.

Treat indicator choice as an evidence question: the correct indicator is the one whose transition range lies within the rapid pH change region of the reaction or titration being observed.