Part 4: Titration curves, indicators and pH measurements

Explain the shape of a neutralisation curve and choose an indicator by its transition range rather than by a memorised colour.

Distinguish equivalence, neutrality and endpoint

Equivalence occurs when acid and base have reacted in their stoichiometric ratio. It does not always occur at pH 7. The endpoint is the observed indicator change or instrumental criterion used to estimate equivalence. A useful indicator changes within the steep part of the curve so a small volume change gives a clear endpoint.

At 25 °C a strong monobasic acid–strong base equivalence is approximately neutral. A weak acid–strong base equivalence is alkaline because A⁻ reacts with water to form some HA and OH⁻. A strong acid–weak base equivalence is acidic because the conjugate acid of the weak base donates protons.

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Read the chemistry at successive points

A weak acid titrated with a strong base begins at the pH set by Ka and initial concentration. Before equivalence, neutralisation creates a buffer. At half-equivalence, equal amounts of HA and A⁻ make [H⁺] ≈ Ka, so pH ≈ pKa. This is an approximate equilibrium relationship, not a claim that the solution is neutral.

At equivalence the weak acid has been converted mainly to its conjugate base, and a buffer formula requiring substantial remaining HA is no longer appropriate. Well beyond equivalence, excess strong alkali controls the pH: calculate excess OH⁻ moles divided by total volume, then use Kw.

For an original 25.0 cm³ sample of 0.100 mol dm⁻³ HA titrated with 0.100 mol dm⁻³ NaOH, equivalence is 25.0 cm³ and half-equivalence 12.5 cm³. If Ka = 1.80 × 10⁻⁵, pH at half-equivalence is about 4.74. At 30.0 cm³ added base, excess OH⁻ is 0.000500 mol in 0.0550 dm³; at 25 °C pH = 11.96.

Why the indicator changes colour

An indicator can be represented as a weak acid HIn with differently coloured HIn and In⁻ forms: HIn ⇌ H⁺ + In⁻. Added acid increases [H⁺] and favours HIn; added alkali removes H⁺ and favours In⁻. Its visible transition covers a pH range rather than one exact pH.

Suppose supplied data show a steep region from pH 7.8 to 10.6. An indicator with range 8.2–10.0 is suitable, while one changing at pH 3.1–4.4 is not. Justifying with the supplied curve is stronger than saying “phenolphthalein is always used for acids”. Use only a few drops so the indicator does not materially change the solution composition.

Measure a curve carefully: PAG11 skills

Calibrate the pH meter with suitable standard buffers according to its instructions, preferably bracketing the range of interest. Rinse the electrode with deionised water between solutions and gently blot excess water; do not rub the delicate glass bulb. Immerse appropriately, stir consistently and wait for a stable reading.

Pipette a known acid volume into the vessel and record pH after measured additions of base from a burette. Use smaller additions near the steep region, allowing mixing and stabilisation each time. Plot pH vertically against cumulative volume horizontally. Temperature affects both chemical equilibria and electrode response, so keep it controlled and report it.

Large additions can skip the sharp change and give a poor equivalence estimate. A dirty or poorly calibrated electrode can bias all pH readings; repeating many readings will not automatically remove that bias. The best improvement names the cause and explains how the change addresses it.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official OCR A mark allocations.

Sources