Selecting Indicators

Spec Mapping — OCR H432 Module 5.1.3 — Acids, bases and buffers, sub-topic on indicator selection: the indicator as a weak acid HIn with two different-coloured conjugate forms, the indicator dissociation constant $K_{In}$KIn​ and its logarithm $\text{p}K_{In}$pKIn​, the colour-change range of approximately $\text{p}K_{In} \pm 1$pKIn​±1, the matching rule that requires $\text{p}K_{In}$pKIn​ to fall within the vertical jump of the titration curve, the unsuitability of universal indicator for endpoint determination, and the contextual link to titration curves from lesson 9 (refer to the official OCR H432 specification document for exact wording). This is the final lesson of Module 5.1.3 — it brings the curve geometry of lesson 9 into a single decision criterion: indicator range vs vertical jump.

This is the concluding lesson of the acids/bases/buffers module — the place where the titration-curve picture (lesson 9) and the $K_a$Ka​ machinery (lesson 4) combine to give a single operational rule: the indicator's colour-change range must coincide with the vertical jump of the titration curve. The lesson covers the chemistry of an indicator (HIn as a weak acid with different-coloured conjugate forms), the quantitative relationship $\text{p}H = \text{p}K_{In}$pH=pKIn​ at the colour midpoint, the four canonical indicator-choice cases (one for each curve shape), why universal indicator is useless for endpoints, and a brief mention of how calorimetric titration ($\Delta H_{\text{neut}}$ΔHneut​) can replace indicators in difficult cases.

Key Equation: $\text{HIn(aq)} \rightleftharpoons \text{H}^+\text{(aq)} + \text{In}^-\text{(aq)}$HIn(aq)⇌H+(aq)+In−(aq) $K_{In} = \frac{[\text{H}^+][\text{In}^-]}{[\text{HIn}]} \quad \Rightarrow \quad \text{at colour midpoint: } [\text{HIn}] = [\text{In}^-] \Rightarrow \text{pH} = \text{p}K_{In}$KIn​=[HIn][H+][In−]​⇒at colour midpoint: [HIn]=[In−]⇒pH=pKIn​ Colour-change range ≈ $\text{p}K_{In} \pm 1$pKIn​±1. The indicator must have its range entirely within the vertical jump of the titration curve.

---

What is an indicator?

An acid-base indicator is a weak acid (or weak base) whose protonated and deprotonated forms have different colours. Writing the indicator generically as HIn:

$\text{HIn(aq)} \rightleftharpoons \text{H}^+\text{(aq)} + \text{In}^-\text{(aq)}$HIn(aq)⇌H+(aq)+In−(aq) $\text{colour A} \qquad\qquad\qquad \text{colour B}$colour Acolour B

HIn is colour A; In⁻ is colour B. The equilibrium position depends on [H⁺] (Le Chatelier):

• In acidic solution (high [H⁺]), the equilibrium is pushed to the left → HIn predominates → colour A.
• In alkaline solution (low [H⁺]), the equilibrium shifts to the right → In⁻ predominates → colour B.

The indicator therefore changes colour as pH passes through a specific range.

Quantifying the change — $K_{In}$KIn​ and $\text{p}K_{In}$pKIn​

The indicator has its own dissociation constant:

$K_{In} = \frac{[\text{H}^+][\text{In}^-]}{[\text{HIn}]}$KIn​=[HIn][H+][In−]​

At the colour midpoint $[\text{HIn}] = [\text{In}^-]$[HIn]=[In−]:

$K_{In} = [\text{H}^+] \quad \Rightarrow \quad \text{pH} = \text{p}K_{In}$KIn​=[H+]⇒pH=pKIn​

So the midpoint of the colour change is at $\text{pH} = \text{p}K_{In}$pH=pKIn​. The eye perceives a single dominant colour once one species is about 10× more abundant than the other, so the practical colour-change range is approximately:

$\text{p}K_{In} - 1 \le \text{pH} \le \text{p}K_{In} + 1$pKIn​−1≤pH≤pKIn​+1

A typical indicator therefore changes colour over about 2 pH units.

Common indicators

Indicator	$\text{p}K_{In}$pKIn​	Range	Acid → Alkali colour
Methyl orange	3.7	3.1 – 4.4	red → yellow
Bromophenol blue	4.0	3.0 – 4.6	yellow → blue
Bromocresol green	4.7	3.8 – 5.4	yellow → blue
Methyl red	5.1	4.4 – 6.2	red → yellow
Bromothymol blue	7.0	6.0 – 7.6	yellow → blue
Phenol red	7.9	6.8 – 8.4	yellow → red
Phenolphthalein	9.3	8.3 – 10.0	colourless → pink
Thymolphthalein	9.9	9.3 – 10.5	colourless → blue

You should know methyl orange and phenolphthalein in detail (their colours, ranges, and standard uses) and be able to use any supplied table. OCR does not require memorisation of the full set.

---

Matching indicators to titration curves — the four cases

The rule is simple and absolute: the indicator's colour-change range (and ideally its $\text{p}K_{In}$pKIn​) must fall entirely within the vertical jump of the titration curve.

Strong acid + strong base (curve 1)

Vertical jump: pH 3.5 → 10.5 (about 7 units). Both methyl orange (3.1–4.4) and phenolphthalein (8.3–10.0) fall within this range. Either works. In practice phenolphthalein is preferred for visual clarity — the colourless → bright pink transition is more obvious than methyl orange's red → yellow.

Weak acid + strong base (curve 2)

Equivalence at pH ~8.7 (e.g. ethanoic + NaOH). Vertical jump: pH 7 → 11 (about 4 units). Phenolphthalein (8.3–10.0) is the correct choice. Methyl orange (3.1–4.4) is wrong — it would change colour in the buffer region of the curve (pH 4–6, where pH changes only slowly), giving a premature, gradual end-point.

Strong acid + weak base (curve 3)

Equivalence at pH ~5.3 (e.g. HCl + NH₃). Vertical jump: pH 3 → 7. Methyl orange (3.1–4.4) is the correct choice. Phenolphthalein (8.3–10.0) is wrong — it would change long after the equivalence point, in the post-equivalence NH₃/NH₄⁺ buffer region.

Weak acid + weak base (curve 4)

No vertical jump exists. No indicator covers a region where the pH is changing fast enough to give a sharp colour change. The colour change would spread over many cm³ of titrant — useless for analytical purposes. For weak-weak titrations, use a pH meter with a glass electrode (PAG 11) or a conductometric titration.

flowchart TD
  A["Identify acid + base types"] --> B{"Both strong?"}
  B -- "Yes" --> C["Either methyl orange OR phenolphthalein; prefer phenolphthalein"]
  B -- "No" --> D{"Weak acid + strong base?"}
  D -- "Yes" --> E["pH at eq > 7 → phenolphthalein"]
  D -- "No" --> F{"Strong acid + weak base?"}
  F -- "Yes" --> G["pH at eq < 7 → methyl orange"]
  F -- "No" --> H["Weak/weak: no indicator works; use pH meter"]

---

Worked examples

Worked example 1 — choosing an indicator

Titrating 25.0 cm³ of 0.100 mol dm⁻³ ethanoic acid with 0.100 mol dm⁻³ NaOH. Which indicator should you use?

Equivalence at pH ~8.72 (lesson 9). Vertical jump: pH 7 → 11. An indicator with $\text{p}K_{In}$pKIn​ near 9 and range fitting 8–10 is needed. Phenolphthalein (range 8.3–10.0) is ideal. Methyl orange (3.1–4.4) is wrong — it falls in the buffer region.

Worked example 2 — why methyl orange fails for weak-acid + strong-base

Methyl orange changes colour between pH 3.1 and 4.4. In a weak acid / strong base titration the equivalence pH is above 7; the pH 3–5 region is the buffer region of the curve, where the pH changes only slowly. Methyl orange would change colour gradually, around 1–2 cm³ of NaOH added (the start of the buffer region), long before the 25 cm³ equivalence point. The end-point would be premature and indistinct.

Worked example 3 — strong acid + weak base

Titrating 25.0 cm³ of 0.100 mol dm⁻³ HCl with 0.100 mol dm⁻³ NH₃. Equivalence at pH ~5.28. Vertical jump: pH 3 → 7. Methyl orange (3.1–4.4) falls inside; phenolphthalein (8.3–10.0) is wrong.

Worked example 4 — exam-style selection

In a titration the pH at equivalence is 8.50 and the vertical region of the curve spans pH 7.0 → 10.5. From the following indicators, which is most suitable? (a) methyl red ($\text{p}K_{In} = 5.1$pKIn​=5.1, range 4.4–6.2); (b) bromothymol blue ($\text{p}K_{In} = 7.0$pKIn​=7.0, range 6.0–7.6); (c) phenolphthalein ($\text{p}K_{In} = 9.3$pKIn​=9.3, range 8.3–10.0); (d) thymolphthalein ($\text{p}K_{In} = 9.9$pKIn​=9.9, range 9.3–10.5).

Check each range against pH 7.0–10.5:

• (a) methyl red 4.4–6.2: entirely outside — colour change before the jump. Unsuitable.
• (b) bromothymol blue 6.0–7.6: partially outside (below 7.0). Unsuitable.
• (c) phenolphthalein 8.3–10.0: fully inside, $\text{p}K_{In}$pKIn​ near the centre. Suitable — best choice.
• (d) thymolphthalein 9.3–10.5: fully inside but upper end at the boundary. Suitable, less ideal than phenolphthalein.

Best choice: phenolphthalein.

Worked example 5 — titration curve overlay with three indicator ranges

Sketch (mentally) a strong-acid / strong-base titration curve (HCl titrated with NaOH, both 0.100 mol dm⁻³, equivalence at 25.0 cm³). The curve runs from pH 1 → 13, with the steep vertical jump from approximately pH 3.5 to pH 10.5 occurring across the addition of 24.95 → 25.05 cm³ NaOH (about 0.1 cm³ total). Now overlay three indicator ranges as horizontal coloured bands:

• Methyl orange (3.1–4.4) — sits at the bottom of the vertical jump. Colour change occurs at ~24.95 cm³ NaOH, which is 99.8 % of the way to equivalence. Error: 0.2 %.
• Bromothymol blue (6.0–7.6) — sits centrally. Colour change occurs at ~25.00 cm³ NaOH, exactly at equivalence. Error: < 0.1 %.
• Phenolphthalein (8.3–10.0) — sits at the top of the vertical jump. Colour change at ~25.05 cm³, i.e. 100.2 %. Error: 0.2 %.

All three are acceptable for SA/SB because the vertical jump is wide enough (~7 pH units) to accommodate any of their ranges. The errors are tiny (≤ 0.2 %) and well below typical burette reading precision (~0.05 cm³ in 25 cm³ = 0.2 %). For weak-acid + strong-base (vertical jump pH 7 → 11), only the upper two would work; methyl orange would now sit outside the jump, in the buffer region — colour change would occur at ~10–15 cm³ NaOH (well before equivalence), giving an error of 40–60 %. This direct geometric overlay is the surest mental model for indicator selection: imagine the indicator band sliding up and down the pH axis; only when the band lies entirely within the steep vertical portion is the indicator valid.

Worked example 6 — why no indicator works for weak/weak

Consider a titration of 0.100 mol dm⁻³ ethanoic acid ($\text{p}K_a$pKa​ = 4.76) with 0.100 mol dm⁻³ aqueous ammonia ($\text{p}K_b$pKb​ = 4.75, $\text{p}K_a$pKa​(NH₄⁺) = 9.25). The equivalence point pH for this weak-weak titration sits near 7.00 (because the salt CH₃COONH₄ is approximately neutral — the cation's acidity nearly cancels the anion's basicity). However, the curve has no steep vertical jump — instead the pH rises continuously and gradually across the entire titration, with no inflection sharper than about 1 pH unit per 1 cm³. None of the standard indicator ranges (typically 2 pH units wide) coincides with such a slow rise; any indicator would change colour gradually over many cm³ of added base, producing an indistinct streaky colour transition rather than a single drop's worth of sharp shift. The diagnostic test: if the steepest portion of the curve has $|d\text{pH}/dV| < 2$∣dpH/dV∣<2 pH/cm³, no visual indicator can give a sharp endpoint. Use a glass-electrode pH meter and locate the equivalence point by the inflection point of the recorded pH-vs-V curve (first-derivative maximum).

---

Why universal indicator does not work for endpoints

Universal indicator is a mixture of dyes that gives a different colour for each pH unit between about 2 and 12. It is useful for estimating the pH of a static solution to the nearest unit or two (the rainbow shift from red → orange → yellow → green → blue → purple gives a coarse pH read-out). However, it does not give a sharp colour change at any specific pH.

In a titration, universal indicator would drift gradually through the rainbow as the pH rises continuously — there is no definite end-point. Even at the steep equivalence region, the colour shift is spread over several pH units and several cm³ of titrant.