How to Choose the Right Indicator for a Titration

Pick the wrong indicator and your titration will look perfectly neat, with sharp colour changes and concordant titres, and still give the wrong answer. That’s what makes indicator choice so important. The mistake doesn’t announce itself.

The good news is that choosing correctly follows a simple rule, and once you understand where the rule comes from, you’ll never need to memorise a table again.

The rule

Choose an indicator whose colour-change range lies within the steep, vertical section of the titration curve.

In practice that means: the indicator should change colour at, or very close to, the pH at the equivalence point.

Why the rule works

An indicator is itself a weak acid, usually written HIn. Its acid form and its conjugate base have different colours:

HIn ⇌ H⁺ + In⁻ (colour 1) (colour 2)

When [HIn] and [In⁻] are equal, pH = pKa of the indicator (often written pKin). Your eye sees a clear change in colour when one form outnumbers the other by about ten to one, so the visible change happens over a range of roughly pKin ± 1.

Near the equivalence point of a good titration, the pH jumps by several units in a single drop. If the indicator’s range sits inside that jump, the colour flips in that one drop, and the end point (colour change) matches the equivalence point (exact neutralisation). If the range sits outside the jump, the colour drifts gradually over several cm³, or changes too early or too late. See equivalence point vs end point.

Common indicators

Indicator	pH range of colour change	Colour in acid	Colour in alkali	pKin (approx.)
Thymol blue (first change)	1.2–2.8	red	yellow	1.7
Methyl orange	3.1–4.4	red	yellow	3.7
Bromophenol blue	3.0–4.6	yellow	blue	4.0
Methyl red	4.4–6.2	red	yellow	5.1
Bromothymol blue	6.0–7.6	yellow	blue	7.1
Phenol red	6.8–8.4	yellow	red	7.9
Phenolphthalein	8.2–10.0	colourless	pink	9.4
Thymolphthalein	9.3–10.5	colourless	blue	9.9

Values vary slightly between sources, and the colour change depends a little on concentration and on your eyes. The ranges above are typical.

The four-step method

Step 1: identify the type of titration

Decide whether the acid and the base are strong or weak. (If you’re unsure, see the strong acids and the strong bases; nearly everything else is weak.)

Step 2: estimate the pH at equivalence

• Strong acid + strong base: pH 7. The salt formed (e.g. NaCl) is neutral.
• Weak acid + strong base: above 7, typically 8–9. The salt’s anion (e.g. ethanoate) is a weak base.
• Strong acid + weak base: below 7, typically 5–6. The salt’s cation (e.g. NH₄⁺) is a weak acid.
• Weak acid + weak base: close to 7 but with no steep section.

Step 3: estimate the vertical section

For 0.1 mol/dm³ solutions:

• Strong/strong: about pH 3.5 to 10.5
• Weak acid/strong base: about pH 7 to 11
• Strong acid/weak base: about pH 3 to 7

The titration curves article shows where these numbers come from.

Step 4: pick an indicator whose range fits inside

Titration	Equivalence pH	Suitable	Unsuitable
HCl + NaOH	7	methyl orange, methyl red, bromothymol blue, phenolphthalein	thymol blue (first range)
CH₃COOH + NaOH	about 8.7	phenolphthalein, thymolphthalein	methyl orange, methyl red
HCl + NH₃	about 5.3	methyl orange, methyl red, bromophenol blue	phenolphthalein
CH₃COOH + NH₃	about 7	none reliable	all (use a pH meter)

Worked example 1

Which indicator should be used to titrate methanoic acid with potassium hydroxide?

• Methanoic acid is a weak acid; potassium hydroxide is a strong base.
• Equivalence pH is above 7 (methanoate is a weak base).
• The vertical section runs from about 7 to 11.
• Phenolphthalein (8.2–10.0) fits. Methyl orange (3.1–4.4) would change in the buffer region, long before equivalence.

Worked example 2

A student titrates ammonia solution with hydrochloric acid using phenolphthalein. Explain why their result will be inaccurate.

• Strong acid + weak base: equivalence pH is about 5.3.
• Phenolphthalein changes between pH 10.0 and 8.2.
• As acid is added, the pH falls through the buffer region (around 10 down to 8.5) well before equivalence, so the pink colour fades gradually and disappears too early.
• The titre will be too small, giving a concentration of ammonia that’s too low. Methyl orange should have been used.

Worked example 3

Why can’t universal indicator be used in a titration?

Universal indicator is a mixture of several indicators designed to change colour continuously across the pH range from about 1 to 14. There’s no single sharp colour change, so there’s no clear end point. It’s ideal for estimating pH, and useless for titrations.

Worked example 4: a diprotic acid

Sodium carbonate is titrated with hydrochloric acid. There are two equivalence points, at about pH 8.3 and pH 3.8. Which indicators mark each?

• First equivalence (CO₃²⁻ → HCO₃⁻, pH 8.3): phenolphthalein (pink to colourless).
• Second equivalence (HCO₃⁻ → H₂CO₃, pH about 3.8): methyl orange (yellow to red, through orange).

Using both indicators in sequence in the same flask is a classic technique for analysing mixtures of carbonate and hydrogencarbonate, known as a double-indicator titration.

How much indicator to use

Two or three drops is standard. More is not better:

• Indicators are weak acids, so adding a lot uses up some titrant and introduces a small error.
• Strong colours are harder to judge. A faint, clean change is easier to spot than a deep one.

Colour-blindness and accessibility

Some colour changes (such as methyl orange’s red to yellow, with orange in between) can be hard to see for students with certain types of colour vision deficiency. Alternatives include:

• Choosing an indicator with a colourless-to-coloured change, such as phenolphthalein or thymolphthalein, where the chemistry allows it.
• Using a pH meter or a data logger and finding the equivalence point from the curve.
• Working in pairs so another person confirms the end point.

Making your own indicator

Red cabbage juice contains anthocyanin pigments that change through red, purple, blue, green and yellow as pH rises. It’s a great demonstration of how indicators work, but like universal indicator, its gradual changes make it poor for accurate titrations.

Exam shortcut

A quick rule of thumb if you are short of time in an exam: phenolphthalein when the base is strong, methyl orange when the acid is strong, and either when both are strong.

Common mistakes

• Always using phenolphthalein out of habit. It’s wrong for strong acid–weak base titrations.
• Choosing by the colours you like rather than by pH range.
• Using universal indicator for a titration.
• Adding too much indicator.
• Forgetting that the titration direction flips the colour change. Adding acid to alkali with phenolphthalein goes pink → colourless.

Key takeaways

• An indicator changes colour over roughly pKin ± 1.
• Choose one whose range lies inside the vertical section of the titration curve.
• Strong/strong: most indicators work. Weak acid/strong base: phenolphthalein. Strong acid/weak base: methyl orange.
• Weak/weak titrations have no sharp end point; use a pH meter.
• For the full practical method, read acid–base titration.