QCE Chemistry - Unit 3 - Chemical equilibrium systems

Acid-Base Indicators | QCE Chemistry

Understand how acid-base indicators work and how to select an indicator for a titration in QCE Chemistry.

Updated 2026-05-18 - 4 min read

Acid-base indicators are weak acids or weak bases whose conjugate forms have different colours. They are useful because their equilibrium position changes as pH changes.

How indicators work

A common simplified indicator equilibrium is:

$ \mathrm{HIn(aq)} \rightleftharpoons \mathrm{H^+(aq)} + \mathrm{In^-(aq)} $

$\mathrm{HIn}$ and $\mathrm{In^-}$ have different colours.

In acidic solution, extra $\mathrm{H^+}$ shifts the equilibrium left, so the $\mathrm{HIn}$ colour dominates.

In basic solution, $\mathrm{H^+}$ is removed by reaction with $\mathrm{OH^-}$, so the equilibrium shifts right and the $\mathrm{In^-}$ colour dominates.

The reason $pK_a$ matters is that an indicator is still an equilibrium system. If the acidic and basic indicator forms are present in roughly equal amounts, the colour is halfway through its change. For a weak acid indicator:

$ K_a = \frac{[\mathrm{H^+}][\mathrm{In^-}]}{[\mathrm{HIn}]} $

At the halfway colour point, $[\mathrm{In^-}] = [\mathrm{HIn}]$, so the fraction becomes 1 and $K_a = [\mathrm{H^+}]$. Taking negative logs gives $pH = pK_a$.

That does not mean the colour suddenly flips at one pH. Human eyes see a gradual mixture of the two colours, so QCE treats the useful colour-change range as approximately one pH unit either side of $pK_a$.

Endpoint versus equivalence point

The equivalence point is the stoichiometric point: acid and base have reacted in the exact mole ratio from the balanced equation.

The endpoint is what you observe: the indicator's colour change.

A good titration uses an indicator whose endpoint is as close as practical to the equivalence point.

Choosing an indicator

Choose an indicator whose colour-change range falls within the steep part of the titration curve.

Typical patterns:

• strong acid + strong base: equivalence point near pH 7; many indicators can work because the vertical region is steep
• weak acid + strong base: equivalence point above pH 7; phenolphthalein is often suitable
• strong acid + weak base: equivalence point below pH 7; methyl orange is often more suitable
• weak acid + weak base: pH change near equivalence is usually too gradual for a sharp indicator endpoint

Here is the thinking process:

1. Locate or estimate the equivalence point of the titration.
2. Look at the steep vertical region near that equivalence point.
3. Choose an indicator whose $pK_a \pm 1$ range sits inside that steep region.
4. Reject indicators that change too early or too late, even if their colours are easy to see.

For strong acid-strong base titrations, the curve is so steep near pH 7 that several indicators can work. For weak acid-strong base titrations, the equivalence point is above pH 7, so a higher-range indicator is usually needed. For strong acid-weak base titrations, the equivalence point is below pH 7, so a lower-range indicator is usually needed.

Weak acid-weak base titrations are the awkward case. The pH change around equivalence is often too gradual, so an indicator may not give a sharp endpoint. In that situation, a pH probe or conductometric method is usually more reliable.

Worked example

Exam traps

Other traps:

• choosing an indicator because the colour is familiar
• assuming all equivalence points are pH 7
• saying indicators are strong acids or bases
• using too much indicator, which can slightly affect the titration
• choosing an indicator whose range touches the curve but misses the steepest part

Quick check

Sources

• QCAA Chemistry subject page
• QCAA Chemistry 2025 syllabus
• OpenStax Chemistry 2e: Acid-Base Titrations
• OpenStax Chemistry 2e: Relative Strengths of Acids and Bases