Titration Principles: Choosing the Right Indicator
The indicator must change colour within the steep part of the curve, and where that steep part lies depends on the strengths involved.
Equivalence point and end point
The equivalence point is where stoichiometrically equivalent amounts have reacted. The end point is where the indicator changes colour. They are not the same thing, and the difference between them is the indicator error.
A well-chosen indicator makes that error negligible; a poorly chosen one makes the titration meaningless. Choosing correctly is therefore the substance of the topic.
Where the equivalence point lies
| Titration | pH at equivalence | Why |
|---|---|---|
| Strong acid with strong base | Neutral | The salt formed does not hydrolyse |
| Weak acid with strong base | Basic | The conjugate base of the weak acid hydrolyses |
| Strong acid with weak base | Acidic | The conjugate acid of the weak base hydrolyses |
| Weak acid with weak base | Depends on relative strengths | Both ions hydrolyse |
Choosing the indicator
An acid–base indicator is itself a weak acid whose conjugate forms differ in colour. It changes over a range of roughly two pH units centred on its own dissociation constant.
The rule is simply that this range must fall within the steep portion of the titration curve.
| Titration | Suitable indicator range |
|---|---|
| Strong acid with strong base | Wide steep region — most indicators work |
| Weak acid with strong base | Basic range only |
| Strong acid with weak base | Acidic range only |
Using an indicator that changes colour in acid for a weak acid titration gives an end point well before the equivalence point, and the resulting concentration is systematically wrong. That is the standard error the question is designed to expose.
The buffer region
In a weak acid titration, the region before the equivalence point is a buffer, and the pH there changes only slowly. At half neutralisation the concentrations of acid and conjugate base are equal, so the pH equals the acid's pK.
That half-equivalence point is therefore a direct route to the dissociation constant, and reading it from a titration curve is a common exercise.
Other titration types
| Type | Reaction | End point detected by |
|---|---|---|
| Redox | Electron transfer | Self-indicating reagent, or a redox indicator |
| Complexometric | Complex formation | Metal ion indicator changing colour on release |
| Precipitation | Insoluble product formed | Indicator forming a coloured precipitate after the equivalence point |
Some redox reagents are self-indicating, since the reagent itself is coloured and its persistence marks the end point. That removes the need for a separate indicator, which is one reason such reagents are popular.
Complexometric titrations require pH control, because the chelating agent's effectiveness depends on how much of it is deprotonated. A buffer is therefore part of the procedure rather than an optional extra, and explaining why is a standard question.
Frequently asked questions
Why is the equivalence point of a weak acid titration basic?
Because the salt formed contains the conjugate base of the weak acid, which hydrolyses to produce hydroxide.
Why can most indicators be used for a strong acid with a strong base?
Because the curve is extremely steep at the equivalence point, spanning many pH units for a very small volume of titrant, so almost any indicator changes within that region.
What does the half-equivalence point give?
The pK of the weak acid, since acid and conjugate base are present in equal concentration there.
Why is a buffer needed in a complexometric titration?
Because the chelating agent competes with hydrogen ions, so its ability to bind the metal depends on pH. Holding the pH constant makes the reaction stoichiometric and reproducible.
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