Explain how acid-base indicators work as weak acid equilibria and select an appropriate indicator for a titration

What this dot point is asking

An acid-base indicator is a substance that changes colour depending on the pH of the solution. Most are weak organic acids (or bases) where the protonated and deprotonated forms absorb different wavelengths of light.

How indicators work

Represent the indicator as a weak acid $\text{HIn}$:

where $\text{HIn}$ is one colour and $\text{In}^-$ is another. In acidic solution the high $[\text{H}^+]$ pushes the equilibrium to the left (Le Chatelier), so the $\text{HIn}$ colour dominates. In basic solution $[\text{H}^+]$ is low, the equilibrium shifts right, and the $\text{In}^-$ colour dominates. For example, with litmus the $\text{HIn}$ form is red and the $\text{In}^-$ form is blue.

The colour change range

Because the colour switches over the range where neither form completely dominates, each indicator has a characteristic transition interval. Common examples:

• Methyl orange: red to yellow, range about pH 3.1 to 4.4.
• Bromothymol blue: yellow to blue, range about pH 6.0 to 7.6.
• Phenolphthalein: colourless to pink, range about pH 8.3 to 10.0.

Choosing an indicator for a titration

The key skill is matching the indicator to the equivalence point of the titration. You want the indicator to change colour exactly where the titration curve is steepest, so the colour change is sharp and corresponds to the equivalence point.

• Strong acid with strong base: equivalence point at pH 7, with a very steep curve from about pH 3 to 11. Almost any of the common indicators works.
• Strong acid with weak base: equivalence point is acidic (below 7). Use methyl orange.
• Weak acid with strong base: equivalence point is basic (above 7). Use phenolphthalein.

End point versus equivalence point

Two terms are easy to confuse. The equivalence point is the point in a titration where the amount of titrant added is stoichiometrically equal to the amount of analyte present, a property of the chemistry alone. The end point is where the indicator actually changes colour, a property of the indicator you chose. They coincide only when the indicator's transition range straddles the equivalence-point pH. Choosing well makes the difference between them negligible, which is why indicator selection is a marked skill in WACE volumetric analysis.

Universal indicator and why titrations avoid it

Universal indicator is a mixture of several indicators that gives a continuous colour change across the whole pH scale, useful for estimating pH but useless for a titration. Because it changes gradually it gives no sharp end point, so it cannot pin down the equivalence volume to within a drop. A single indicator with a narrow transition range, matched to the equivalence-point pH, changes colour over the addition of roughly one drop on the steep part of the curve, giving the precise end point a titration needs.

Why this matters

Choosing the right indicator is essential for accurate volumetric analysis, the quantitative technique used to find unknown concentrations. A mismatched indicator gives a false end point and a wrong result, so this dot point links directly to the titration calculations elsewhere in Unit 3.