Acid–Base Indicators and Their Selection

Introduction

The colour change that ends a titration is itself a chemical equilibrium. An indicator is usually a weak acid or base whose conjugate forms have different colours. It does not know when analyte and titrant moles are equal; it responds to pH. Good selection places its visible transition within the sharp pH change around equivalence, making the observed endpoint a useful approximation.

Core explanation

Represent a simple indicator as HIn ⇌ H⁺ + In⁻. The acid form HIn and base form In⁻ have different colours. Its dissociation constant gives pH = pKₐ(In) + log([In⁻]/[HIn]) under the usual dilute-solution approximation. Near pH = pKₐ(In), the forms have comparable concentrations. A visible transition generally spans a range around the indicator pKₐ because the eye needs a substantial change in their ratio to perceive a new colour. The exact interval depends on indicator, solvent, concentration and viewing conditions.

Suppose a titration curve jumps from pH 4 to pH 10 over a tiny volume around equivalence. Many indicators with transitions within that interval could work. For a weak acid titrated by strong base, equivalence is usually basic because the conjugate base hydrolyses; an indicator transitioning in the basic region may be suitable. For a weak base titrated by strong acid, equivalence is acidic, so the transition should be in an acidic part of the steep region. A strong acid–strong base titration often has a broad steep span around pH 7 at common concentrations, allowing more choices. Very dilute titrations may have a less sharp jump and demand more care.

An indicator should be added in only a small amount. It is itself an acid or base and can consume some titrant if used excessively. The chosen endpoint convention matters: the first persistent faint colour may differ from a deep colour, and consistent lighting and swirling improve repeatability. A blank titration or endpoint correction can sometimes quantify an indicator contribution.

An instrumental pH electrode can measure the curve without relying on human colour judgment, which helps for coloured or opaque samples. Yet an instrument still needs calibration and a rule for locating equivalence. Neither visual nor instrumental endpoint is inherently perfect; both infer a stoichiometric point from a signal.

Step-by-step reasoning

1. Identify acid and base strengths and predict whether equivalence is acidic, neutral or basic. 2. Sketch or inspect the pH-versus-volume curve near equivalence. 3. Choose an indicator whose transition interval lies largely within the steep region. 4. Use a small indicator amount and a consistent colour-change convention. 5. Compare endpoint with equivalence and consider correction if required accuracy is high.

Visual explanation

Draw a weak-acid/strong-base pH curve with an equivalence point above pH 7. Shade one indicator transition interval across the steep vertical segment and another far below it. The first interval gives a small volume error; the second changes colour too early. Add a miniature HIn ⇌ In⁻ colour key beside the curve.

Real-world analogy

An indicator is a warning lamp that turns on when the pH enters its chosen range. It is useful if that range coincides with the narrow volume interval near chemical equivalence. A lamp set to switch on much earlier or later is functioning correctly as a pH sensor but is the wrong sensor for this titration.

Real-world example

In acetic-acid titration with sodium hydroxide, the equivalence solution contains acetate, which makes it basic. An indicator with a basic transition can mark the sharp part of the curve. An indicator that changes far into the acidic buffer region would consistently stop titration before enough NaOH had reacted.

Why?

Why is the colour change spread over a pH interval rather than occurring at one pH? Both indicator forms coexist in equilibrium. As pH changes, their concentration ratio changes continuously. The eye perceives one colour as dominant only when the ratio crosses a practical threshold, so visible transition occupies a range.

Common misconception

“Every acid–base titration reaches equivalence at pH 7” is false for weak-acid or weak-base systems. Another mistake is selecting an indicator solely because it has a dramatic colour; the transition pH relative to the titration curve is what determines endpoint error.

Worked example

Suppose an indicator has pKₐ = 8.5. At pH 8.5, [In⁻]/[HIn] = 10^(8.5−8.5) = 1. At pH 9.5, the ratio is 10. If a weak-acid/strong-base titration has its steep region from roughly pH 7 to 10, this indicator can be useful. If equivalence occurs in an acidic steep region from pH 3 to 5, its colour change would occur too late.

Quick check

1. Why might a basic-range indicator suit weak-acid/strong-base titration? Answer: The equivalence solution contains the weak acid's conjugate base, so its pH is usually above 7. A basic-range transition can lie in the steep region near that equivalence point.

Exam focus

Write the indicator equilibrium and relate the form ratio to pH. Select from the curve's steep region, not from a memorised “neutral point.” State that endpoint and equivalence are distinct. Explain how excessive indicator can introduce a small but systematic titrant demand.

Advanced insight

For coloured, turbid or strongly buffered samples, an indicator's visual transition may be difficult to detect even if its pH interval is theoretically suitable. Potentiometric detection can give a more reproducible signal, but electrode response and junction potentials then become new uncertainty sources. Method choice shifts, rather than eliminates, the measurement problem.

Summary

An acid–base indicator changes colour as its protonated and deprotonated forms shift in equilibrium. Its transition spans a pH range. Select that range to fall within the steep titration-curve region near stoichiometric equivalence, and use a clear endpoint convention to limit indicator error.

Practice questions

1. An indicator has pKₐ 5.0. What is [In⁻]/[HIn] at pH 6.0? Answer: Henderson–Hasselbalch gives log ratio = 6.0 − 5.0 = 1.0, so [In⁻]/[HIn] = 10.

2. Why is a low-pH indicator generally poor for a weak-acid/strong-base titration with basic equivalence? Answer: It may change colour in the buffer region before the sharp rise near equivalence, causing an early endpoint and a low titrant volume.

3. Does a pH electrode remove every endpoint uncertainty? Answer: No. It replaces colour judgment with electrode calibration, response, sampling and curve-analysis issues. The stoichiometric equivalence still has to be inferred from measured data.