Acid-Base Indicators

An acid-base indicator is a weak organic acid or base whose conjugate form has a distinctly different colour from the original form. Because the colour change depends on the relative concentrations of the two forms, indicators respond visibly to changes in the pH of a solution. Common examples include litmus, phenolphthalein, and methyl orange — each changing colour across a specific, narrow pH range called its transition interval. The theoretical foundation for indicator behaviour lies in ionic equilibrium. An indicator molecule HIn dissociates partially in water to release a proton and form its conjugate base In⁻. The ratio [In⁻]/[HIn] is governed by the acid dissociation constant KIn. When pH ≪ pKIn the indicator exists predominantly in the HIn form (colour A); when pH ≫ pKIn it exists mainly as In⁻ (colour B). The eye detects an apparent colour change roughly one pH unit on either side of pKIn, giving a transition interval of about two pH units. Selecting the correct indicator for a titration is essential for accurate results. The endpoint (the moment the indicator changes colour) must coincide as closely as possible with the equivalence point (the stoichiometric completion of the neutralisation reaction). Because the pH at the equivalence point differs depending on whether the titration involves a strong or weak acid and base, different indicators are appropriate for different titration types. Choosing the wrong indicator can introduce a systematic error of several millilitres in a burette reading.

Key takeaways

The Indicator Equilibrium
Structural Basis of Colour Change
Transition Intervals of Common Indicators
Matching Indicator to Titration Type
Minimising Indicator Error

The Indicator Equilibrium

Every indicator participates in the equilibrium HIn(aq) ⇌ H⁺(aq) + In⁻(aq). The acid dissociation constant is KIn = [H⁺][In⁻]/[HIn]. Taking logarithms gives pH = pKIn + log([In⁻]/[HIn]). When [In⁻] = [HIn], pH = pKIn exactly, and both colours contribute equally — this is the midpoint of the transition interval. The human eye detects the colour shift when the ratio [In⁻]/[HIn] lies roughly between 1/10 and 10/1, corresponding to a pH range of pKIn ± 1.

Structural Basis of Colour Change

Indicators are typically large conjugated organic molecules. The protonation state alters the extent of the π-conjugated system, shifting the wavelength of maximum absorption (λmax) and hence the observed colour. Phenolphthalein is colourless in its fully protonated lactone form (ring closed, no extended conjugation) and becomes deep magenta-pink in its dianion form (ring open, fully conjugated quinonoid structure) above pH ≈ 8.2.

Transition Intervals of Common Indicators

Methyl orange: pKIn ≈ 3.5, transition interval pH 3.1–4.4 (red → orange → yellow). Methyl red: pKIn ≈ 5.1, transition interval pH 4.4–6.2 (red → yellow). Bromothymol blue: pKIn ≈ 7.0, transition interval pH 6.0–7.6 (yellow → blue). Phenolphthalein: pKIn ≈ 9.1, transition interval pH 8.2–10.0 (colourless → pink). Universal indicator is a mixture of several indicators that shows a continuous rainbow of colours across pH 1–14.

Matching Indicator to Titration Type

Strong acid–strong base titrations (e.g., HCl vs. NaOH): the equivalence point is pH 7.0, and the pH curve is nearly vertical from pH 3 to pH 11, so methyl orange, methyl red, bromothymol blue, or phenolphthalein all give accurate endpoints. Weak acid–strong base titrations (e.g., CH₃COOH vs. NaOH): equivalence point is alkaline (pH ≈ 8.7 for 0.1 mol dm⁻³ solutions) because the salt CH₃COO⁻ undergoes hydrolysis; phenolphthalein (pKIn ≈ 9.1) is the correct choice. Strong acid–weak base titrations (e.g., HCl vs. NH₃): equivalence point is acidic (pH ≈ 5.3); methyl orange or methyl red is appropriate. Weak acid–weak base titrations have no sharp vertical section on the pH curve; no indicator gives a sharp colour change, making them impractical to titrate using an indicator.

Minimising Indicator Error

Indicators are themselves weak acids or bases that consume a tiny amount of the titrant. To minimise this error, use only 2–3 drops of indicator solution. The volume of indicator solution should be negligible compared with the titration volume. For very precise work, a pH meter is used instead of an indicator, as it introduces no chemical error and can pinpoint the equivalence point exactly from the inflection of the pH vs. volume curve.

The Key Equation

\text{pH} = \text{p}K_{\text{In}} + \log\!\left(\frac{[\text{In}^-]}{[\text{HIn}]}\right)

\text{pH} — pH of solution (dimensionless): Negative logarithm (base 10) of the hydrogen ion activity; measures acidity or alkalinity of the solution.
\text{p}K_{\text{In}} — Indicator pKa (dimensionless): Negative logarithm of the indicator's acid dissociation constant KIn; equals the pH at which both indicator forms are present in equal concentrations, i.e., the colour midpoint.
[\text{In}^-] — Conjugate base concentration (mol dm⁻³): Molar concentration of the deprotonated (basic) form of the indicator, which displays colour B.
[\text{HIn}] — Acid form concentration (mol dm⁻³): Molar concentration of the protonated (acidic) form of the indicator, which displays colour A.

Worked Example

Problem. A student titrates 25.00 cm³ of 0.100 mol dm⁻³ ethanoic acid (Ka = 1.8 × 10⁻⁵ mol dm⁻³) with 0.100 mol dm⁻³ sodium hydroxide. Calculate the pH at the equivalence point and select an appropriate indicator from: methyl orange (transition pH 3.1–4.4), bromothymol blue (pH 6.0–7.6), and phenolphthalein (pH 8.2–10.0).

\text{pH}_{\text{eq}} \approx 8.72; \quad \text{Indicator: phenolphthalein (pH 8.2–10.0)}

Answer. pH at equivalence point ≈ 8.72; phenolphthalein is the appropriate indicator.

Figure

Diagram

Origins · 1 minFriedrich Wilhelm Ostwald and how this was discovered

Friedrich Wilhelm Ostwald

1853–1932 · Baltic German (Latvian-born)

Ostwald developed the quantitative theory of acid-base indicators in the 1890s, explaining indicator colour changes as a direct consequence of the equilibrium between the protonated and deprotonated forms of a weak organic acid. He derived the relationship between indicator colour ratio and [H⁺], which is the conceptual forerunner of the Henderson–Hasselbalch expression applied to indicators. His broader work on ionic equilibria, electrolytic dissociation, and the dilution law (Ostwald's dilution law) earned him the Nobel Prize in Chemistry in 1909. Ostwald was also a passionate philosopher of science and an enthusiastic painter — he even developed his own colour theory and classification system for pigments, publishing a book on the science of colour in 1917.

Acid Dissociation Constant (Ka) — KIn is itself a Ka value; understanding Ka and pKa is essential for predicting the pH at which an indicator changes colour.
Henderson–Hasselbalch Equation — The indicator transition equation pH = pKIn + log([In⁻]/[HIn]) is a direct application of the Henderson–Hasselbalch equation to the indicator equilibrium.
Acid-Base Titrations — Indicators are the primary tool for detecting the endpoint of acid-base titrations; choosing the right indicator requires knowing the pH at the equivalence point.
Buffer Solutions — Buffers resist pH change near pKa, the same region where indicators change colour — understanding both concepts together clarifies why the indicator transition interval spans ≈ 2 pH units.
Salt Hydrolysis — The alkaline equivalence point in a weak acid–strong base titration results from hydrolysis of the conjugate base salt, which determines which indicator is appropriate.