Theory of Indicators: How Acid-Base Indicators Work

Indicators are organic dyes (weak acids or weak bases) that change colour over a specific pH range. In acid-base titrations, they signal the endpoint — the moment when the amount of titrant exactly neutralises the analyte. The theory of indicators explains why these colour changes occur and how the dissociation equilibrium or structural rearrangement of the indicator molecule responds to pH variations.

Try the experiment yourself first. Run a real titration below — pick an acid-base pair and an indicator, add titrant drop by drop, and watch the flask colour and the pH curve respond. Then read on to understand the theory behind what you just saw.

What is an Indicator? Types & Principles

An indicator is typically a weak organic acid or base whose undissociated form has a different colour from its dissociated (ionised) form. Indicators are classified into two main categories:

• Weak acid indicators (e.g., phenolphthalein, bromothymol blue)
• Weak base indicators (e.g., methyl orange, methyl red)

Additionally, a universal indicator is a mixture of several indicators that produces a continuous colour change across a wide pH range (typically pH 4–10), allowing approximate pH determination. It works on the same principle: each component responds to a different pH interval.

Common Indicators and Their pH Ranges

Two Major Theories of Indicator Action

Two complementary theories explain the mechanism of colour change:

1. Ostwald’s Theory (Ionisation Theory)

Proposed by Friedrich Wilhelm Ostwald, this theory states that the colour change is due to the ionisation equilibrium of the indicator. Postulates:

• The indicator exists in two forms: undissociated (molecular) and dissociated (ionic).
• These two forms have distinctly different colours.
• In acidic solution, the equilibrium shifts toward the undissociated form; in basic solution, it shifts toward the dissociated form.
• For a weak acid indicator (HIn): HIn ⇌ H⁺ + In⁻. Colour of HIn ≠ colour of In⁻.
• For a weak base indicator (InOH): InOH ⇌ In⁺ + OH⁻. Again, two different colours.

Example – Phenolphthalein (weak acid indicator): Represented as HPh. In acidic medium, high [H⁺] suppresses dissociation (Le Chatelier’s principle), so colourless HPh dominates. In alkaline medium, H⁺ is removed, the equilibrium shifts right, producing pink Ph⁻ ions.

Example – Methyl orange (weak base indicator): Represented as MeOH. In acidic medium, OH⁻ ions are neutralised, shifting the equilibrium toward the red Me⁺ form. In alkaline medium, excess OH⁻ pushes the equilibrium left, giving the yellow unionised MeOH.

2. Quinonoid Theory (Resonance / Tautomeric Theory)

This theory explains colour changes through structural isomerism (tautomerism). Key postulates:

• Indicators can exist in two tautomeric forms: benzenoid (colourless or pale) and quinonoid (highly coloured).
• The two forms are in dynamic equilibrium.
• Acidic or alkaline conditions favour one form over the other, causing a visible colour shift.
• The quinonoid form has an extended conjugated double-bond system that absorbs visible light.

Phenolphthalein: In acidic or neutral medium, it adopts the colourless benzenoid (lactone) structure. In alkaline medium, the lactone ring opens and the molecule rearranges to the pink quinonoid form.

Methyl orange: In alkaline solution, the benzenoid (azo) form is stable and appears yellow. In acidic solution, protonation produces a quinonoid structure that appears red.

Comparison of Ostwald and Quinonoid Theories

Practical Applications & Selection of Indicators

Choosing the correct indicator is crucial for accurate titrations. The rule is simple: the indicator’s transition range must fall on the steep, near-vertical portion of the titration curve, so that the colour change (endpoint) coincides with the equivalence point.

• Strong acid + strong base (equivalence pH = 7): the steep region spans roughly pH 4–10, so both phenolphthalein and methyl orange work, though phenolphthalein gives a sharper endpoint.
• Weak acid + strong base (equivalence pH ≈ 8.7): only phenolphthalein is suitable; methyl orange changes colour far too early, in the buffer region.
• Strong acid + weak base (equivalence pH ≈ 5.3): methyl orange is the right choice; phenolphthalein would change too late.

Do not just memorise these rules — verify them yourself in the virtual lab at the top of this page. Try titrating acetic acid with NaOH using methyl orange and watch the colour change happen far away from the steep jump in the curve.

Summary of Key Equations & Concepts