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An acid-base indicator contains a weak acid (HIn) that converts to its conjugate base (In−) as pH changes. Each form has a distinct color. When hydronium concentration is high, the equilibrium shifts left, favoring HIn and its color. When base is added, hydronium decreases, shifting equilibrium right toward In− and its color. The perceived color depends on the ratio of these two species.

The pH interval is the range of pH values over which an indicator visibly changes color, typically spanning approximately pKa ± 1. When pH is one unit below pKa, the HIn color dominates. When pH is one unit above pKa, the In− color dominates. Between these points, the solution displays a mixed color. This relationship is described by the henderson hasselbalch equation calculating buffers.

Strong acid-strong base titrations produce steep titration curves with rapid pH changes near the equivalence point. This steep region allows a broad range of indicators to work effectively. For example, during titration of 0.1 M hydrochloric acid with 0.1 M sodium hydroxide, pH jumps from 3 to 11, so both phenolphthalein (pH 8.3–10) and methyl red (pH 4.2–6) are suitable because their endpoints overlap the steep portion.

Weak acid-base titrations produce less steep titration curves than strong acid-strong base reactions, creating a narrower region of rapid pH change. This limits which indicators can be used effectively. During titration of 0.1 M acetic acid with 0.1 M sodium hydroxide, pH rises from 7 to 11. Phenolphthalein works because its endpoint overlaps the equivalence point, but methyl red fails because its endpoint misses this critical region.

The endpoint is the point at which an indicator visibly changes color during titration. The ideal indicator has an endpoint positioned near the equivalence point, where the analyte is completely neutralized. For accurate titration results, the indicator's pH interval must overlap the steep portion of the titration curve. Misalignment between endpoint and equivalence point leads to systematic errors in volume measurements.

Universal indicators and pH paper contain a mixture of multiple indicators, each with different pH intervals and colors. This combination allows them to display distinct colors across a wide range of pH values, enabling approximate pH determination of unknown solutions through visual comparison. Single-color indicators like phenolphthalein or methyl red change color over a narrow, specific pH range and are better suited for precise titration endpoints.

An indicator's equilibrium between HIn and In− responds to changes in hydronium concentration according to Le Châtelier's principle. Adding acid increases hydronium, shifting equilibrium left toward HIn and its associated color. Adding base decreases hydronium, shifting equilibrium right toward In− and its color. This reversible color change makes indicators useful for monitoring pH changes in titrations and other acid-base reactions.