Acid-Base Titrations and Indicators Explained Simply

An acid–base titration finds the unknown concentration of an acid or base by reacting it with a solution of known concentration, adding the titrant drop by drop until the reaction is just complete. The point where stoichiometrically equivalent amounts have reacted is the equivalence point, while the point where the indicator changes colour is the end point — and a good titration keeps the two as close as possible.

Titrations are the most practised quantitative experiment in school chemistry because they turn a colour change into a precise number. Choosing the right indicator for each acid–base combination is the conceptual heart of the topic and a guaranteed exam question.

Key Concepts

• Equivalence point: moles of acid = moles of base (adjusted for stoichiometry); detected by a sharp pH jump.
• End point: the indicator's colour change; ideally coincides with the equivalence point.
• Indicator ranges: phenolphthalein 8.3–10.0 (colourless → pink); methyl orange 3.1–4.4 (red → yellow); litmus ~5–8.
• Correct pairings: strong acid–strong base: phenolphthalein or methyl orange; weak acid–strong base: phenolphthalein (equivalence pH > 7); strong acid–weak base: methyl orange (equivalence pH < 7); weak–weak: no sharp jump, no suitable indicator.

An indicator is a weak acid/base (HIn ⇌ H⁺ + In⁻) whose two forms differ in colour, switching near its pKa. In a strong–strong titration the pH jumps from ~4 to ~10 in a drop, so either indicator works. Acetic acid vs NaOH ends above pH 7 (salt hydrolysis) → phenolphthalein only; ammonia vs HCl ends below pH 7 → methyl orange only.

Formulas

• Titration equation: M₁V₁ = M₂V₂ for 1:1 reactions (e.g. HCl vs NaOH), where M is molarity and V is volume in consistent units.
• General form: (M₁V₁)/n₁ = (M₂V₂)/n₂, where n₁, n₂ are the stoichiometric coefficients — e.g. n = 2 for H₂SO₄ vs NaOH.
• Indicator rule: choose an indicator whose colour-change range brackets the equivalence-point pH of that titration.

Worked Example

25.0 mL of 0.100 M NaOH exactly neutralises 20.0 mL of dilute HCl using phenolphthalein. Find the molarity of the HCl. Step 1: the reaction HCl + NaOH → NaCl + H₂O has 1:1 stoichiometry. Step 2: moles of NaOH = 0.100 × 0.0250 = 0.00250 mol. Step 3: moles of HCl = 0.00250 mol (1:1). Step 4: molarity of HCl = 0.00250 ÷ 0.0200 = 0.125 M. So the hydrochloric acid is 0.125 M.

Common Misconceptions

• "The indicator tells you the pH of the solution." No — it only signals that the pH has crossed its narrow change range; it is a switch, not a meter.
• "Phenolphthalein works for every titration." Wrong — in a strong acid vs weak base titration the equivalence pH is below 7, where phenolphthalein stays colourless, so methyl orange is required.
• "End point and equivalence point are the same thing." They are close in a good titration but conceptually different — one is observed (colour), the other is stoichiometric (moles).

Real-World Applications

• Industry: titrations check the acidity of vinegar, the saponification value of oils, and the purity of pharmaceutical raw materials.
• Environment: water-testing labs titrate to find hardness, dissolved CO₂ and residual chlorine in drinking water.
• Medicine: antacid strength and vitamin C content in tablets are assayed by titration in quality-control labs.

Titration converts a colour change into a concentration through the M₁V₁ = M₂V₂ relationship, and the indicator must match the titration's equivalence pH — phenolphthalein for weak acid–strong base, methyl orange for strong acid–weak base, either for strong–strong. Learn the four pairings and the stoichiometry, and titration numericals become free marks.