Titrations & Indicators

Add base one drop at a time, watch the pH, and read the story the curve tells — including the trap of assuming 'neutralised' means pH 7.

High schoolIntro Gen ChemUni Year 1
⏱️ About 22 min

A titration looks calm — a burette dripping base into acid — but the pH is doing something dramatic. For most of the run it barely budges, then near one special point it leaps several units in a single drop. Learn to read that leap and you can measure an unknown concentration exactly.

💡
The big idea: A titration curve plots pH against the volume of titrant added. The steep jump marks the equivalence point — where acid and base are in stoichiometric balance. Crucially, that point is only pH 7 for a strong acid–strong base; weak partners shift it away from 7.
🎯 By the end, you'll be able to
  • Read a titration curve and locate the equivalence point
  • Distinguish the equivalence point (stoichiometric) from the endpoint (indicator colour change)
  • Explain why the equivalence pH is 7 only for strong acid–strong base
  • Choose an indicator whose colour change matches the equivalence pH
📎 Helpful to know first

What a titration measures

In a titration you slowly add a solution of known concentration (the titrant, usually in a burette) to a measured volume of the unknown, tracking pH as you go. The goal is the equivalence point: the moment when you've added exactly enough titrant to react with all the unknown, mole for mole.

Plot pH versus volume added and you get the classic S-shaped titration curve — flat, then a near-vertical cliff at equivalence, then flat again.

\[ \ce{HCl + NaOH -> NaCl + H2O} \]
A strong acid neutralised by a strong base. At equivalence the moles of acid and base are equal.
🔑 Equivalence point vs endpoint
The equivalence point is the true stoichiometric balance (equal moles of acid and base). The endpoint is where your indicator changes colour. A well-chosen indicator makes the two practically coincide — but they are different ideas, and a badly matched indicator makes them disagree.
⚠️ Neutralised ≠ pH 7
The biggest misconception: reaching the equivalence point does not always mean pH 7. It's 7 only when a strong acid meets a strong base, giving a neutral salt. A weak acid–strong base titration finishes at the equivalence point with pH above 7 (the conjugate base left behind is basic); a strong acid–weak base finishes below 7.

Why the weak-acid curve ends up basic

Titrate acetic acid (weak) with NaOH (strong). At the equivalence point every acetic acid molecule has been converted to acetate, A⁻. But acetate is a conjugate base — it reacts a little with water to make OH⁻. So the solution at equivalence is a solution of a weak base, and its pH lands above 7 (typically around 8–9).

Notice too that the first half of this curve is a buffer region: you have both HA and A⁻ present, so the pH climbs only gently — exactly the Henderson–Hasselbalch behaviour from the last lesson. At the half-equivalence point, pH = pKa.

✨ Half-equivalence: a free pKa
At the half-equivalence point of a weak acid titration, exactly half the acid has been converted, so [HA] = [A⁻]. Henderson–Hasselbalch then gives pH = pKa. Reading the pH at half-equivalence off the curve is a quick way to measure an unknown acid's pKa.

Choosing an indicator

An indicator is itself a weak acid whose acid and base forms are different colours; it flips colour over a narrow pH range. The rule is simple: pick an indicator whose colour-change range straddles the equivalence pH of your titration.

  • Strong acid–strong base (equivalence ≈ 7): bromothymol blue works.
  • Weak acid–strong base (equivalence > 7): phenolphthalein (changes around pH 8–10) is ideal.
  • Strong acid–weak base (equivalence < 7): methyl orange (changes around pH 3–4) fits.
📝 Worked example: You titrate 25.0 mL of 0.100 M HCl with 0.100 M NaOH. What volume of NaOH reaches the equivalence point, and what is the pH there?
  1. Moles of HCl = 0.0250 L × 0.100 mol/L = 2.50×10⁻³ mol.
  2. Equivalence needs equal moles of NaOH: 2.50×10⁻³ mol. Volume = mol ÷ concentration = 2.50×10⁻³ ÷ 0.100 = 0.0250 L = 25.0 mL.
  3. Both are strong, so the salt (NaCl) is neutral → pH = 7.00 at equivalence.
✓ 25.0 mL of NaOH; pH = 7.00 (strong–strong, so neutral).
✏️ Practice: You titrate 20.0 mL of 0.150 M HCl with 0.100 M NaOH. What volume of NaOH (in mL) is needed to reach the equivalence point?
mL
Solution
  1. Moles of HCl = 0.0200 L × 0.150 mol/L = 3.00×10⁻³ mol.
  2. Equivalence needs the same moles of NaOH: 3.00×10⁻³ mol.
  3. Volume = 3.00×10⁻³ mol ÷ 0.100 mol/L = 0.0300 L = 30.0 mL.

Check your understanding

1. You reach the equivalence point of a weak acid–strong base titration. The pH is most likely…
At equivalence the weak acid has become its conjugate base, which reacts with water to give OH⁻. So the solution is basic and pH > 7. Only strong–strong gives pH 7.
2. What is the difference between the equivalence point and the endpoint?
The equivalence point is where moles of acid and base match. The endpoint is where the indicator changes colour. A good indicator makes them nearly coincide, but they're distinct.
3. At the half-equivalence point of a weak acid titration, the pH equals…
At half-equivalence [HA] = [A⁻], so Henderson–Hasselbalch gives pH = pKa. That's why the curve's half-equivalence pH is a handy way to read off pKa.
🔎 Related lessons
Previous: Buffers & Henderson–Hasselbalch Next: Solubility Equilibria (Ksp) Strong vs Weak Acids & Bases
✅ Key takeaways
  • A titration curve plots pH against titrant volume; equivalence is the steep jump.
  • Equivalence point = equal moles of acid and base; endpoint = indicator colour change.
  • Equivalence pH is 7 only for strong acid–strong base.
  • Weak acid–strong base finishes above 7; strong acid–weak base finishes below 7.
  • At half-equivalence pH = pKa; choose an indicator whose range spans the equivalence pH.
➡️ Titrations exploit reactions that go essentially to completion. But some ionic compounds barely dissolve at all, sitting in equilibrium with a trace of dissolved ions. Quantifying that with Ksp is the final stop.
Want to test yourself on this? Try the Chemistry practice test →