A titration is just a controlled neutralization: drip a base of known concentration into an acid (or vice versa) until the reaction is complete, then back out the unknown concentration from how much you used. The reason the lab feels tricky is the curve — the pH shoots up at the equivalence point, the indicator flips, and you have to know which feature to read. This guide on acids, bases, and titration sorts strong from weak, defines the equivalence point precisely, and walks through one titration curve end to end.
Strong vs. Weak Acids and Bases
Acid–base strength is about how completely a species ionizes in water, not how dangerous it feels.
A strong acid ionizes essentially 100% in water — HCl, HNO₃, HClO₄, H₂SO₄ (first proton), HBr, and HI. For 0.10 M HCl, you can write [H⁺] ≈ 0.10 M, no equilibrium calculation needed. A strong base does the same — NaOH, KOH, and the heavier Group 1 hydroxides ionize fully into OH⁻.
A weak acid ionizes only partially. Acetic acid (CH₃COOH, the acid in vinegar) is the canonical example: in a 0.10 M solution, only a small fraction donates its proton, and the rest sits as undissociated molecules in equilibrium with the ions. The fraction that ionizes is set by K(a), the acid dissociation constant. The same logic gives weak bases a K(b). Ammonia (NH₃) is the standard weak base.
That distinction matters in titration because the shape of the curve depends on it. A strong-acid–strong-base titration produces a steep, symmetric jump centered at pH 7. A weak-acid–strong-base titration produces an asymmetric curve with a buffer plateau and an equivalence point above pH 7.
The Equivalence Point Is Not the Endpoint
Two terms get used interchangeably and should not be.
The equivalence point is the point in the titration where the moles of added titrant exactly equal the moles of analyte being neutralized. It is a chemical condition: stoichiometrically complete reaction.
The endpoint is the point you actually observe — usually when an indicator changes color. A well-chosen indicator changes color so close to the equivalence point that the difference is negligible. A badly chosen one (one whose pK(a) is far from the equivalence-point pH) flips early or late, and you record a wrong concentration.
For a strong-acid–strong-base titration, the equivalence point is at pH 7, and any indicator whose color change overlaps the steep rise — phenolphthalein, bromothymol blue — works. For a weak-acid–strong-base titration (e.g. acetic acid with NaOH), the equivalence point is above 7 because the conjugate base of the weak acid (acetate) makes the solution basic, so phenolphthalein (changes around pH 8.3) is the right choice. Match the indicator's pK(In) to the equivalence-point pH within about one unit.
Reading a Titration Curve: Four Regions
A weak-acid–strong-base titration curve has four distinct regions, and each one calls for a different calculation.
- Initial point (no titrant added). Only the weak acid is present. pH is set by the K(a) of the acid and its initial concentration: solve an ICE table or use the shortcut pH ≈ ½ (pK(a) − log C). For acetic acid (pK(a) = 4.74) at 0.10 M, pH ≈ 2.87.
- Buffer region (partway to equivalence). Both the weak acid and its conjugate base are present in meaningful amounts, so the Henderson–Hasselbalch equation works: pH = pK(a) + log ([A⁻]/[HA]). The midpoint of this region, when exactly half the acid has been neutralized, is the most useful point on the whole curve: pH = pK(a). That is the cleanest experimental way to find pK(a) — read the pH at half-equivalence.
- Equivalence point. All the weak acid has been converted to its conjugate base. The pH is determined by the conjugate base hydrolyzing water — slightly basic. Compute it from K(b) = K(w)/K(a) and the diluted concentration of conjugate base.
- Past equivalence. Excess strong base controls the pH. Treat it as a simple strong-base problem: pOH = −log(excess OH⁻ concentration), then pH = 14 − pOH.
A strong-acid–strong-base curve drops two of those regions — there is no buffer plateau and no hydrolysis at the equivalence point — but the past-equivalence region is the same.
Worked Titration: 25.00 mL of 0.100 M HCl with 0.100 M NaOH
The reaction: HCl + NaOH → NaCl + H₂O. Both species are strong, so neutralization is complete.
Moles of HCl initially: 0.02500 L × 0.100 M = 0.00250 mol.
To reach the equivalence point, you need 0.00250 mol of NaOH. With 0.100 M NaOH, that is 0.02500 L = 25.00 mL of titrant. Equal volumes, because concentrations are equal.
At equivalence, the only ions present are Na⁺ and Cl⁻ (a neutral salt) in 50.00 mL of solution. The water itself sets the pH at 7.00. That is exactly where the steep mid-rise of the curve crosses.
If the unknown were the acid concentration instead, you would reverse the algebra. Suppose 24.85 mL of 0.100 M NaOH reached the equivalence point on a 25.00 mL acid sample. Moles of acid = 0.02485 × 0.100 = 0.002485 mol; concentration = 0.002485 / 0.02500 = 0.0994 M. That is the structure of every titration calculation: moles of titrant from volume × molarity, then back out the analyte using the reaction stoichiometry.
Getting Help
Titrations sit on top of equilibrium math. If pH and pOH still feels shaky, that is the prerequisite. The buffer plateau in the middle of a weak-acid titration is the same chemistry as a deliberately made buffer solution, and the Henderson–Hasselbalch equation is the same in both places.
Conclusion
A titration is a controlled reaction with a numerical payoff. Tell strong from weak first — the curve shape and the equivalence-point pH depend on it. The equivalence point is set by stoichiometry; the endpoint is set by the indicator, and a good indicator is one whose color change brackets the equivalence-point pH. Each region of the curve has its own calculation, but they all begin the same way: moles of titrant added × concentration, then compare to moles of analyte.
