Acid-Base Titration: How to Do the Calculation, Pick the Right Indicator, and Read the Curve

An acid-base titration determines the concentration of an unknown acid or base by slowly adding a solution of known concentration (the titrant) until the reaction reaches the equivalence point — where moles of acid exactly equal moles of base. At equivalence, M_acid x V_acid = M_base x V_base for monoprotic acids and bases. The indicator you choose must change color at a pH near the equivalence point, which depends on whether you are titrating a strong acid with a strong base (equivalence at pH 7), a weak acid with a strong base (equivalence above pH 7), or a weak base with a strong acid (equivalence below pH 7).

For a simple monoprotic acid-base titration, the math at the equivalence point is straightforward. Since every mole of acid reacts with one mole of base, the moles must be equal: M_acid x V_acid = M_base x V_base. If you know three of these four values, you can solve for the fourth.

Example: You titrate 25.00 mL of an unknown HCl solution with 0.100 M NaOH. It takes 31.25 mL of NaOH to reach the equivalence point. What is the HCl concentration? M_HCl = (0.100 M x 31.25 mL) / 25.00 mL = 0.125 M.

For polyprotic acids (like H2SO4 or H3PO4), you need to account for the stoichiometry. Sulfuric acid donates two protons per molecule, so the relationship becomes 2 x M_acid x V_acid = M_base x V_base. A 0.050 M H2SO4 solution behaves like 0.100 M in terms of available protons. Forgetting this factor of 2 is one of the most common titration mistakes on exams. Always write the balanced equation first and identify the mole ratio before plugging numbers into a formula.

A titration curve plots pH (y-axis) against volume of titrant added (x-axis). The shape of this curve tells you everything about the reaction chemistry.

For a strong acid titrated with a strong base, the curve starts at a low pH (around 1-2 for typical concentrations), stays relatively flat as base is added (the buffer zone does not exist for strong acid-strong base combinations, so pH rises slowly), then shoots nearly vertical at the equivalence point — rising from about pH 4 to pH 10 in a single drop of titrant. This steep vertical section is why the equivalence point is easy to detect. After equivalence, pH rises gradually as excess base accumulates.

For a weak acid titrated with a strong base, the curve looks different in two key ways. First, the initial pH is higher (a 0.1 M acetic acid solution starts around pH 2.9, not pH 1) because the weak acid is only partially dissociated. Second, there is a buffer region — a long, relatively flat section before the equivalence point where pH changes slowly because the solution contains both the weak acid and its conjugate base, forming a buffer. The midpoint of this buffer region is the half-equivalence point, where pH equals pKa. This is a high-yield fact for exams: at the half-equivalence point, half the acid has been neutralized, so the acid and conjugate base concentrations are equal, and the Henderson-Hasselbalch equation simplifies to pH = pKa.

The equivalence point for a weak acid-strong base titration is above pH 7 — typically around pH 8-9. This is because the product at equivalence is the conjugate base of the weak acid, which hydrolyzes water to produce OH-. This shift is the reason you cannot use the same indicator for every titration.

An indicator is a weak acid itself that changes color at a specific pH range. The key rule: the indicator's color change range must overlap with the steep portion of the titration curve near the equivalence point. If it does not, you will either overshoot or undershoot.

For strong acid-strong base titrations (equivalence at pH 7), the steep region spans roughly pH 4 to pH 10, so almost any indicator works — phenolphthalein (pH 8.2-10), methyl red (pH 4.4-6.2), or bromothymol blue (pH 6.0-7.6) are all acceptable. Phenolphthalein is the traditional choice because the color change from clear to pink is easy to see.

For weak acid-strong base titrations (equivalence around pH 8-9), you need an indicator that changes in the basic range. Phenolphthalein (8.2-10) is appropriate. Methyl red (4.4-6.2) would change color too early — before you reach equivalence — giving a false endpoint and an inaccurate concentration.

For weak base-strong acid titrations (equivalence around pH 5-6), you need an indicator that changes in the acidic range. Methyl red or methyl orange (pH 3.1-4.4) works. Phenolphthalein would change too late.

The exam question is almost always: given this titration type, which indicator gives the most accurate endpoint? The answer comes from matching the indicator range to the equivalence pH. ChemistryIQ has practice problems that present titration curve shapes and ask you to select the correct indicator from a list.

The most frequent error is confusing the equivalence point with the endpoint. The equivalence point is the theoretical point where moles of acid equal moles of base — it is a calculated value. The endpoint is the experimentally observed point where the indicator changes color. In a well-designed titration, the endpoint is very close to the equivalence point, but they are not identical. Exams love to test whether you know this distinction.

Second, students often forget that the pH at the equivalence point is not always 7. It is only 7 for strong acid-strong base. For weak acid-strong base, the equivalence pH is above 7. For weak base-strong acid, it is below 7. If an exam question asks for the pH at equivalence for an acetic acid and NaOH titration, the answer is not 7 — you need to calculate the hydrolysis of the conjugate base (acetate ion).

Third, watch your sig figs on molarity calculations. If the buret reads to 0.01 mL precision, your volume measurements have four significant figures. Do not round intermediate calculations — carry extra digits and round only the final answer. Premature rounding is a point-killer on quantitative titration problems.