Acid-Base Titration — Procedure, Calculations & Indicators

Acid-base titration is a method for finding an unknown acid or base concentration by reacting it with a standard solution of known concentration. You add the standard solution until the reaction reaches the equivalence point, where the acid and base have reacted in the exact mole ratio from the balanced equation.

In a typical lab setup, you measure the sample, add titrant from a burette, watch for the endpoint, and use the titrant volume to calculate the unknown concentration. If the reaction is $1{:}1$, the shortcut $C_aV_a = C_bV_b$ works. If it is not $1{:}1$, you must use the balanced equation first.

Why Acid-Base Titration Works

The method works because neutralization follows stoichiometry. At equivalence, the reacting amounts are not just "close." They match the balanced chemical equation.

That condition matters. HCl and NaOH react in a $1{:}1$ ratio, but sulfuric acid and sodium hydroxide react in a $1{:}2$ ratio:

$$\mathrm{H_2SO_4} + 2\mathrm{NaOH} \rightarrow \mathrm{Na_2SO_4} + 2\mathrm{H_2O}$$

So volume alone is never enough. You also need the mole ratio from the equation.

Acid-Base Titration Procedure Step by Step

A standard classroom or lab titration usually follows these steps:

1. Rinse and fill the burette with the standard solution.
2. Pipette a known volume of the unknown solution into a flask.
3. Add a few drops of a suitable indicator, or use a pH probe.
4. Add titrant while swirling, then go dropwise near the endpoint.
5. Record the burette reading when the endpoint appears.
6. Repeat until you get consistent titres.

The endpoint is the visible signal, such as a color change. The equivalence point is the stoichiometric point from the chemistry. In textbook problems they are often treated as the same, but in real lab work they are only close to each other.

Acid-Base Titration Formula and Calculation

The safest approach is to start from moles and the balanced equation:

$$\frac{n_{\text{acid}}}{a} = \frac{n_{\text{base}}}{b}$$

where $a$ and $b$ are the stoichiometric coefficients from the balanced equation.

Because $n = CV$, this becomes:

$$\frac{C_{\text{acid}}V_{\text{acid}}}{a} = \frac{C_{\text{base}}V_{\text{base}}}{b}$$

This is the general acid-base titration relation.

The shortcut

$$C_aV_a = C_bV_b$$

works only when the reaction is $1{:}1$.

Worked Example: Finding an Unknown HCl Concentration

Suppose $25.0\ \mathrm{mL}$ of hydrochloric acid is titrated with $0.100\ \mathrm{mol/L}$ sodium hydroxide. The endpoint is reached after $18.4\ \mathrm{mL}$ of NaOH is added. Find the concentration of the HCl.

First write the reaction:

$$\mathrm{HCl} + \mathrm{NaOH} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}$$

This is a $1{:}1$ reaction, so the moles of HCl and NaOH are equal at equivalence.

Find the moles of NaOH added:

$$n(\mathrm{NaOH}) = CV = 0.100 \times 0.0184 = 0.00184\ \mathrm{mol}$$

So the HCl sample also contained:

$$n(\mathrm{HCl}) = 0.00184\ \mathrm{mol}$$

Now divide by the acid volume in liters:

$$C(\mathrm{HCl}) = \frac{0.00184}{0.0250} = 0.0736\ \mathrm{mol/L}$$

So the acid concentration is:

$$0.0736\ \mathrm{mol/L}$$

This example shows the logic of most titration problems: known concentration plus measured volume gives moles, and moles give the unknown concentration.

How To Choose the Right Indicator

An indicator should change color over the steep part of the titration curve. If its transition range sits far from the sharp pH change, it can signal the endpoint too early or too late.

For a strong acid-strong base titration, several common indicators can work because the pH jump near equivalence is large. For a weak acid-strong base titration, an indicator with a transition range on the basic side is usually more suitable, because the equivalence point is typically above $7$ in dilute aqueous solution. For a strong acid-weak base titration, an indicator that changes in the acidic range is often a better fit.

The practical rule is simple: match the indicator to the titration curve, not just the reagent names.

Common Acid-Base Titration Mistakes

Using $C_aV_a = C_bV_b$ Automatically

That shortcut is valid only for a $1{:}1$ reaction. If the coefficients differ, use the balanced equation first.

Mixing Up Endpoint and Equivalence Point

The endpoint is what you observe. The equivalence point is what the stoichiometry defines. They should be close, but they are not identical by definition.

Overshooting Near the Endpoint

Most of the useful volume is added safely before the last few drops. Near the endpoint, one extra drop can matter.

Mixing Volume Units

If you calculate moles with $n = CV$, use concentration in $\mathrm{mol/L}$ and volume in $\mathrm{L}$. If you use the ratio form on both sides, the volume units must still match.

When Acid-Base Titration Is Used

Acid-base titration is used to find unknown concentrations, standardize solutions, and check the concentration of a sample. It also turns the idea of neutralization into a measurable lab method.

It is most useful when the neutralization reaction is well defined and the endpoint can be detected clearly.

Try a Similar Problem

Change the worked example so the NaOH volume is $22.0\ \mathrm{mL}$ instead of $18.4\ \mathrm{mL}$, then solve it again. After that, try a non-$1{:}1$ case such as $\mathrm{H_2SO_4}$ with NaOH and notice that the setup changes before the arithmetic does.