Acid Base Titration A Level Chemistry Guide

Acid-base titration at A level is the precise technique of adding a solution of known concentration (the titrant) from a burette to a measured volume of another solution until the reaction is exactly complete, allowing you to calculate the unknown concentration of the second solution. You enter your known concentration, volumes, and the stoichiometric ratio from the balanced equation, and the calculator returns the unknown concentration in mol dm�³. The result is only as good as your burette readings and your choice of indicator, so this guide covers the full method, the calculations, and the common traps that lose marks in exams.

Formula

The core calculation is a rearrangement of the definition of concentration: concentration (mol dm�³) = moles ÷ volume (dm³). For a titration, the moles of the known solution (titrant) are found first, then the mole ratio from the balanced equation gives the moles of the unknown, and finally the unknown concentration is calculated.

Step 1: Moles of known solution = c� × (v� ÷ 1000)

Step 2: Moles of unknown solution = moles of known × r

Step 3: Concentration of unknown = moles of unknown ÷ (v₂ ÷ 1000)

Where:

• câ‚� = concentration of the known solution (mol dmâ�»Â³)
• vâ‚� = volume of known solution from the burette (cm³, converted to dm³ by dividing by 1000)
• vâ‚‚ = volume of unknown solution pipetted into the flask (cm³)
• r = stoichiometric ratio: moles of unknown ÷ moles of known, taken from the balanced equation

Worked Example

Using the default values: a 0.100 mol dm�³ solution of hydrochloric acid is titrated against 25.00 cm³ of sodium hydroxide solution, and the endpoint is reached when 25.00 cm³ of acid has been added. The balanced equation is HCl + NaOH → NaCl + H₂O, so the ratio r = 1.

1. Moles of HCl = 0.100 × (25.00 ÷ 1000) = 0.00250 mol
2. Moles of NaOH = 0.00250 × 1 = 0.00250 mol
3. Concentration of NaOH = 0.00250 ÷ (25.00 ÷ 1000) = 0.1000 mol dm�³

The calculator gives 0.1000 mol dm�³, matching the manual calculation. Notice that the volumes are equal and the ratio is 1, so the concentrations are equal – a useful sanity check.

The Practical Method

A level titration requires precision at every step. The standard procedure is:

• Rinse the burette with the titrant (the known solution) and fill it, ensuring no air bubbles in the jet.
• Use a pipette to transfer an exact volume (usually 25.00 cm³) of the unknown solution into a conical flask.
• Add 2–3 drops of indicator (see the next section for choosing one).
• Run the titrant from the burette while swirling the flask continuously, slowing to a drop-by-drop pace near the endpoint.
• Record the initial and final burette readings to two decimal places (e.g., 0.00 cm³ and 24.80 cm³).
• Repeat until you have two concordant results – typically within 0.10 cm³ of each other – and use the mean titre.

The first titration is often a rough run to see the approximate endpoint, then subsequent runs are done carefully. The mean is calculated only from concordant results, never from the rough run.

Choosing the Right Indicator

The indicator must change colour at the pH of the equivalence point, which depends on the strength of the acid and base. The pH at equivalence is exactly 7 only for a strong acid–strong base titration. For other combinations, the equivalence point is not neutral:

• Strong acid + strong base: equivalence pH = 7. Use methyl orange (3.1–4.4) or phenolphthalein (8.3–10.0) – both work because the pH jump crosses 7 steeply.
• Strong acid + weak base: equivalence pH < 7 (e.g., ammonium chloride from HCl + NH₃). Use methyl orange.
• Weak acid + strong base: equivalence pH > 7 (e.g., sodium ethanoate from ethanoic acid + NaOH). Use phenolphthalein.
• Weak acid + weak base: no sharp pH jump, so no simple indicator works; this combination is rarely titrated at A level.

A common mistake is assuming the endpoint is always pH 7. It is not – the indicator is chosen to match the equivalence pH, not the neutral pH.

Common Mistakes and How to Avoid Them

Exam markers repeatedly see the same errors. Watch for these:

• Forgetting to convert cm³ to dm³: always divide volumes by 1000 before using them in concentration calculations.
• Using the wrong mole ratio: for a diprotic acid like Hâ‚‚SOâ‚„ reacting with NaOH, the ratio is 1:2, not 1:1. Write the balanced equation first.
• Reading the burette incorrectly: read at eye level, from the bottom of the meniscus, and record to 0.05 cm³ (two decimal places).
• Using the rough titre in the mean: the rough run is only for estimating the endpoint, never for the final calculation.
• Not rinsing equipment properly: the burette must be rinsed with the titrant, the pipette with the unknown solution, but the conical flask should be rinsed with distilled water only – any residual solution would change the moles present.
• Overshooting the endpoint: if you add too much titrant, the result is invalid. Start again rather than trying to correct it.

Interpreting the Result and Its Limits

The calculated concentration is only as accurate as the least precise measurement. A burette reading is typically ±0.05 cm³, so a titre of 25.00 cm³ has an uncertainty of about ±0.2%. The pipette is usually accurate to ±0.06 cm³, and the balance (if you prepare the known solution by weighing) adds its own uncertainty. The final answer should be quoted to the same number of significant figures as the least precise data – usually 3 or 4. Also note that the formula assumes the reaction goes to completion and that the indicator changes colour exactly at the equivalence point; in practice, a small overshoot (the endpoint) is slightly different from the equivalence point, but the error is negligible for strong acid–strong base titrations.

Alternatives and Extensions

If you cannot use a burette, a pH meter or data logger can plot the pH against volume added, giving a curve that shows the equivalence point as the steepest part. This method removes the need for an indicator and works for weak acid–weak base pairs. A conductivity titration measures the change in ionic conductivity instead of pH, which is useful for very dilute or coloured solutions where indicators fail. At A level, you may also encounter back titration, where an excess of a known reagent is added to the unknown, and the remaining excess is titrated – this is used when the reaction is slow or the substance is insoluble.

Frequently Asked Questions

Why do I divide the volume by 1000? Concentration is defined in mol dm�³, but burettes and pipettes measure in cm³. Since 1 dm³ = 1000 cm³, you must convert the volume to dm³ before calculating moles. Forgetting this step is the single most common calculation error in A level titration questions.

What is the difference between endpoint and equivalence point? The equivalence point is when the moles of acid and base are exactly stoichiometrically equal. The endpoint is when the indicator changes colour. A good indicator is chosen so the endpoint is as close as possible to the equivalence point, but they are technically different moments.

Can I use any indicator for any titration? No. The indicator must have a colour change range that includes the pH at the equivalence point. Using phenolphthalein for a strong acid–weak base titration would give a result that is too high because the colour change happens after the equivalence point.

What does a concordant result mean? Concordant titres are readings that agree within 0.10 cm³ of each other. You should repeat the titration until you have at least two concordant results, then average only those. The rough titre is never included in the mean.

How do I find the mole ratio if the equation is not given? You must write and balance the full equation. For example, H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O shows that 1 mole of sulfuric acid reacts with 2 moles of sodium hydroxide, so the ratio r = 2 for the unknown (NaOH) relative to the known (H₂SO₄).

Why is my calculated concentration slightly different from the expected value? Small differences come from measurement uncertainties (burette, pipette, indicator overshoot) and from the assumption that the known solution is exactly the concentration you think it is – if you prepared it by weighing, the solid may have absorbed moisture. These are normal sources of experimental error.

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in c1 | Concentration of known solution | mol/dm³ | 0.100 | 0
in v1 | Volume of known solution used | cm³ | 25.00 | 0
in v2 | Volume of unknown solution pipetted | cm³ | 25.00 | 0
in r | Stoichiometric ratio (mol unknown : mol known) | - | 1 | 0
out c2 | Concentration of unknown solution | mol/dm³ | (c1 * v1) / (v2 * r) | 4
out moles1 | Moles of known solution | mol | (c1 * v1) / 1000 | 5
out moles2 | Moles of unknown solution | mol | ((c1 * v1) / 1000) * r | 5