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Titration is the classic chemistry practical. It’s how you find the exact concentration of an acid or alkali by carefully neutralising it with a solution whose concentration you already know. It’s been used for over two centuries to test everything from vinegar and wine to drinking water and medicines, and it remains a standard method in quality-control labs today.
Done well, a school titration can measure concentration to within about 0.5%. This guide explains how to get there.
The principle
You react a known volume of one solution with a second solution of known concentration (the standard solution), adding it gradually until the reaction is exactly complete. An indicator changes colour to show when that point has been reached. From the volume used, you calculate the unknown concentration.
For hydrochloric acid and sodium hydroxide:
HCl + NaOH → NaCl + H₂O
The two react in a 1 : 1 ratio. At the equivalence point, the moles of acid added equal the moles of alkali present.
Equipment
| Item | Purpose |
|---|---|
| Burette (50.00 cm³) | Delivers the titrant in precisely measured amounts |
| Volumetric pipette (25.00 cm³) and pipette filler | Measures a fixed volume of the solution being analysed |
| Conical flask (250 cm³) | Holds the solution being titrated; its shape lets you swirl without splashing |
| Funnel | Filling the burette (removed before titrating) |
| White tile | Makes the colour change easier to see |
| Clamp stand and burette clamp | Holds the burette vertical |
| Wash bottle of distilled water | Rinsing |
| Indicator | Shows the end point |
Eye protection is essential throughout. Even dilute acids and alkalis are harmful in the eyes.
Preparing the equipment
Rinsing is where many accuracy problems start, so get it right.
- Burette: rinse with distilled water, then with a small amount of the solution it will hold. Drain the rinse through the tap. This removes water that would dilute the titrant.
- Pipette: rinse with distilled water, then with the solution it will measure.
- Conical flask: rinse with distilled water only. Water left in the flask doesn’t matter, because it doesn’t change the number of moles you pipette in.
The rule: equipment that measures a solution is rinsed with that solution; the flask that simply holds the reaction is rinsed with water.
Method
- Fill the burette using a funnel, with the tap closed, to above the 0.00 cm³ mark. Remove the funnel (drips from it can change the reading). Run some liquid through the tap to fill the tip and remove air bubbles.
- Read the initial volume at eye level, from the bottom of the meniscus, to the nearest 0.05 cm³. It doesn’t have to be exactly 0.00.
- Pipette 25.00 cm³ of the solution to be analysed into the conical flask. Let it drain freely and touch the tip against the inside of the flask; don’t blow out the last drop, because pipettes are calibrated to leave it behind.
- Add 2–3 drops of indicator.
- Rough titration: add titrant fairly quickly while swirling until the colour changes. Note the volume. This tells you roughly where the end point is.
- Accurate titrations: refill and repeat. Add titrant quickly until about 2 cm³ before the rough value, then drop by drop, swirling after each drop. Rinse the inside walls of the flask with a little distilled water near the end, to wash down any splashes. Stop at the first permanent colour change.
- Record the final reading and calculate the titre (final − initial).
- Repeat until you have at least two concordant results: titres within 0.10 cm³ of each other.
Recording results
Use a clear table with all readings to two decimal places ending in 0 or 5:
| Rough | 1 | 2 | 3 | |
|---|---|---|---|---|
| Final reading / cm³ | 24.10 | 23.85 | 47.30 | 24.00 |
| Initial reading / cm³ | 0.00 | 0.40 | 23.85 | 0.50 |
| Titre / cm³ | 24.10 | 23.45 | 23.45 | 23.50 |
Mean titre = (23.45 + 23.45 + 23.50) ÷ 3 = 23.47 cm³
Only concordant titres go into the mean. The rough titre is always left out.
Worked calculation
25.00 cm³ of sodium hydroxide solution of unknown concentration was titrated with 0.100 mol/dm³ hydrochloric acid. The mean titre was 23.47 cm³. Calculate the concentration of the sodium hydroxide.
Step 1. Moles of the known solution. Moles HCl = 0.100 × (23.47 ÷ 1000) = 2.347 × 10⁻³ mol
Step 2. Use the mole ratio. HCl : NaOH = 1 : 1, so moles NaOH = 2.347 × 10⁻³ mol
Step 3. Concentration of the unknown. [NaOH] = 2.347 × 10⁻³ ÷ (25.00 ÷ 1000) = 0.0939 mol/dm³
If the ratio isn’t 1 : 1, step 2 changes. For sulfuric acid and sodium hydroxide (H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O), moles of NaOH are twice the moles of H₂SO₄. More practice is in titration calculations and the molarity calculator.
Choosing the indicator
The indicator must change colour at a pH within the steep part of the titration curve.
| Titration | Indicator | Colour change (acid → alkali) |
|---|---|---|
| Strong acid + strong base | Phenolphthalein or methyl orange | colourless → pink, or red → yellow |
| Weak acid + strong base | Phenolphthalein | colourless → pink |
| Strong acid + weak base | Methyl orange | red → yellow |
| Weak acid + weak base | No simple indicator works well | use a pH meter |
Universal indicator is never used for titrations: it changes colour gradually across the whole pH range, so there’s no sharp end point. The reasons behind these choices are explained in how to choose an indicator.
Sources of error and how to reduce them
Systematic errors push every result in the same direction:
- Burette not rinsed with titrant → titrant diluted → titre too high.
- Pipette not rinsed with the solution → sample diluted → fewer moles → titre too low.
- Air bubble in the burette tip that escapes during the titration → volume reading larger than liquid actually delivered.
- Consistently overshooting the end point by a drop or two.
- Reading the meniscus from above or below eye level (parallax).
Random errors vary from run to run:
- Judging the exact colour change.
- Reading the burette to ±0.05 cm³ at each end.
Measurement uncertainty
A typical burette reading has an uncertainty of ±0.05 cm³. Each titre uses two readings, so its uncertainty is ±0.10 cm³. For a titre of 23.45 cm³:
Percentage uncertainty = 0.10 ÷ 23.45 × 100 = 0.43%
A 25.00 cm³ pipette typically has an uncertainty of ±0.06 cm³ (0.24%). To reduce percentage uncertainty, use concentrations that give titres of at least 20 cm³, so the fixed reading errors are a smaller fraction of the total.
Tips from experienced chemists
- Swirl constantly. Unmixed titrant can give a temporary colour change that fades.
- Near the end point, use half-drops: let a drop hang from the tip and wash it into the flask with distilled water.
- Put a sheet of white paper behind the burette to make the meniscus easier to read.
- Record readings immediately, in ink. Never “tidy up” a result afterwards.
- Always do a rough titration first. It saves time and reduces overshooting.
Key takeaways
- Titration finds an unknown concentration by exact neutralisation with a standard solution.
- Rinse the burette and pipette with the solutions they will hold; rinse the flask with water.
- Record readings to ±0.05 cm³ and average only concordant titres (within 0.10 cm³).
- Calculate moles of the known solution, apply the mole ratio, then divide by the volume of the unknown.
- Choose an indicator that changes colour in the steep part of the pH curve.
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