How to do a salt preparation by titration

Preparing a salt by titration is the one salt-preparation method every SPM student must know for the sodium, potassium and ammonium salts. It looks fiddly at first, but it follows a fixed sequence, and once you understand why each step is there, it is hard to forget. This guide walks through the full method and the thinking behind it.

Why titration is needed for these salts

For most soluble salts you react an acid with an excess of an insoluble solid, then filter the leftover away. That trick is impossible for a sodium, potassium or ammonium salt, because the base you would use, sodium hydroxide, potassium hydroxide or ammonia solution, is itself soluble. There is no solid left over to filter, so you can never tell by looking when the acid has been exactly neutralised.

Titration solves this by measuring the exact volume of acid that neutralises a fixed volume of alkali, using an indicator to signal the end point. The reaction is a straightforward neutralisation of acid and alkali: for sodium chloride, NaOH + HCl → NaCl + H₂O.

The apparatus you need

• A pipette (usually 25.0 cm³) to measure a fixed, accurate volume of the alkali.
• A burette (50.0 cm³) to deliver the acid a drop at a time and measure how much is used.
• A conical flask to hold the alkali and swirl the mixture.
• A suitable indicator, phenolphthalein or methyl orange, and a white tile to see the colour change clearly.

Stage 1: Find the exact volume by titration (with indicator)

1. Use the pipette to transfer 25.0 cm³ of the alkali (say sodium hydroxide) into the conical flask. Add two or three drops of phenolphthalein, which turns the alkali pink.
2. Fill the burette with the acid (dilute hydrochloric acid) and record the initial reading at eye level.
3. Add the acid slowly, swirling the flask constantly. As the end point nears, the pink colour flashes and fades where the acid lands.
4. Add the last drops one at a time until a single drop turns the solution from pink to colourless, this is the end point. Record the final reading.
5. The volume of acid used is final minus initial. Repeat the titration until you get consistent readings, and take the average as the exact volume, V cm³.

Stage 2: Repeat without indicator to get a pure salt

The solution you just made contains indicator, so it is contaminated and cannot be crystallised as a pure salt. To fix this, repeat the neutralisation using the same volumes but no indicator:

6. Pipette a fresh 25.0 cm³ of the alkali into a clean conical flask, this time add no indicator.
7. From the burette, run in exactly V cm³ of the acid, the volume you found in Stage 1. Because you already know the neutralising volume, the mixture is now neutral salt solution with nothing else added.

Stage 3: Crystallise the salt

8. Pour the neutral solution into an evaporating dish and heat gently to evaporate about one-third to a half of the water, until a hot saturated solution forms. Test by dipping a cold glass rod; crystals on the rod mean it is ready.
9. Stop heating and leave the solution to cool slowly. Slow cooling grows larger, well-formed crystals of sodium chloride.
10. Filter off the crystals, wash with a very small amount of cold distilled water, and dry them between filter papers.

Worked idea: why not just evaporate to dryness?

For a salt like sodium chloride you could evaporate fully, but for hydrated salts the water of crystallisation matters, so the standard, safe instruction is to evaporate only to the saturated point and let crystals form on cooling. Boiling a solution dry can also make crystals spit or decompose. Stopping at the saturated point is the habit to build for every salt you crystallise.

The exam logic in one line

The whole method exists to answer two problems at once: how do I know when neutralisation is complete (answer: an indicator and a titration), and how do I keep the salt pure (answer: repeat without indicator). If you can state those two reasons, you understand the method rather than just memorising it. You can see the titration technique demonstrated step by step in the acid–base titration experiment.

Linking the practical to the calculation

Salt preparation by titration is the practical twin of the number work in a titration calculation. The same burette and pipette readings that make a pure salt in the lab also feed the mole calculation that finds an unknown concentration in Paper 2. Practising both together, the procedure and the sums, means one topic earns marks across the practical paper and the written paper, which is why the acids, bases and salts chapter is such good value to master.

Where students lose marks

• Forgetting the second, indicator-free stage, the most common slip, and it costs the purity marks.
• Reading the burette wrongly, always read the bottom of the meniscus at eye level.
• Rinsing the pipette or burette with water only before use, which dilutes the solution; rinse each with the solution it will hold.
• Evaporating to dryness instead of stopping at the saturated point.

If the sequence keeps scrambling under pressure, a teacher can rehearse it with you against past Paper 3 questions until it runs automatically. Our online one-to-one lessons run in English from RM50 an hour, with a paid one-hour trial, so the method is taught and drilled directly on the questions you find hardest.