SingaporeCombined ScienceSyllabus dot point

How do we choose a method to prepare a pure, dry sample of a salt, and how does solubility decide that choice?

Describe the preparation of soluble and insoluble salts by titration, excess solid and precipitation, using the solubility rules to select the method

A focused answer to the O-Level Combined Science outcome on making salts. The solubility rules, preparing soluble salts by excess solid and by titration, making insoluble salts by precipitation, and obtaining a pure dry sample.

Generated by Claude Opus 4.810 min answerUpdated 2026-06-06

Reviewed by: AI editorial process; not yet individually human-reviewed

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What this dot point is asking
The answer
Examples in context
Try this

What this dot point is asking

SEAB wants you to use solubility rules to decide whether a salt is soluble or insoluble, then choose and describe the correct preparation method: excess solid or titration for soluble salts, and precipitation for insoluble salts. The practical steps and the reasons behind them (especially using excess and drying gently) are frequent structured-question marks.

The answer

The solubility rules

You need a working set of rules:

All sodium, potassium and ammonium salts are soluble.
All nitrates are soluble.
Most chlorides are soluble (except silver chloride and lead chloride).
Most sulfates are soluble (except barium sulfate, lead sulfate and calcium sulfate).
Most carbonates and hydroxides are insoluble (except those of sodium, potassium and ammonium).

These rules decide which preparation method applies.

Preparing a soluble salt by excess solid

When the salt is soluble and the base, metal or carbonate is insoluble, add it in excess to the acid:

add the solid (metal, insoluble base or carbonate) to warm acid until no more reacts,
filter off the unreacted excess solid,
evaporate the filtrate to a saturated solution, then cool to crystallise,
filter off the crystals and dry them gently.

Using excess guarantees all the acid is used up, so the product is not contaminated with leftover acid.

Preparing a soluble salt by titration

When both reactants are soluble (e.g. a soluble alkali such as sodium hydroxide with an acid), you cannot use excess, because any leftover would dissolve and contaminate the salt. Instead:

titrate the acid against the alkali using an indicator to find the exact volume needed,
repeat with the same volumes but no indicator (so the salt is pure),
evaporate and crystallise as before.

Preparing an insoluble salt by precipitation

When the salt is insoluble, mix two soluble solutions that supply the right ions; the insoluble salt drops out as a precipitate:

\text{Ba}^{2+}(\text{aq}) + \text{SO}_4^{2-}(\text{aq}) \rightarrow \text{BaSO}_4(\text{s})

Then filter to collect the precipitate, wash it with distilled water to remove soluble impurities, and dry it.

Worked example

Choose and outline a method to prepare a pure, dry sample of lead(II) iodide, given that it is insoluble in water. Name two suitable starting solutions and write the ionic equation.

Step 1: Check the solubility

Lead(II) iodide is insoluble, so the correct method is precipitation.

Step 2: Choose two soluble solutions

Choose solutions that supply lead ions and iodide ions and are themselves soluble, for example lead(II) nitrate solution (provides $\text{Pb}^{2+}$ ) and potassium iodide solution (provides $\text{I}^-$ ).

Step 3: Mix and precipitate

On mixing, a bright yellow precipitate of lead(II) iodide forms:

\text{Pb}^{2+}(\text{aq}) + 2\text{I}^-(\text{aq}) \rightarrow \text{PbI}_2(\text{s})

Step 4: Filter, wash and dry

Filter to collect the precipitate, wash the residue with distilled water to remove soluble potassium nitrate, then dry it in a warm oven.

Examples in context

Example 1. Barium meals for X-rays. Patients drink a suspension of barium sulfate before an abdominal X-ray because it is insoluble and so not absorbed or toxic, yet it blocks X-rays to show the gut. It is made industrially by precipitation, the same insoluble-salt method used in the lab.

Example 2. Growing pure crystals for fertiliser. Soluble salts such as ammonium sulfate are made on a large scale by neutralising acid with a base, then crystallising. Slow cooling gives larger, purer crystals, which is why the lab method of evaporating to saturation and cooling matters for product quality.

Try this

Q1. State the solubility rule for nitrates and for sodium salts. [2 marks]

Cue. All nitrates are soluble; all sodium salts are soluble.

Q2. A salt is soluble and is made from an insoluble carbonate. Name the preparation method and say why excess carbonate is used. [2 marks]

Cue. Excess solid method; excess ensures all the acid reacts, so no acid contaminates the salt.

Q3. Write the ionic equation for the precipitation of silver chloride and state its colour. [2 marks]

Cue. $\text{Ag}^+ + \text{Cl}^- \rightarrow \text{AgCl}$ ; the precipitate is white.

Exam-style practice questions

Practice questions written in the style of SEAB exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

Original5 marksA student is asked to prepare a pure, dry sample of copper(II) sulfate crystals from copper(II) oxide and dilute sulfuric acid. Describe the steps the student should follow and explain why excess copper(II) oxide is used.

Show worked answer →

Warm some dilute sulfuric acid and add copper(II) oxide a little at a time, stirring, until no more dissolves (the oxide is in excess). Excess is used to make sure all the acid has reacted, so no acid is left to contaminate the salt.
Filter the mixture to remove the unreacted copper(II) oxide. The filtrate is copper(II) sulfate solution.
Heat the filtrate gently to evaporate some water and form a saturated solution (until crystals start to appear at the edge).
Leave the solution to cool and crystallise slowly.
Filter off the crystals and dry them between sheets of filter paper (not by strong heating, which would remove the water of crystallisation).

Markers reward adding the oxide in excess until no more dissolves, filtering off the excess, crystallising by evaporation and cooling, drying gently, and the explanation that excess ensures all the acid is used up so the salt is pure.

Original4 marksBarium sulfate is an insoluble salt. Describe how a pure, dry sample of barium sulfate could be prepared by precipitation, naming two suitable soluble starting solutions.

Show worked answer →

Use two soluble solutions that together provide barium ions and sulfate ions, for example barium chloride solution and sodium sulfate solution (or barium nitrate and dilute sulfuric acid).

Mix the two solutions; a white precipitate of barium sulfate forms at once: $\text{Ba}^{2+} + \text{SO}_4^{2-} \rightarrow \text{BaSO}_4$ .

Filter to collect the precipitate. Wash the residue with distilled water to remove soluble impurities, then dry it in a warm oven or between filter papers.

Markers reward choosing two soluble solutions that supply the right ions, mixing to precipitate, and filtering then washing and drying the residue.