How to prepare soluble and insoluble salts

Salt preparation is one of the most practical topics in SPM Chemistry, and Paper 3 loves to ask you to describe a full method. The secret is that you cannot choose a method until you know whether the salt is soluble or insoluble, that single fact decides everything. This guide shows you how to make that decision and then carry out each method cleanly, step by step.

First, decide: soluble or insoluble?

Before writing a single step, classify the salt using the solubility rules. The rules you must know are:

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

If the salt is soluble, you make it in solution and then crystallise it. If it is insoluble, you make it fall out of solution as a precipitate. These are two completely different procedures.

Preparing a soluble salt (not sodium, potassium or ammonium)

For a soluble salt of a metal like copper, zinc, magnesium or iron, the method uses an acid and an excess of an insoluble solid. The insoluble solid can be a reactive metal, a metal oxide (base) or a metal carbonate. Using excess ensures all the acid reacts, and because the solid is insoluble you can simply filter off what is left over.

To prepare copper(II) sulfate from copper(II) oxide and sulfuric acid:

1. Warm dilute sulfuric acid in a beaker.
2. Add excess copper(II) oxide and stir until no more dissolves, the leftover black solid tells you all the acid has reacted: CuO + H₂SO₄ → CuSO₄ + H₂O.
3. Filter to remove the excess copper(II) oxide. The blue filtrate is copper(II) sulfate solution.
4. Evaporate the filtrate until a saturated solution forms, test by dipping a cold glass rod and looking for crystals.
5. Cool the concentrated solution slowly so crystals grow; this is crystallisation.
6. Filter the crystals, wash with a little cold distilled water, and dry between filter papers.

The same three-part logic, react with excess, filter off the excess, then crystallise, works whether you start from a metal, an oxide or a carbonate. You can see the full procedure written out in the preparation of a soluble salt experiment.

The exception: soluble salts of sodium, potassium and ammonium

The excess-solid method fails for sodium, potassium and ammonium salts, because their starting bases (sodium hydroxide, potassium hydroxide, ammonia) are themselves soluble. You cannot add an excess and filter it off, there is nothing to filter. Instead these salts are made by titration, measuring the exact volume of acid and alkali that react, then crystallising the product. Because both reactants are solutions, an indicator is needed to find the end point. This titration route has its own careful steps, which are worth learning on their own.

Preparing an insoluble salt by precipitation

An insoluble salt is made by precipitation, also called double decomposition. You mix two soluble solutions, each supplying one of the ions you need, and the insoluble salt forms instantly as a solid.

To prepare lead(II) iodide, a bright yellow insoluble salt:

1. Prepare a solution of a soluble lead salt, such as lead(II) nitrate, and a solution of a soluble iodide, such as potassium iodide.
2. Mix the two solutions. The insoluble salt precipitates at once: Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃.
3. Filter to collect the yellow precipitate; the other product (potassium nitrate) stays in solution and passes through.
4. Wash the residue on the filter paper with distilled water to remove any soluble impurities clinging to it.
5. Dry the precipitate between filter papers or in a warm oven.

The choice of starting solutions matters: both must be soluble, and the by-product must also be soluble so that only your target salt is left as the solid. The preparation of an insoluble salt by precipitation experiment shows more examples, such as barium sulfate and silver chloride.

A decision flow you can rely on

When a question names a salt, work through it in order:

1. Is it a sodium, potassium or ammonium salt? If yes, prepare it by titration.
2. Is it any other soluble salt? If yes, use acid plus excess metal, oxide or carbonate, then filter and crystallise.
3. Is it insoluble? If yes, use precipitation by mixing two soluble solutions, then filter, wash and dry.

That three-way split covers every salt the syllabus asks about, which is why the acids, bases and salts chapter treats the classification step as the heart of the topic.

Common mistakes in Paper 3 answers

• Wrong method for the family. Trying the excess-solid method on sodium chloride loses marks, it must be titration.
• Forgetting to filter the excess. For a soluble salt from an insoluble solid, the filtration step is what makes the product pure.
• Not washing the precipitate. An unwashed insoluble salt keeps soluble impurities and is not pure.
• Evaporating to dryness. For hydrated salts you crystallise, you must stop at the saturated point, not boil off all the water, or the crystals decompose or lose their shape.

Getting the method to stick

These procedures are easy to muddle under exam pressure, especially the point where soluble and insoluble methods diverge. Writing out each method from memory, then checking it against the experiment pages, fixes most errors. If the steps keep slipping, often because the solubility classification underneath is shaky, a teacher can walk you through them against real questions. Our online one-to-one lessons run in English from RM50 an hour, with a paid one-hour trial, so the whole method is taught directly against the salts you find hardest.