What are the common methods for the preparation of salts in Cambridge IGCSE Chemistry?

In Cambridge IGCSE Chemistry, salts can be prepared using several methods, including the reaction of an acid with a base (neutralization), reaction of an acid with a metal, reaction of an acid with a carbonate, and precipitation reactions. Each method is chosen based on the solubility and reactivity of the reactants and the desired salt.

How do you prepare a salt using an acid and a base?

To prepare a salt using an acid and a base, you perform a neutralization reaction. For example, when hydrochloric acid (HCl) reacts with sodium hydroxide (NaOH), sodium chloride (NaCl) and water (H2O) are formed. The general equation is: Acid + Base → Salt + Water. This method is often used to prepare soluble salts.

What is the role of titration in the preparation of salts?

Titration is a technique used to prepare salts by accurately measuring the volume of acid needed to completely react with a base. This method is particularly useful for preparing soluble salts and ensures that the acid and base react in the correct stoichiometric proportions. It involves adding the acid to the base until the endpoint is reached, indicated by a color change of an indicator.

Can you explain how a precipitation reaction is used to prepare salts?

A precipitation reaction involves mixing two solutions containing ions that react to form an insoluble salt, which precipitates out of the solution. For example, when solutions of barium chloride (BaCl2) and sulfuric acid (H2SO4) are mixed, barium sulfate (BaSO4), an insoluble salt, precipitates. This method is used to prepare insoluble salts.

Why is it important to consider the solubility of salts in their preparation?

The solubility of salts is crucial in their preparation because it determines the method used. Soluble salts can be prepared by neutralization or titration, while insoluble salts are typically prepared by precipitation. Understanding solubility helps in choosing the correct reactants and method to efficiently produce the desired salt without unnecessary byproducts.

Why is "Using filtering for a soluble salt prep with soluble reactants" a common mistake in Preparation of Salts?

Why it happens: Mixing methods. How to avoid it: Soluble + soluble → titration. Soluble acid + INSOLUBLE base → excess + filter. Soluble + soluble giving INSOLUBLE product → precipitation. Source: 0620/42 — recurring

Why is "Skipping the acidification step in ion tests" a common mistake in Preparation of Salts?

Why it happens: Speed. How to avoid it: Acidify FIRST to remove carbonates that would also give white precipitates.

Why is "Not washing the precipitate before drying" a common mistake in Preparation of Salts?

Why it happens: Forgetting purity. How to avoid it: Wash the residue with DISTILLED water to remove dissolved impurities.

Why is "Boiling the solution to dryness for crystallisation" a common mistake in Preparation of Salts?

Why it happens: Trying to be quick. How to avoid it: Evaporate SOME water; leave to crystallise. Boiling dry damages crystal shape and may decompose the salt.

Acids, Bases and SaltsCambridge IGCSE Chemistry (0620)

Preparation of Salts

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Preparation of Salts — Cambridge IGCSE 0620 Chemistry Extended (2026)

Three main methods: insoluble base + acid (titration not needed), insoluble salt by precipitation, and soluble salt by titration. Choose the right route for the salt you want.

What you’ll learn

Mapped to the Cambridge IGCSE 0620 syllabus (2026-2028).

9.3 — Describe methods of salt preparation including titration, precipitation and direct combination.
9.3 — Apply solubility rules to predict whether a salt is soluble.

Solubility rules

Memorise the soluble groups; everything else needs checking.

Always soluble.

Group 1 metal salts (Li⁺, Na⁺, K⁺): always.
Ammonium ( $\text{NH}_{4}^{+}$ ) salts: always.
All NITRATES: always.

Mostly soluble (with exceptions).

CHLORIDES: soluble EXCEPT silver ( $\text{AgCl}$ ) and lead ( $\text{PbCl}_{2}$ ).
SULFATES: soluble EXCEPT barium ( $\text{BaSO}_{4}$ ), lead ( $\text{PbSO}_{4}$ ), and calcium (sparingly).

Mostly insoluble (with exceptions).

CARBONATES: insoluble EXCEPT Group 1 + ammonium.
HYDROXIDES: insoluble EXCEPT Group 1 + ammonium + Ca(OH)₂ (sparingly).
OXIDES: insoluble EXCEPT Group 1 metal oxides and CaO.

Worked example. Is sodium nitrate soluble? Yes — Group 1 + nitrate (both rules say soluble).

Worked example. Is silver chloride soluble? NO — silver is the chloride exception.

Worked example. Is iron(II) carbonate soluble? NO — carbonates are insoluble except Group 1/ammonium.

Group 1 + ammonium: always soluble.
Nitrates: always soluble.
Chlorides: usually soluble (not Ag, Pb).
Sulfates: usually soluble (not Ba, Pb).
Carbonates / hydroxides / oxides: usually NOT soluble.

Method 1 — Soluble salt from insoluble base / metal / carbonate

Add excess solid to acid; filter off excess; evaporate to get crystals. Used when the base/metal/carbonate is insoluble.

When to use. The salt is SOLUBLE; the base / metal / carbonate is INSOLUBLE.

Examples.

Copper(II) sulfate from CuO + H₂SO₄ (CuO is insoluble).
Magnesium chloride from Mg + HCl.
Calcium nitrate from CaCO₃ + HNO₃.

Steps.

Warm the acid in a beaker.
Add the solid (insoluble base / metal / carbonate) IN EXCESS — keep adding until no more reacts (you'll see undissolved solid).
FILTER off the excess solid.
The FILTRATE is the salt solution.
Evaporate gently in an evaporating basin until ABOUT HALF the water remains (crystallisation point).
Cool, leave to crystallise.
Filter / decant off mother liquor; pat crystals dry.

Add solid in excess, filter it off, then evaporate the filtrate to crystallise the soluble salt.

Worked. Make copper(II) sulfate.

Add black CuO to warm dilute H₂SO₄ until no more reacts. Solution turns blue.
Filter off excess CuO.
Evaporate the blue filtrate, cool. Blue crystals of $\text{CuSO}_{4}\cdot 5\text{H}_{2}\text{O}$ form.

Cambridge tip. Add solid IN EXCESS so all the acid is used up; otherwise the residue includes acid impurity.

Excess insoluble solid + acid → salt solution.
Filter to remove excess.
Evaporate, crystallise.
Examples: CuSO₄, MgCl₂, Ca(NO₃)₂.

Method 2 — Soluble salt by titration (acid + alkali)

Both reactants soluble. Use indicator to find exact volumes; repeat without indicator to make pure salt.

When to use. Both acid and base are SOLUBLE (e.g. NaOH alkali). You need to know exactly when neutralisation completes.

Examples.

Sodium chloride from HCl + NaOH.
Potassium nitrate from HNO₃ + KOH.
Sodium sulfate from H₂SO₄ + NaOH.

Steps.

Pipette a known volume of alkali into a conical flask.
Add a few drops of indicator (e.g. methyl orange or phenolphthalein).
Titrate with the acid from a burette until the colour just changes (end-point).
Record the titre. Repeat for concordance.
Repeat the titration WITHOUT indicator, using the average volumes — gives a pure, indicator-free salt solution.
Evaporate gently and crystallise.

Why repeat without indicator? Indicator dye would contaminate the salt crystals.

Add acid from the burette until the indicator just changes — then repeat with no indicator for a pure salt.

Worked. Make pure sodium chloride.

Titrate $25\,\text{cm}^{3}$ of $0.1\,\text{mol/dm}^{3}$ NaOH against $0.1\,\text{mol/dm}^{3}$ HCl. End-point at $\sim 25\,\text{cm}^{3}$ HCl.
Repeat with the same volumes but no indicator.
Evaporate the resulting NaCl(aq) → NaCl crystals.

Both reactants soluble: use titration.
Indicator finds end-point.
Repeat WITHOUT indicator for pure salt.
Common indicators: methyl orange, phenolphthalein.

Method 3 — Insoluble salt by precipitation

Mix two soluble salts whose ions form your insoluble target. Filter, wash, dry.

When to use. The salt is INSOLUBLE.

Examples.

Silver chloride: $\text{AgNO}_{3}(aq) + \text{NaCl}(aq) \to \text{AgCl}(s) + \text{NaNO}_{3}(aq)$ . AgCl precipitates as white solid.
Barium sulfate: $\text{BaCl}_{2}(aq) + \text{Na}_{2}\text{SO}_{4}(aq) \to \text{BaSO}_{4}(s) + 2\text{NaCl}(aq)$ . White precipitate.
Lead(II) iodide: $\text{Pb(NO}_{3}\text{)}_{2}(aq) + 2\text{KI}(aq) \to \text{PbI}_{2}(s) + 2\text{KNO}_{3}(aq)$ . Bright yellow precipitate ("golden rain").

Steps.

Mix solutions of two soluble salts that contain the wanted cation and anion.
The insoluble salt PRECIPITATES.
FILTER off the solid.
WASH with distilled water (to remove soluble salt impurities).
DRY (in oven or air).

Cambridge tip. Cambridge often asks "why wash with water?" Answer: to remove soluble impurities (the spectator ions and any unreacted reagent).

Two soluble salts mix; the insoluble product drops out as a precipitate to be filtered, washed and dried.

Mix two soluble salts → precipitate insoluble salt.
Filter, wash, dry.
Examples: AgCl, BaSO₄, PbI₂.
Wash to remove spectator ions.

How it’s examined

Salt preparation appears every Paper 4 (8-10 marks): identify the right method, write balanced equation, list practical steps. Examiner reports flag students leaving out 'in excess' for Method 1 and forgetting the indicator-free repeat for Method 2.

Sources: Cambridge IGCSE Chemistry 0620 syllabus 2026-2028 (9.3); 0620/42 Oct/Nov 2024 — Q15 (salt preparation); 0620 Examiner Reports 2022-2024. Last reviewed 2026-05-07.

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Step-by-step worked examples — Preparation of Salts

Step-by-step solutions to past-paper-style questions on preparation of salts, written exactly the way a tutor would explain them at the board.

1Soluble salt by acid + insoluble base (titration not needed)
Extended• Adapted from 0620/42 May/Jun 2024 Q10• soluble salt
Question
Describe how to prepare pure copper(II) sulfate crystals from copper(II) oxide and dilute sulfuric acid.
Step-by-step solution
1. Step 1
  Warm the dilute $\text{H}_{2}\text{SO}_{4}$ and add EXCESS CuO until no more dissolves.
2. Step 2
  FILTER off unreacted CuO.
3. Step 3
  Heat the filtrate to evaporate some water → leave to crystallise.
4. Step 4
  Filter, then dry the crystals between filter paper.
Answer
Acid + excess insoluble base → filter → evaporate to crystallise → dry.
2Insoluble salt by precipitation
Extended• precipitation
Question
Describe how to prepare pure barium sulfate.
Step-by-step solution
1. Step 1
  Mix solutions of barium chloride and sodium sulfate.
  $\text{BaCl}_{2} + \text{Na}_{2}\text{SO}_{4} \to \text{BaSO}_{4}\downarrow + 2\text{NaCl}$
2. Step 2
  Filter.
3. Step 3
  Wash the residue with distilled water to remove soluble salts.
4. Step 4
  Dry the residue.
Answer
Mix two suitable solutions → filter → wash → dry the precipitate.
3Soluble salt from acid + alkali (titration)
Extended• titration
Question
Describe how to prepare pure sodium chloride from HCl and NaOH (both soluble).
Step-by-step solution
1. Step 1
  Titrate to find the exact volume of HCl needed to neutralise the NaOH (using indicator).
2. Step 2
  Repeat WITHOUT indicator using the recorded volume.
3. Step 3
  Heat to crystallise the resulting NaCl solution.
Answer
Titrate (with indicator) → repeat without indicator → crystallise.
4Identify a sulfate ion
Extended• ion test
Question
Describe a chemical test for the sulfate ion.
Step-by-step solution
1. Step 1
  Acidify with dilute hydrochloric acid (to remove carbonates).
2. Step 2
  Add aqueous barium chloride/nitrate.
3. Step 3
  WHITE precipitate of $\text{BaSO}_{4}$ confirms sulfate.
Answer
Acidify with HCl, add $\text{BaCl}_{2}$ → white precipitate.

Key Definitions and Keywords — Preparation of Salts

Definitions to memorise and the exact keywords mark schemes credit for preparation of salts answers — sharpened from recent examiner reports for the 2026 0620 sitting.

Salt
Examiner keyword
Compound formed when the H of an acid is replaced by a metal ion (or $\text{NH}_{4}^{+}$ ).
Precipitation
Examiner keyword
Formation of an insoluble solid (precipitate) when two solutions react.
Crystallisation
Examiner keyword
Slow formation of crystals from a saturated solution as solvent evaporates.

Common Mistakes and Misconceptions — Preparation of Salts

The traps other students keep falling into on preparation of salts questions — taken from recent Cambridge IGCSE 0620 examiner reports and mark schemes — and how to avoid them.

✕Using filtering for a soluble salt prep with soluble reactants
0620/42 — recurring
Why it happens
Mixing methods.
How to avoid it
Soluble + soluble → titration. Soluble acid + INSOLUBLE base → excess + filter. Soluble + soluble giving INSOLUBLE product → precipitation.
✕Skipping the acidification step in ion tests
Why it happens
Speed.
How to avoid it
Acidify FIRST to remove carbonates that would also give white precipitates.
✕Not washing the precipitate before drying
Why it happens
Forgetting purity.
How to avoid it
Wash the residue with DISTILLED water to remove dissolved impurities.
✕Boiling the solution to dryness for crystallisation
Why it happens
Trying to be quick.
How to avoid it
Evaporate SOME water; leave to crystallise. Boiling dry damages crystal shape and may decompose the salt.

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Preparation of Salts — frequently asked questions

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Preparation of Salts

Short Study Notes in the form of Slides

Detailed Notes

What you’ll learn

How it’s examined

1Soluble salt by acid + insoluble base (titration not needed)

2Insoluble salt by precipitation

3Soluble salt from acid + alkali (titration)

4Identify a sulfate ion

Salt

Precipitation

Crystallisation

✕Using filtering for a soluble salt prep with soluble reactants

✕Skipping the acidification step in ion tests

✕Not washing the precipitate before drying

✕Boiling the solution to dryness for crystallisation

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Preparation of Salts — frequently asked questions