What are the key differences between acids and bases according to the Edexcel IGCSE Chemistry syllabus?

In the Edexcel IGCSE Chemistry syllabus, acids are substances that release hydrogen ions (H⁺) when dissolved in water, while bases are substances that can neutralize acids by accepting hydrogen ions. Acids typically have a pH less than 7, taste sour, and turn blue litmus paper red. Bases, on the other hand, have a pH greater than 7, feel slippery, and turn red litmus paper blue.

How is the pH scale used to measure acidity and alkalinity in the context of IGCSE Chemistry?

The pH scale is a numerical scale ranging from 0 to 14 used to measure the acidity or alkalinity of a solution. In IGCSE Chemistry, a pH less than 7 indicates an acidic solution, a pH of 7 is neutral, and a pH greater than 7 indicates an alkaline solution. The scale is logarithmic, meaning each whole number change represents a tenfold change in hydrogen ion concentration.

What are the common methods for preparing salts in IGCSE Chemistry?

In IGCSE Chemistry, salts can be prepared using several methods, including neutralization reactions between acids and bases, reaction of acids with metals, reaction of acids with metal carbonates, and precipitation reactions. Each method involves specific reactants and conditions to produce the desired salt.

Why is it important to understand the concept of neutralization in IGCSE Chemistry?

Neutralization is a fundamental concept in IGCSE Chemistry because it involves the reaction between an acid and a base to form a salt and water. This process is crucial for understanding how to control pH levels in various applications, such as agriculture, medicine, and environmental science. It also forms the basis for many laboratory experiments and industrial processes.

Can you explain the role of indicators in determining the pH of a solution as per the Edexcel IGCSE Chemistry curriculum?

Indicators are substances that change color depending on the pH of the solution they are in, and they are essential tools in the Edexcel IGCSE Chemistry curriculum for determining the acidity or alkalinity of a solution. Common indicators include litmus, phenolphthalein, and methyl orange, each with specific color changes at different pH levels. They help students visually identify the pH range of a solution during experiments.

Why is "Boiling the solution 'to dryness' to obtain crystals" a common mistake in Acids, Bases and Salt Preparations?

Why it happens: Trying to skip the slow-cooling step. How to avoid it: Never boil dry. (a) The crystals lose their water of crystallisation → wrong product. (b) Crystals are small, irregular and impure. ALWAYS evaporate to SATURATED only, then COOL SLOWLY for large pure crystals. Source: 4CH1 Examiner Reports 2022-2024

Why is "Forgetting to state that the base / metal / carbonate is added IN EXCESS" a common mistake in Acids, Bases and Salt Preparations?

Why it happens: Just saying 'add CuO until reaction stops'. How to avoid it: The word 'EXCESS' is a mark-scheme keyword. State 'add CuO IN EXCESS until no more dissolves' — this guarantees no acid remains. Source: 4CH1 Examiner Reports 2022-2024

Why is "Forgetting to WASH the precipitate / crystals" a common mistake in Acids, Bases and Salt Preparations?

Why it happens: Assuming the filter step removes all impurities. How to avoid it: After filtering, ALWAYS wash with COLD distilled water (cold so the salt doesn't redissolve) to remove soluble impurities (K⁺, NO₃⁻, etc.) from the surface of the solid. This is a separate marked step.

Why is "Trying to make an insoluble salt by acid + base + crystallisation" a common mistake in Acids, Bases and Salt Preparations?

Why it happens: Using the 'soluble salt' method as a default. How to avoid it: Check solubility FIRST. Insoluble (AgCl, BaSO₄, PbI₂) → MUST use precipitation by mixing two soluble solutions. Soluble salts → use one of acid + insoluble base / titration / metal / carbonate.

Why is "Saying 'all sulfates are soluble' or 'all chlorides are soluble'" a common mistake in Acids, Bases and Salt Preparations?

Why it happens: Memorising the general rule but forgetting exceptions. How to avoid it: Learn the exceptions: SULFATES insoluble: BaSO₄, PbSO₄ (CaSO₄ slightly). CHLORIDES insoluble: AgCl, PbCl₂. Same for Br⁻, I⁻ with Ag and Pb. CARBONATES mostly INSOLUBLE — only Group 1 + NH₄⁺ soluble. Source: 4CH1 Examiner Reports 2022-2024

Why is "Using 'add insoluble base + filter' method to make Na₂SO₄" a common mistake in Acids, Bases and Salt Preparations?

Why it happens: Forgetting that NaOH and Na₂CO₃ are soluble. How to avoid it: If both the acid AND the base are soluble (Na, K, NH₄⁺), there is no solid to filter off. Use TITRATION to measure the exact volume needed; then repeat without indicator.

Inorganic ChemistryPearson Edexcel IGCSE Chemistry 4CH1 Higher

Acids, Bases and Salt Preparations

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Acids, Bases and Salt Preparations — Pearson Edexcel International GCSE Chemistry 4CH1 Study Notes (2026 onwards, Spec Issue 4)

Selecting the right preparation method based on salt solubility. SOLUBLE SALT methods: acid + insoluble base + filter + crystallise (e.g. CuSO₄ from CuO + H₂SO₄); acid + alkali by titration (when both reactants soluble); acid + reactive metal; acid + carbonate. INSOLUBLE SALT method: precipitation by mixing two soluble solutions (e.g. PbI₂ from Pb(NO₃)₂ + KI). Edexcel solubility rules + the role of 'excess', filtration, slow cooling crystallisation, washing and drying. Hydrated salts and water of crystallisation.

At a glance

Step 1 — solubility check of target salt determines method.
Soluble salt + insoluble reactant available (CuO, ZnCO₃, Zn metal): acid + reactant in EXCESS → filter → crystallise.
Soluble salt + only soluble reactants (Na, K, NH₄): TITRATION — measure volume needed, repeat without indicator, crystallise.
Insoluble salt: PRECIPITATION — mix two soluble solutions → filter → wash → dry.
Solubility rules: Group 1 + NH₄⁺ + NO₃⁻ all soluble. Most chlorides soluble (except AgCl, PbCl₂). Most sulfates soluble (except BaSO₄, PbSO₄). Most carbonates and hydroxides insoluble (except Group 1, NH₄⁺).
Crystallisation: evaporate to SATURATED (not dry); cool slowly for large crystals; filter, wash, pat dry.
Water of crystallisation: chemically-bound water in crystal lattice (CuSO₄·5H₂O). Heat → loses water → anhydrous.

What you’ll learn

Mapped to the Cambridge IGCSE 4CH1 syllabus (2026 onwards).

2.46 — Describe the preparation of pure, dry samples of named soluble salts: from a soluble acid + an insoluble reactant (metal, base, carbonate).
2.47 — Describe the preparation of soluble salts where both reactants are soluble using a titration procedure (acid + alkali).
2.48 — Recall the general solubility rules for salts in water (Group 1, ammonium, nitrates → soluble; halide/sulfate exceptions; carbonates/hydroxides mostly insoluble).
2.49 — Describe the preparation of an insoluble salt by precipitation, including filtering, washing and drying.
2.50 — Understand 'water of crystallisation' and the formula notation X·nH₂O. Heating hydrated → anhydrous.

Choosing the preparation method (spec 2.46, 2.47, 2.49)

Solubility of target salt + availability of soluble/insoluble reactants determines method.

Salt preparation choice depends on TWO questions:

Question 1: Is the TARGET SALT soluble or insoluble?

Question 2: Are the REACTANTS soluble or insoluble?

Decision tree.

                       Is target salt soluble?
                                 |
              +------------------+------------------+
              | YES                                  | NO
              v                                      v
   Is there an INSOLUBLE                  PRECIPITATION
   reactant available?                    Mix two soluble
              |                            solutions whose
   +----------+----------+                 ions form
   | YES                  | NO            insoluble salt
   v                      v               → filter, wash, dry
ACID + INSOLUBLE       TITRATION
BASE/METAL/CARBONATE    (acid + alkali)
+ filter + crystallise

Choose the preparation route from the solubility of the target salt and whether an insoluble reactant exists.

Method 1: Acid + insoluble base / metal / carbonate (soluble salt + insoluble reactant).

Used for salts like CuSO₄, ZnCl₂, MgSO₄, FeCl₂. The base/metal/carbonate is INSOLUBLE in water (so excess can be filtered off).

Add reactant in EXCESS to warm acid; stir.
Stop when no more reactant dissolves.
Filter to remove unreacted excess.
Evaporate the filtrate to saturated solution.
Cool slowly for crystals; filter; wash with cold distilled water; pat dry.

Why excess? Guarantees ALL the acid has been neutralised — no acid contamination of the salt.

Method 2: Titration (soluble salt + two soluble reactants).

Used for sodium / potassium / ammonium salts (no insoluble reactant available because Group 1 and NH₄⁺ salts are all soluble).

Titrate acid against alkali to find the exact volume needed (using indicator).
Note the volume of acid required (the 'titre').
Repeat WITHOUT indicator: mix the same volume of acid and alkali in a clean flask (no indicator → no contamination of crystals).
Evaporate to saturated solution.
Cool slowly to crystallise; filter; wash; dry.

Method 3: Precipitation (insoluble salt).

Used for AgCl, BaSO₄, PbI₂.

Pick TWO SOLUBLE solutions whose ions combine to give the insoluble salt:
- Want AgCl? → Use AgNO₃(aq) + NaCl(aq) (both soluble).
- Want BaSO₄? → Use BaCl₂(aq) + Na₂SO₄(aq).
- Want PbI₂? → Use Pb(NO₃)₂(aq) + KI(aq).
Mix the solutions → insoluble salt precipitates immediately.
Filter to collect precipitate as residue.
Wash with distilled water (removes soluble byproducts like NaNO₃).
Dry in a warm oven or air-dry.

TARGET INSOLUBLE → PRECIPITATION (mix two soluble solutions).
Target SOLUBLE + insoluble reactant available → acid + base/metal/carbonate (excess) → filter → crystallise.
Target SOLUBLE + all soluble reactants (Na, K, NH₄) → TITRATION.
Crystallisation: evaporate to saturated; cool slowly; filter + wash + dry.

See the full worked example for acids, bases and salt preparations →

Preparing a soluble salt from an insoluble base (spec 2.46)

Classic 6-step method using CuSO₄ as worked example.

Worked procedure — CuSO₄·5H₂O from CuO + dilute H₂SO₄.

Equation: $\text{CuO(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{CuSO}_4\text{(aq)} + \text{H}_2\text{O(l)}$ .

Why this method? Cu is BELOW H — Cu metal doesn't react with dilute H₂SO₄ (no fizz). CuO is an INSOLUBLE BASE — can be filtered off. CuSO₄ is SOLUBLE — can crystallise from solution.

Step-by-step.

Measure ~50 cm³ of dilute H₂SO₄ into a 100 cm³ beaker. Warm gently with a Bunsen burner + tripod + gauze (or on a hot-water bath) to ~50 °C. Warming speeds the reaction. Do NOT boil.
Add black CuO powder a spatula at a time, stirring after each. The CuO will gradually dissolve, giving a blue solution of CuSO₄. Continue adding until NO MORE CuO DISSOLVES — this is the visual sign that ALL the acid has reacted. Adding CuO IN EXCESS guarantees no leftover acid (otherwise it would crystallise with the salt).
Filter to remove excess CuO. Set up a filter funnel with filter paper in a clean conical flask (or evaporating basin). Pour the warm mixture through. Unreacted CuO is RESIDUE on the filter paper; CuSO₄ solution is the FILTRATE.
Evaporate the filtrate to a saturated solution. Heat gently in an evaporating basin on a water bath (or low Bunsen flame). Watch for the solution becoming thicker / cloudier. Test for saturation: dip a clean glass rod into the solution; if blue crystals form on the rod as it dries, the solution is saturated. STOP heating here — DO NOT evaporate to dryness (would lose water of crystallisation, decompose to anhydrous CuSO₄).
Cool slowly to crystallise. Leave to cool to room temperature undisturbed (ideally overnight). Slow cooling = large, well-formed CuSO₄·5H₂O crystals (blue). The solubility decreases as the solution cools, so crystals come out of solution.
Filter, wash, dry. Filter the crystals from the remaining 'mother liquor'. Wash with a SMALL volume of COLD distilled water (cold so crystals don't redissolve; small to minimise yield loss; distilled to avoid contamination). Pat dry between two sheets of filter paper.

Final product: pure, dry, blue CuSO₄·5H₂O crystals (hydrated copper(II) sulfate pentahydrate).

Excess base then filter, evaporate to saturation, cool slowly, then filter, wash and dry — never evaporate to dryness.

Variations.

Use ZnO + H₂SO₄ → ZnSO₄·7H₂O (white crystals).
Use MgO + 2HCl → MgCl₂·6H₂O (colourless).
Use FeO + H₂SO₄ → FeSO₄·7H₂O (pale green).
Replace insoluble base with insoluble carbonate (e.g. ZnCO₃ + H₂SO₄ → ZnSO₄ + H₂O + CO₂) — same method.
Replace insoluble base with reactive metal (e.g. Zn + H₂SO₄ → ZnSO₄ + H₂) — note H₂ gas.

Equation: insoluble base + acid → soluble salt + water.
Add base IN EXCESS until no more dissolves.
Filter, evaporate to saturated, cool slowly, filter, wash, dry.
Never evaporate to dryness — lose water of crystallisation.

See the full worked example for acids, bases and salt preparations →

Preparing an insoluble salt by precipitation (spec 2.49)

Mix two soluble solutions → precipitate insoluble salt → filter, wash, dry.

Insoluble salts (AgCl, BaSO₄, PbI₂, PbCl₂) cannot be made by 'acid + base + crystallise' methods. Instead, they are made by PRECIPITATION — mixing two soluble solutions whose ions combine to form the insoluble salt.

Worked example — PbI₂ ('golden rain').

Equation: $\text{Pb(NO}_3)_2\text{(aq)} + 2\text{KI(aq)} \rightarrow \text{PbI}_2\text{(s)} + 2\text{KNO}_3\text{(aq)}$ .

Mixing two soluble solutions precipitates the insoluble salt, which is then filtered, washed and dried.

Choose the reactants:

Need Pb²⁺: use Pb(NO₃)₂ (all nitrates soluble). NOT PbCl₂ or PbSO₄ — those are insoluble.
Need I⁻: use KI (all K salts soluble). NaI also OK; AgI is NOT (AgI insoluble).
The other ions (K⁺, NO₃⁻) must form a SOLUBLE compound (KNO₃) so they wash away in the filtrate.

Step-by-step.

Mix equal volumes of Pb(NO₃)₂(aq) and KI(aq) in a beaker. A bright yellow precipitate of PbI₂ forms immediately — described as 'golden rain' because of the shimmering appearance.
Filter through filter paper. The yellow PbI₂ stays as RESIDUE on the filter; the colourless KNO₃ solution passes through as FILTRATE (discarded).
Wash the residue ON the filter paper with several portions of distilled water. This removes any soluble impurities (K⁺, NO₃⁻, excess Pb²⁺ or I⁻ from one reactant being in slight excess). Doesn't dissolve PbI₂ — it's insoluble.
Dry. Transfer to a watch glass and either dry in a warm oven or leave at room temperature.

Final product: pure, dry, yellow PbI₂ powder.

Other precipitation examples.

Salt	Soluble reactants	Other product (soluble)
AgCl (white)	AgNO₃ + NaCl	NaNO₃
AgBr (cream)	AgNO₃ + NaBr	NaNO₃
AgI (yellow)	AgNO₃ + NaI	NaNO₃
BaSO₄ (white)	BaCl₂ + Na₂SO₄	NaCl
PbSO₄ (white)	Pb(NO₃)₂ + Na₂SO₄	NaNO₃
PbCl₂ (white)	Pb(NO₃)₂ + NaCl	NaNO₃

Ionic equation (omits spectator ions).

For PbI₂: $\text{Pb}^{2+}\text{(aq)} + 2\text{I}^-\text{(aq)} \rightarrow \text{PbI}_2\text{(s)}$ .

The K⁺ and NO₃⁻ are spectator ions — they appear on both sides of the full equation and don't actually take part in the precipitate-forming step.

Insoluble salt → make by precipitation (NOT crystallisation).
Choose two SOLUBLE reactants providing the ions you need.
Mix → filter → WASH with distilled water → dry.
Ionic equation: Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s).

See the full worked example for acids, bases and salt preparations →

Solubility rules (spec 2.48)

Memorise: Group 1, NH₄⁺, NO₃⁻ always soluble. Sulfate / halide exceptions.

Edexcel-required solubility rules — memorise these for both calculation and method-choice questions.

Always SOLUBLE.

Group of salts	Notes
All Group 1 (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺)	No exceptions
All ammonium (NH₄⁺)	No exceptions
All nitrates (NO₃⁻)	No exceptions — including AgNO₃, Pb(NO₃)₂, Cu(NO₃)₂, Fe(NO₃)₃
All ethanoates (CH₃COO⁻)	(Not always required by 4CH1 but useful)

MOSTLY soluble, with exceptions:

Salt type	Soluble	INSOLUBLE exceptions
Chlorides (Cl⁻)	Most metal chlorides	AgCl (white), PbCl₂ (white, soluble in HOT water)
Bromides (Br⁻)	Most metal bromides	AgBr (cream), PbBr₂ (white)
Iodides (I⁻)	Most metal iodides	AgI (yellow), PbI₂ (yellow)
Sulfates (SO₄²⁻)	Most metal sulfates	BaSO₄ (white), PbSO₄ (white), CaSO₄ (slightly)

MOSTLY insoluble, with exceptions:

Salt type	Insoluble	SOLUBLE exceptions
Carbonates (CO₃²⁻)	Most metal carbonates	Group 1 (Li₂CO₃, Na₂CO₃, K₂CO₃) + (NH₄)₂CO₃
Hydroxides (OH⁻)	Most metal hydroxides	Group 1 (LiOH, NaOH, KOH) + Ca(OH)₂ slightly + NH₃ solution
Sulfides (S²⁻)	Most metal sulfides	Group 1 + Group 2 + (NH₄)₂S

Practical use in method choice.

Want CuSO₄ → SOLUBLE (most sulfates). Use insoluble-base method.
Want NaCl → SOLUBLE (Group 1). Use titration method.
Want AgCl → INSOLUBLE (exception). Use precipitation.
Want BaSO₄ → INSOLUBLE (exception). Use precipitation.
Want PbI₂ → INSOLUBLE (exception). Use precipitation.

Memory tricks.

'SAN-N' — Sodium, Ammonium, Nitrate (always soluble) + 'and Nitrate again' to emphasise.
'AgCl, AgBr, AgI' — all silver halides insoluble (the basis for the halide test).
'BaSO₄' — barium sulfate insoluble (the basis for the sulfate test).
Carbonates 'all insoluble except K, Na, NH₄' (heat them → metal oxide + CO₂; but Group 1 carbonates are too stable to decompose easily).

Group 1, NH₄⁺, NO₃⁻ → ALL soluble.
Sulfates soluble EXCEPT BaSO₄, PbSO₄ (CaSO₄ slightly).
Halides soluble EXCEPT AgX and PbX₂.
Carbonates and hydroxides mostly INSOLUBLE; only Group 1 + NH₄⁺ soluble.

See the full worked example for acids, bases and salt preparations →

Hydrated salts and water of crystallisation (spec 2.50)

X·nH₂O. Heat → loses water → anhydrous. Reverse: add water → back to blue.

Water of crystallisation = water molecules chemically bound into the crystal lattice of a salt. They form a regular part of the crystal structure — NOT 'free' water trapped in cracks, but specific H₂O units coordinated to the metal cation.

Formula notation. Write the salt formula, then a dot, then the number of H₂O molecules: $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ means each formula unit of CuSO₄ has 5 water molecules associated with it.

Common hydrated salts.

Salt	Formula	Colour	Common name
Copper(II) sulfate pentahydrate	CuSO₄·5H₂O	Bright blue	'Blue vitriol'
Magnesium sulfate heptahydrate	MgSO₄·7H₂O	Colourless	Epsom salt
Sodium carbonate decahydrate	Na₂CO₃·10H₂O	White	Washing soda
Calcium sulfate dihydrate	CaSO₄·2H₂O	White	Gypsum
Iron(II) sulfate heptahydrate	FeSO₄·7H₂O	Pale green	Green vitriol
Cobalt(II) chloride hexahydrate	CoCl₂·6H₂O	Pink	(Cobalt chloride paper)

Heating drives off water — hydrated → anhydrous.

When a hydrated salt is heated, the water of crystallisation is driven off as steam, leaving the anhydrous (water-free) salt:

$\text{CuSO}_4 \cdot 5\text{H}_2\text{O(s)} \xrightarrow{\Delta} \text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O(g)}$

Colour change observed: blue → white (for CuSO₄). Condensation forms on cold parts of the apparatus (the steam).

Reverse — adding water reforms the hydrate.

If water is added to anhydrous (white) CuSO₄, the blue colour returns and heat is given out (exothermic):

$\text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O(l)} \rightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O(s)} + \text{heat}$

Tests using hydrated/anhydrous salts.

Anhydrous CuSO₄ test for water: white CuSO₄ turns BLUE on adding water. Used to detect water in unknown liquids.
Cobalt(II) chloride paper test for water: blue CoCl₂ paper turns PINK on adding water.

Both confirm the PRESENCE of water but NOT purity (a salt solution would also turn the paper pink).

Why slow cooling matters for crystallisation. Hydrated salts crystallise out of solution with their full water of crystallisation when given time. Fast cooling gives small, irregular crystals; slow cooling gives large, well-formed crystals with the correct hydration state.

Water of crystallisation = chemically bonded water in crystal lattice (CuSO₄·5H₂O).
Heat → drives off water → anhydrous salt (blue → white for CuSO₄).
Add water → reforms hydrate (white → blue, exothermic).
Test for water: anhydrous CuSO₄ → blue OR CoCl₂ paper → pink.

See the full worked example for acids, bases and salt preparations →

How it’s examined

Salt preparations are a 4CH1 STAPLE — appear on EVERY paper. Paper 1 (4CH1/1C) items: (a) describe a method to prepare a named salt (5-6 marks — required practical); (b) state solubility rules and predict whether a salt is soluble or insoluble (3-4 marks); (c) write balanced equation + state observations (3-4 marks). Paper 2 (4CH1/2C) BOLD 'P' items: (d) choose appropriate method given starting materials + solubility of target (4-5 marks); (e) describe precipitation reactions including ionic equations (3-4 marks). Watch the words 'in EXCESS' and 'COOL SLOWLY' — these are mark-scheme keywords.

Sources: Pearson Edexcel International GCSE Chemistry 4CH1 specification — Issue 4, September 2024 (valid for 2026 onwards); Pearson Edexcel 4CH1 Examiner Reports 2022-2024; 4CH1/1C May/Jun 2024 question paper and mark scheme; 4CH1/1C January 2024 question paper and mark scheme; 4CH1/2C May/Jun 2024 question paper and mark scheme; 4CH1/2C January 2024 question paper and mark scheme. Last reviewed 2026-05-12.

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Step-by-step worked examples — Acids, Bases and Salt Preparations

Step-by-step solutions to past-paper-style questions on acids, bases and salt preparations, written exactly the way a tutor would explain them at the board.

1Preparing pure dry copper(II) sulfate crystals from CuO + H₂SO₄ (6 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q4 (required practical)• soluble salt, insoluble base, spec-2.46, spec-2.47
Question
Describe how to prepare a pure, dry sample of copper(II) sulfate crystals starting from copper(II) oxide (black powder) and dilute sulfuric acid. (6 marks)
Step-by-step solution
1. Step 1
  Choose method (1 M). Copper is BELOW H in the reactivity series → Cu metal doesn't react with dilute H₂SO₄. CuO is an INSOLUBLE BASE → use the 'acid + insoluble base + filter + crystallise' method.
2. Step 2
  Add CuO in EXCESS to warm acid (1 A). Warm ~25 cm³ of dilute H₂SO₄ in a beaker (Bunsen + gauze) — gentle warming, NOT boiling. Add black CuO powder a SPATULA at a time, stirring after each. The CuO dissolves to give a blue solution: $\text{CuO} + \text{H}_2\text{SO}_4 \rightarrow \text{CuSO}_4 + \text{H}_2\text{O}$ . Continue adding until no more CuO dissolves — this guarantees the acid is fully neutralised (EXCESS base ensures all acid reacted).
3. Step 3
  Filter to remove unreacted CuO (1 A). Filter the mixture through a filter paper in a filter funnel into a clean conical flask / evaporating basin. The black CuO solid stays in the filter paper (RESIDUE); the blue CuSO₄ solution passes through (FILTRATE). This step is essential — any unreacted base would contaminate the salt.
4. Step 4
  Evaporate to concentrate (1 B). Heat the filtrate in an evaporating basin GENTLY (above a water bath OR with a low Bunsen flame) until ~half the volume has evaporated — the solution becomes SATURATED. Test for saturation: dip a glass rod and lift — crystals form on the rod. STOP heating before all water evaporates (otherwise the crystals decompose / lose water of crystallisation).
5. Step 5
  Crystallise (1 B). Leave the saturated solution to cool slowly to room temperature (overnight is best). As temperature falls, solubility decreases → blue CuSO₄·5H₂O crystals form. Slow cooling = large, well-formed crystals.
6. Step 6
  Filter, wash, dry (1 B). Filter the crystals from the remaining solution (the 'mother liquor'). Wash with a SMALL volume of cold distilled water (to remove impurities without redissolving the crystals). Pat dry between two pieces of filter paper. Final product: pure, dry, blue CuSO₄·5H₂O crystals.
Answer
Method: (1) Warm dilute H₂SO₄ in a beaker. (2) Add CuO in EXCESS, stirring, until no more dissolves. (3) Filter to remove unreacted CuO. (4) Evaporate the blue filtrate until saturated (crystals form on a glass rod). (5) Cool slowly to room temperature → CuSO₄·5H₂O crystals form. (6) Filter, wash with cold distilled water, pat dry between filter papers. Equation: CuO + H₂SO₄ → CuSO₄ + H₂O.
Examiner tip
Edexcel 6-mark mark scheme: 1 each for (a) warm acid + add base in excess; (b) ensure no more dissolves / acid fully reacted; (c) filter to remove excess base; (d) evaporate to saturated solution; (e) cool slowly to crystallise; (f) filter + wash + dry. Common error: 'boil to dryness' — this loses water of crystallisation and decomposes the salt. ALWAYS cool slowly for crystallisation.
2Preparing insoluble lead(II) iodide by precipitation (5 marks, Paper 1)
Core• Adapted from 4CH1/1C Jan 2024 Q3• insoluble salt, precipitation, spec-2.49, spec-2.50
Question
Lead(II) iodide is an insoluble yellow salt. Describe how to prepare a pure, dry sample of lead(II) iodide starting from solutions of lead(II) nitrate and potassium iodide. (5 marks)
Step-by-step solution
1. Step 1
  Choose two SOLUBLE reactants (1 M). Use lead(II) nitrate Pb(NO₃)₂ (all nitrates soluble) and potassium iodide KI (all K salts soluble). Both contain the ions we need (Pb²⁺ and I⁻) AND the by-products (K⁺, NO₃⁻) form a soluble compound (KNO₃) so they wash away.
  $\text{Pb(NO}_3)_2\text{(aq)} + 2\text{KI(aq)} \rightarrow \text{PbI}_2\text{(s)} + 2\text{KNO}_3\text{(aq)}$
2. Step 2
  Mix the solutions (1 A). Mix equal volumes of dilute Pb(NO₃)₂(aq) and KI(aq) in a beaker. An immediate YELLOW PRECIPITATE of PbI₂ forms — sometimes called the 'golden rain' reaction because of the characteristic shimmering appearance.
3. Step 3
  Filter (1 A). Filter the mixture through a filter paper. The yellow PbI₂ stays as RESIDUE on the filter paper; the colourless KNO₃ solution passes through as FILTRATE and is discarded.
4. Step 4
  Wash (1 B). Wash the precipitate ON THE FILTER PAPER with several portions of distilled water. This rinses off any soluble impurities (K⁺, NO₃⁻, excess Pb²⁺ or I⁻) without dissolving the PbI₂ (it's insoluble).
5. Step 5
  Dry (1 B). Transfer the washed precipitate to a watch glass and dry in a warm oven (or simply leave to dry at room temperature). Final product: pure, dry, yellow PbI₂ powder.
Answer
Method: (1) Mix Pb(NO₃)₂(aq) and KI(aq) — yellow PbI₂ precipitate forms immediately. (2) Filter through filter paper — PbI₂ stays as residue. (3) Wash the residue with distilled water to remove soluble KNO₃, excess ions. (4) Dry in a warm oven (or air-dry). Equation: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq).
Examiner tip
Edexcel 5-mark mark scheme: 1 for choosing two SOLUBLE reactants providing Pb²⁺ and I⁻; 1 for balanced equation (with state symbols); 1 for filter to collect precipitate; 1 for WASH with distilled water; 1 for dry. Both reactants MUST be soluble (the whole point of precipitation method). Common error: forgetting the 'wash' step → impurities remain.
3Applying solubility rules to predict precipitates (4 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q3• solubility rules, spec-2.48
Question
State whether each of the following salts is SOLUBLE or INSOLUBLE in water, giving the solubility rule used: (a) sodium carbonate, (b) barium sulfate, (c) silver chloride, (d) ammonium nitrate. (4 marks)
Step-by-step solution
1. Step 1
  (a) Sodium carbonate, Na₂CO₃ (1 M). SOLUBLE. Rule: most carbonates are INSOLUBLE, EXCEPT those of Group 1 metals (Na, K, Li) and ammonium. Sodium is a Group 1 metal, so Na₂CO₃ is soluble.
2. Step 2
  (b) Barium sulfate, BaSO₄ (1 A). INSOLUBLE. Rule: most sulfates are soluble, EXCEPT BaSO₄, PbSO₄ (and CaSO₄ is slightly soluble). Barium sulfate is the classic 'insoluble sulfate' — it's the basis for the chemical test for sulfate ions (white precipitate with Ba²⁺).
3. Step 3
  (c) Silver chloride, AgCl (1 A). INSOLUBLE. Rule: most chlorides are soluble, EXCEPT AgCl and PbCl₂. Silver chloride is white and is the basis for the test for chloride ions (white precipitate with AgNO₃).
4. Step 4
  (d) Ammonium nitrate, NH₄NO₃ (1 B). SOLUBLE. TWO rules apply: (i) all ammonium (NH₄⁺) salts are soluble; (ii) all nitrate (NO₃⁻) salts are soluble. Either rule alone is enough — both apply here.
Answer
(a) Na₂CO₃: SOLUBLE (Group 1 carbonate). (b) BaSO₄: INSOLUBLE (exception to soluble-sulfates rule). (c) AgCl: INSOLUBLE (exception to soluble-chlorides rule). (d) NH₄NO₃: SOLUBLE (ammonium AND nitrate — both always soluble).
Examiner tip
Edexcel 4-mark mark scheme: 1 mark per correct answer + rule. State the RULE not just the answer. Common errors: stating BaSO₄ soluble (because most sulfates are) — note the exception. AgCl is the only common silver salt that is insoluble — AgNO₃ is soluble.
4Choosing the correct salt-preparation method for ZnSO₄ from given starting materials (5 marks, Paper 2)
Extended• Adapted from 4CH1/2C Jan 2024 Q5• method choice, soluble salt, reactive metal, spec-2.46, spec-2.47
Question
A student wants to prepare a pure, dry sample of zinc sulfate (ZnSO₄·7H₂O). The available starting materials are: zinc metal granules, zinc oxide powder, zinc carbonate powder, dilute sulfuric acid. (a) Identify TWO different acid + reactant combinations that would work. (b) Choose the BEST method (with reason) and write the balanced equation. (5 marks)
Step-by-step solution
1. Step 1
  Solubility check (1 M). ZnSO₄ is soluble (most sulfates soluble; Zn is not an exception). So we need an 'acid + insoluble base' or 'acid + reactive metal' or 'acid + metal carbonate' method (NOT precipitation).
2. Step 2
  (a) Combination 1: Zn metal + H₂SO₄ (1 A). Zinc is ABOVE H in the reactivity series, so it reacts with acid to give H₂ gas: $\text{Zn} + \text{H}_2\text{SO}_4 \rightarrow \text{ZnSO}_4 + \text{H}_2$ . Add Zn granules in excess to warm dilute H₂SO₄ until fizzing stops; filter; crystallise.
3. Step 3
  (a) Combination 2: ZnO + H₂SO₄ (1 A). Zinc oxide is an insoluble base. $\text{ZnO} + \text{H}_2\text{SO}_4 \rightarrow \text{ZnSO}_4 + \text{H}_2\text{O}$ . Add ZnO in excess to warm acid; filter; crystallise. (ZnCO₃ + H₂SO₄ → ZnSO₄ + H₂O + CO₂ also works.)
4. Step 4
  (b) Best method (1 B). Using ZnO + H₂SO₄ (or ZnCO₃) is preferred over zinc metal because: (i) Zn metal + acid gives off H₂ gas (slight loss of mass / fire risk); (ii) the reaction with ZnO is more controllable; (iii) it's easier to tell when all the acid has reacted (no more solid dissolves) than to know when to stop adding metal. Some textbooks prefer Zn metal for simplicity, but ZnO/ZnCO₃ are safer.
5. Step 5
  (b) Method with equation (1 B). Warm dilute H₂SO₄, add ZnO in EXCESS while stirring until no more dissolves, filter off excess ZnO, evaporate to saturated solution, cool slowly to crystallise ZnSO₄·7H₂O, filter, wash with cold water, pat dry. Equation: $\text{ZnO(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{H}_2\text{O(l)}$ .
Answer
(a) Two valid combinations: (i) Zn + H₂SO₄ → ZnSO₄ + H₂; (ii) ZnO + H₂SO₄ → ZnSO₄ + H₂O (or ZnCO₃ + H₂SO₄). (b) BEST: ZnO + acid (insoluble-base method) — controllable, no flammable H₂ gas, easy to know when complete. Method as above. The crystals are ZnSO₄·7H₂O (water of crystallisation).
Examiner tip
Edexcel 5-mark mark scheme: 1 for solubility identification of ZnSO₄; 1 each for two valid combinations; 1 for choosing best method WITH REASON; 1 for full procedure summary. Either ZnO or ZnCO₃ is acceptable as best. Zn metal works too but examiners reward the safety argument.

Model Answers — Acids, Bases and Salt Preparations

High-scoring sample answers for acids, bases and salt preparations on the Pearson Edexcel IGCSE 4CH1 paper, with examiner-style notes mapping each response to the mark scheme and assessment objectives.

Question 1
4CH1/1C-style — required-practical salt preparation6 marks
Describe in detail the procedure to prepare a pure, dry sample of hydrated copper(II) sulfate crystals (CuSO₄·5H₂O) starting from copper(II) oxide and dilute sulfuric acid. Include a balanced equation, the role of each step, and one safety precaution. (6 marks)
Model answer
Equation. $\text{CuO(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{CuSO}_4\text{(aq)} + \text{H}_2\text{O(l)}$ .

Why this method? Copper is BELOW hydrogen in the reactivity series, so copper metal does NOT react with dilute sulfuric acid. CuO is an INSOLUBLE BASE, so the 'acid + insoluble base in excess + filter + crystallise' method is used.

Step-by-step procedure.

Warm the acid. Measure ~50 cm³ of dilute H₂SO₄ into a 100 cm³ beaker. Place on a tripod + gauze and warm gently with a Bunsen flame (or on a hot water bath) to ~50 °C — do NOT boil. Warming speeds the reaction.

Add CuO in excess. Add black CuO powder a spatula at a time, stirring after each addition. The CuO will react with the acid to give a blue solution of CuSO₄. Keep adding CuO until no more dissolves — this is the visual cue that ALL the acid has been neutralised. Using EXCESS base guarantees no acid remains in the final product (which would contaminate the crystals).

Filter to remove excess CuO. Set up a filter funnel with filter paper over a clean conical flask or evaporating basin. Pour the warm mixture through. The unreacted black CuO is the RESIDUE (caught on filter paper); the blue copper(II) sulfate solution is the FILTRATE.

Evaporate to concentrate. Heat the filtrate in an evaporating basin GENTLY (on a water bath OR with a low Bunsen flame). Stop when the solution is saturated — test by dipping a cold glass rod: if crystals form on the rod, the solution is saturated. Do NOT evaporate to dryness — this would (a) lose the water of crystallisation, (b) potentially decompose the salt with overheating.

Cool slowly to crystallise. Leave the saturated solution to cool to room temperature, ideally overnight. As the temperature falls, the solubility of CuSO₄ decreases and large, well-formed blue crystals of CuSO₄·5H₂O (hydrated copper(II) sulfate pentahydrate) form. Slow cooling = large crystals.

Filter, wash, dry. Filter off the crystals from the mother liquor. Wash with a SMALL volume of COLD distilled water (cold so the crystals don't redissolve; minimum volume so impurities are removed but yield is preserved). Pat the crystals dry between two sheets of filter paper.

Result. Pure, dry, blue CuSO₄·5H₂O crystals. The blue colour comes from the [Cu(H₂O)₆]²⁺ complex ion — if you heated them strongly they would lose water and turn white (anhydrous CuSO₄).

Safety precaution. Wear safety goggles throughout (acid can damage eyes); use a heat-resistant mat under the Bunsen; allow hot apparatus to cool before handling.
Why this scores
Edexcel 6-mark mark scheme: 1 for balanced equation; 1 for adding base in excess until no more dissolves; 1 for filter to remove excess; 1 for evaporate to saturated (not to dryness); 1 for cool slowly to crystallise; 1 for wash + dry + safety. State 'EXCESS' base — this is the key word ensuring no acid remains. Don't say 'boil dry' — loses water of crystallisation.

Question 2

4CH1/2C-style — method choice based on solubility6 marks

Three different salts are to be prepared: zinc chloride (ZnCl₂), silver chloride (AgCl), and sodium sulfate (Na₂SO₄). For EACH salt, name the most appropriate preparation method and explain your choice. (6 marks)

Model answer

Reference solubility rules.

All Group 1 (Na, K, Li) and ammonium salts → SOLUBLE.
All nitrates → SOLUBLE.
Most chlorides → soluble; EXCEPT AgCl, PbCl₂.
Most sulfates → soluble; EXCEPT BaSO₄, PbSO₄.
Most carbonates → INSOLUBLE; EXCEPT Group 1 + NH₄⁺.

(1) Zinc chloride (ZnCl₂) — soluble salt.

Method: acid + insoluble base (ZnO) OR acid + reactive metal (Zn) OR acid + carbonate (ZnCO₃).
Best: ZnO (or ZnCO₃) + dilute HCl → ZnCl₂ + H₂O (or ZnCl₂ + H₂O + CO₂).
Reasoning: ZnCl₂ is soluble, so we can crystallise it from solution. Zn is reactive enough to give a soluble salt by neutralising HCl. The 'add excess + filter + crystallise' method gives high purity.
Equation: $\text{ZnO} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2\text{O}$ .

(2) Silver chloride (AgCl) — INSOLUBLE salt.

Method: PRECIPITATION by mixing two soluble solutions.
Reactants: silver nitrate AgNO₃(aq) (soluble — all Ag NO₃ salts are) + sodium chloride NaCl(aq) (soluble — all Na salts).
Reasoning: AgCl is INSOLUBLE, so it precipitates immediately on mixing. The by-product NaNO₃ is soluble and washes away. We can't make AgCl by acid + base because AgCl wouldn't crystallise from solution — it would just precipitate at random concentration.
Equation: $\text{AgNO}_3\text{(aq)} + \text{NaCl(aq)} \rightarrow \text{AgCl(s)} + \text{NaNO}_3\text{(aq)}$ .
Procedure: Mix solutions → white precipitate of AgCl → filter → wash with distilled water → dry.

(3) Sodium sulfate (Na₂SO₄) — soluble salt, but starting reactants are alkali + acid.

Method: TITRATION (acid + alkali) followed by crystallisation. Na is Group 1, so any sodium base (NaOH) is SOLUBLE — there's no insoluble Na compound to filter off, so we can't use the 'add excess + filter' approach.
Reactants: NaOH(aq) + H₂SO₄(aq) → Na₂SO₄(aq) + H₂O.
Reasoning: We must measure the EXACT volume of acid needed to react with a known volume of alkali (or vice versa) using indicator + titration. Then we know how much acid to add WITHOUT indicator, so the indicator doesn't contaminate the crystals.
Procedure: (1) Titration with indicator to find exact volume; (2) repeat WITHOUT indicator using that exact volume; (3) evaporate to saturated solution; (4) crystallise; (5) filter + dry.
Equation: $2\text{NaOH} + \text{H}_2\text{SO}_4 \rightarrow \text{Na}_2\text{SO}_4 + 2\text{H}_2\text{O}$ .

Summary table.

Salt	Solubility	Method	Why
ZnCl₂	Soluble	Acid + insoluble base	ZnO insoluble → can filter excess
AgCl	Insoluble	Precipitation	Insoluble → precipitates on mixing
Na₂SO₄	Soluble	Titration	Both reactants soluble → no solid to filter

Why this scores

Edexcel 6-mark mark scheme: 2 for each salt (1 method, 1 reasoning). The KEY decision tree: (i) Is the target salt SOLUBLE or INSOLUBLE? If insoluble → precipitation. (ii) If soluble, is there an INSOLUBLE reactant available? If yes → acid + insoluble base/carbonate/metal. (iii) If only soluble reactants → titration.

Question 3
4CH1/1C-style — solubility-rules recall5 marks
State the solubility rules for the following types of salts in water, including any common exceptions: (a) Group 1 metal salts and ammonium salts; (b) nitrates; (c) chlorides; (d) sulfates; (e) carbonates and hydroxides. (5 marks)
Model answer
(a) Group 1 metals and ammonium salts. ALL SOLUBLE without exception. This includes lithium (Li⁺), sodium (Na⁺), potassium (K⁺), rubidium (Rb⁺), caesium (Cs⁺) salts AND ammonium (NH₄⁺) salts. Examples: NaCl, KOH, Na₂CO₃, (NH₄)₂SO₄. No exceptions.

(b) Nitrates (NO₃⁻). ALL SOLUBLE without exception. Every metal nitrate dissolves in water — from NaNO₃ to AgNO₃ to Pb(NO₃)₂ to Fe(NO₃)₃. This is why nitrates are used as the 'soluble form' of a metal cation for precipitation reactions.

(c) Chlorides (Cl⁻). Most chlorides are SOLUBLE, EXCEPT:

Silver chloride (AgCl) — white precipitate (used as a test for Cl⁻).

Lead(II) chloride (PbCl₂) — white precipitate (soluble in HOT water; reprecipitates on cooling).

The same applies to bromides (AgBr cream, PbBr₂ white) and iodides (AgI yellow, PbI₂ yellow).

(d) Sulfates (SO₄²⁻). Most sulfates are SOLUBLE, EXCEPT:

Barium sulfate (BaSO₄) — white precipitate (used as a test for SO₄²⁻).

Lead(II) sulfate (PbSO₄) — white precipitate.

Calcium sulfate (CaSO₄) — SLIGHTLY SOLUBLE (treated as sparingly soluble; can form a precipitate at high concentrations).

(e) Carbonates (CO₃²⁻) and hydroxides (OH⁻). Most are INSOLUBLE, EXCEPT:

Carbonates: Group 1 (Li₂CO₃, Na₂CO₃, K₂CO₃) AND ammonium carbonate (NH₄)₂CO₃ are SOLUBLE.

Hydroxides: Group 1 (LiOH, NaOH, KOH) are SOLUBLE; Ca(OH)₂ is SLIGHTLY SOLUBLE (gives limewater); ammonia solution (NH₃ + H₂O) gives NH₄OH which is soluble. All others (Cu(OH)₂, Fe(OH)₂, Fe(OH)₃, Mg(OH)₂, Al(OH)₃, etc.) are INSOLUBLE.

Memory tip — 'SAN P' rule. Soluble: all Sodium, Ammonium, Nitrate. Plus: most chlorides + most sulfates (with the listed exceptions). The pattern most often examined: 'insoluble metal hydroxides' (the basis of the NaOH ion test in Chemical Tests).
Why this scores
Edexcel 5-mark mark scheme: 1 mark each for (a), (b), (c), (d), (e). The exceptions are essential — examiners specifically target them. AgCl, PbI₂, BaSO₄ are the three most-tested insoluble salts in 4CH1 papers.
Question 4
4CH1/2C-style — water of crystallisation5 marks
Hydrated copper(II) sulfate has the formula CuSO₄·5H₂O. (a) Explain what is meant by 'water of crystallisation'. (b) Describe the colour change observed when hydrated CuSO₄ is heated strongly, and write a word equation for the change. (c) State how the original colour can be restored. (5 marks)
Model answer
(a) Water of crystallisation. The water molecules that are CHEMICALLY BONDED into a salt's crystal lattice when it crystallises from aqueous solution. They form part of the regular crystalline structure — they are NOT 'free' water trapped in the crystal, but specific H₂O units coordinated to the metal cation. They are written as a 'dot' formula: $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ means each formula unit of CuSO₄ has 5 water molecules associated with it. Many salts form hydrates: MgSO₄·7H₂O (Epsom salt), Na₂CO₃·10H₂O (washing soda), CaSO₄·2H₂O (gypsum), CuSO₄·5H₂O (blue vitriol), FeSO₄·7H₂O (green vitriol).

(b) Colour change on heating. When hydrated copper(II) sulfate (CuSO₄·5H₂O, BLUE) is heated strongly, the water of crystallisation is driven off as steam. The remaining anhydrous copper(II) sulfate (CuSO₄, WHITE / pale grey) is left.

Word equation:

$\text{hydrated copper(II) sulfate} \xrightarrow{\Delta} \text{anhydrous copper(II) sulfate} + \text{water}$

Symbol equation:

$\text{CuSO}_4 \cdot 5\text{H}_2\text{O(s)} \xrightarrow{\Delta} \text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O(g)}$

Observations: blue powder turns white; condensation (steam) forms on the cool upper part of the test tube — droplets of colourless liquid. (The anhydrous salt is the basis of the chemical test for WATER: white CuSO₄ turns BLUE when water is added.)

(c) Restoring the colour. Add a few drops of water to the white anhydrous CuSO₄ — it turns BLUE again as the water rehydrates the salt:

$\text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O(l)} \rightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O(s)} + \text{heat (exothermic)}$

This colour change (white → blue on adding water) is the standard 4CH1 test for the presence of water.
Why this scores
Edexcel 5-mark mark scheme: 1 for definition of water of crystallisation; 1 for blue → white colour change; 1 for word equation; 1 for adding water; 1 for blue colour restored on adding water (rehydration). The 'water of crystallisation' definition often loses marks — must state CHEMICALLY BONDED / part of crystal structure, not just 'water in the salt'.

Key Definitions and Keywords — Acids, Bases and Salt Preparations

Definitions to memorise and the exact keywords mark schemes credit for acids, bases and salt preparations answers — sharpened from recent examiner reports for the 2026 Pearson Edexcel IGCSE 4CH1 sitting.

Salt
Examiner keyword
An ionic compound formed when the H⁺ ions of an acid are replaced by a metal cation (or ammonium NH₄⁺). The anion comes from the acid (e.g. HCl → chloride; H₂SO₄ → sulfate; HNO₃ → nitrate).
Soluble salt
Examiner keyword
A salt that dissolves in water (e.g. NaCl, CuSO₄, KNO₃, ZnSO₄). Prepared by: acid + insoluble base, acid + alkali (titration), acid + metal, acid + carbonate.
Insoluble salt
Examiner keyword
A salt that does NOT dissolve in water (e.g. AgCl, BaSO₄, PbI₂). Prepared by PRECIPITATION — mixing two soluble solutions whose ions combine to give the insoluble salt.
Precipitation reaction
Examiner keyword
A reaction between two solutions producing an INSOLUBLE solid (the precipitate). e.g. Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s). The starting solutions are both soluble; the product is insoluble.
Precipitate
An insoluble solid formed in solution during a precipitation reaction; appears as a cloudiness or fine suspension. Filtered off, washed and dried to give the pure salt.
Crystallisation
Examiner keyword
Forming pure crystals of a solid from a saturated solution by slow cooling (or slow evaporation at room temperature). Used to obtain pure soluble salt crystals, often with water of crystallisation.
Saturated solution
Examiner keyword
A solution containing the MAXIMUM amount of dissolved solute at a given temperature. No more solute will dissolve. Test: crystals form on a cool glass rod dipped in the solution.
Water of crystallisation
Examiner keyword
Water molecules chemically bound into the crystal lattice of a hydrated salt. Written with a dot, e.g. CuSO₄·5H₂O. Removed on heating → anhydrous salt (white CuSO₄).
Hydrated / anhydrous
Examiner keyword
HYDRATED = contains water of crystallisation (e.g. blue CuSO₄·5H₂O). ANHYDROUS = NO water (e.g. white CuSO₄, formed by heating the hydrated salt).
Residue / filtrate
RESIDUE: the solid left on the filter paper. FILTRATE: the liquid that passes through. e.g. in CuSO₄ preparation, excess CuO is residue; CuSO₄ solution is filtrate.
In excess (during salt preparation)
Examiner keyword
Adding more of the base/metal/carbonate than the acid needs, so that ALL the acid is consumed. The unreacted excess is filtered off. Ensures no acid remains in the final product.

Common Mistakes and Misconceptions — Acids, Bases and Salt Preparations

The traps other students keep falling into on acids, bases and salt preparations questions — taken from recent Pearson Edexcel IGCSE 4CH1 examiner reports and mark schemes — and how to avoid them.

✕Boiling the solution 'to dryness' to obtain crystals
4CH1 Examiner Reports 2022-2024
Why it happens
Trying to skip the slow-cooling step.
How to avoid it
Never boil dry. (a) The crystals lose their water of crystallisation → wrong product. (b) Crystals are small, irregular and impure. ALWAYS evaporate to SATURATED only, then COOL SLOWLY for large pure crystals.
✕Forgetting to state that the base / metal / carbonate is added IN EXCESS
4CH1 Examiner Reports 2022-2024
Why it happens
Just saying 'add CuO until reaction stops'.
How to avoid it
The word 'EXCESS' is a mark-scheme keyword. State 'add CuO IN EXCESS until no more dissolves' — this guarantees no acid remains.
✕Forgetting to WASH the precipitate / crystals
Why it happens
Assuming the filter step removes all impurities.
How to avoid it
After filtering, ALWAYS wash with COLD distilled water (cold so the salt doesn't redissolve) to remove soluble impurities (K⁺, NO₃⁻, etc.) from the surface of the solid. This is a separate marked step.
✕Trying to make an insoluble salt by acid + base + crystallisation
Why it happens
Using the 'soluble salt' method as a default.
How to avoid it
Check solubility FIRST. Insoluble (AgCl, BaSO₄, PbI₂) → MUST use precipitation by mixing two soluble solutions. Soluble salts → use one of acid + insoluble base / titration / metal / carbonate.
✕Saying 'all sulfates are soluble' or 'all chlorides are soluble'
4CH1 Examiner Reports 2022-2024
Why it happens
Memorising the general rule but forgetting exceptions.
How to avoid it
Learn the exceptions: SULFATES insoluble: BaSO₄, PbSO₄ (CaSO₄ slightly). CHLORIDES insoluble: AgCl, PbCl₂. Same for Br⁻, I⁻ with Ag and Pb. CARBONATES mostly INSOLUBLE — only Group 1 + NH₄⁺ soluble.
✕Using 'add insoluble base + filter' method to make Na₂SO₄
Why it happens
Forgetting that NaOH and Na₂CO₃ are soluble.
How to avoid it
If both the acid AND the base are soluble (Na, K, NH₄⁺), there is no solid to filter off. Use TITRATION to measure the exact volume needed; then repeat without indicator.

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

Get a report showing which sub-topics you've nailed and which ones still need work.

Video lesson

Short walkthrough of the concepts students most often get stuck on.

Acids, Bases and Salt Preparations — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

Acids, Bases and Salt Preparations

Short Study Notes in the form of Slides

Short Study Notes — Acids, Bases and Salt Preparations

Detailed Notes

What you’ll learn

How it’s examined

1Preparing pure dry copper(II) sulfate crystals from CuO + H₂SO₄ (6 marks, Paper 1)

2Preparing insoluble lead(II) iodide by precipitation (5 marks, Paper 1)

3Applying solubility rules to predict precipitates (4 marks, Paper 1)

4Choosing the correct salt-preparation method for ZnSO₄ from given starting materials (5 marks, Paper 2)

Salt

Soluble salt

Insoluble salt

Precipitation reaction

Precipitate

Crystallisation

Saturated solution

Water of crystallisation

Hydrated / anhydrous

Residue / filtrate

In excess (during salt preparation)

✕Boiling the solution 'to dryness' to obtain crystals

✕Forgetting to state that the base / metal / carbonate is added IN EXCESS

✕Forgetting to WASH the precipitate / crystals

✕Trying to make an insoluble salt by acid + base + crystallisation

✕Saying 'all sulfates are soluble' or 'all chlorides are soluble'

✕Using 'add insoluble base + filter' method to make Na₂SO₄

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Acids, Bases and Salt Preparations — frequently asked questions