Pearson Edexcel International GCSE Chemistry 4CH1: Acids, alkalis and titrations; acids, bases and salt preparations – Revision Notes

These notes condense sub-topics (f) Acids, alkalis and titrations and (g) Acids, bases and salt preparations of Topic 2, Inorganic chemistry, in the Pearson Edexcel International GCSE Chemistry (4CH1) specification, Issue 3 (September 2024). They cover points 2.28–2.43. The qualification is untiered; points 2.33C, 2.40C, 2.41C and 2.43C are bold with a C reference, so they are Paper 2 only. For full explanations and worked examples, use the acids, alkalis and salts study guide.

Links: 4CH1 Chemistry course hub, printable 4CH1 checklist, free 4CH1 all-topics diagnostic. For exam-style questions, use the Groups 1 and 7, acids, salts and tests practice set. Neighbouring content is in the Reactivity and group trends revision notes.

2.28 Indicator colours

Indicator	Acid	Alkali
Litmus	red	blue
Phenolphthalein	colourless	pink
Methyl orange	red	yellow

Memory hook: only phenolphthalein is colourless in acid; only methyl orange is yellow in alkali.

2.29–2.30 pH bands and universal indicator

pH	Class	Typical universal indicator colour
0–3	strongly acidic	red
4–6	weakly acidic	orange to yellow
7	neutral	green
8–10	weakly alkaline	blue
11–14	strongly alkaline	purple

Method in steps: measuring pH

1. Add a few drops of universal indicator (or use the paper).
2. Compare the colour with a pH colour chart.
3. Read off the nearest whole-number pH and name its band.

Universal indicator gives an approximate pH. It is a mixture, so its colour changes gradually.

2.31–2.32 Ions and neutralisation

• Acids in aqueous solution are a source of H⁺ ions.
• Alkalis in aqueous solution are a source of OH⁻ ions.
• Alkalis neutralise acids: H⁺(aq) + OH⁻(aq) → H₂O(l)

2.33C Titration (Paper 2 only)

Method in steps

1. Pipette 25.0 cm³ of alkali into a conical flask (pipette filler).
2. Add a few drops of phenolphthalein or methyl orange.
3. Fill the burette with acid; read the start volume at eye level.
4. Add acid with swirling; dropwise near the end point.
5. Stop at the first permanent colour change; titre = final − initial.
6. Rough run first, then repeat until titres are concordant.

Key points

• Burette readings are usually taken to the nearest 0.05 cm³.
• Concordant titres: usually within 0.10 cm³ of each other.
• Mean titre uses only concordant results, never the rough one.
• Use a white tile. Never universal indicator (no sharp end point).

2.34 Solubility rules

Always or usually soluble	Exceptions (insoluble)
Sodium, potassium, ammonium compounds	none
All nitrates	none
Chlorides	silver, lead(II)
Sulfates	barium, calcium, lead(II)

Usually insoluble	Exceptions (soluble)
Carbonates	sodium, potassium, ammonium
Hydroxides	sodium, potassium, calcium (calcium hydroxide only slightly)

Lead(II) appears in two exception lists: its chloride and its sulfate are both insoluble.

2.35–2.36 Proton transfer

• H⁺ is a proton.
• Acid = proton donor. Base = proton acceptor.
• In a neutralisation the base takes H⁺ from the acid.

Small reminders:

OH⁻ + H⁺ → H₂O          (hydroxide ion accepts a proton)
NH₃ + H⁺ → NH₄⁺         (ammonia accepts a proton)

2.38 Bases and alkalis

• Bases: metal oxides, metal hydroxides, ammonia.
• Alkali = a base that is soluble in water.
• All alkalis are bases; not all bases are alkalis.

2.37 Making salts from acids

Acid	Salt name ending
hydrochloric, HCl	chloride
sulfuric, H₂SO₄	sulfate
nitric, HNO₃	nitrate

acid + metal     → salt + hydrogen          (not nitric acid: excluded)
acid + base      → salt + water
acid + carbonate → salt + water + carbon dioxide

Small worked reminder:

Fe(s) + H₂SO₄(aq) → FeSO₄(aq) + H₂(g)
CuO(s) + 2HCl(aq) → CuCl₂(aq) + H₂O(l)
Na₂CO₃(aq) + 2HNO₃(aq) → 2NaNO₃(aq) + H₂O(l) + CO₂(g)

Choosing a salt preparation

Salt wanted	Reactants	Method	Spec
Soluble; base, metal or carbonate insoluble	acid + excess solid	add excess, filter, crystallise	2.39, 2.42
Soluble; Na, K or NH₄ salt	acid + alkali	titrate, repeat without indicator, crystallise	2.40C
Insoluble	two soluble solutions	precipitate, filter, wash, dry	2.41C, 2.43C

2.39 and 2.42 Excess-solid method

Method in steps (required practical: copper(II) oxide + dilute sulfuric acid → hydrated copper(II) sulfate)

1. Warm the dilute acid.
2. Add the solid a little at a time, stirring, until some is left over (excess).
3. Filter off the excess solid; keep the filtrate.
4. Evaporate some water until the solution is saturated.
5. Leave to cool and crystallise.
6. Filter off the crystals; dry between filter papers.

• Excess solid means all the acid is used up.
• Copper(II) sulfate: black solid dissolves, solution turns blue, blue crystals of CuSO₄·5H₂O.
• Do not evaporate to dryness: the crystals lose their water of crystallisation.

2.40C Acid + alkali (Paper 2 only)

1. Titrate to find the exact volume of acid needed.
2. Mix the same volumes again with NO indicator.
3. Evaporate some water, crystallise, filter, dry.

Why: both reactants are soluble, so an excess could not be filtered off, and indicator would contaminate the salt.

2.41C and 2.43C Precipitation (Paper 2 only)

1. Mix two solutions: one gives the cation, one gives the anion.
2. Filter off the precipitate.
3. Wash it with distilled water (removes the soluble by-product).
4. Dry it in a warm oven or between filter papers.

Lead(II) sulfate: lead(II) nitrate solution + dilute sulfuric acid or a soluble sulfate. Ionic equation: Pb²⁺(aq) + SO₄²⁻(aq) → PbSO₄(s). Lead compounds are toxic.

Must-know distinctions

• Base vs alkali: an alkali is a base that dissolves in water.
• Acid definitions: source of H⁺ in water (2.31) vs proton donor (2.36). Both are correct; use the one the question names.
• Filtrate vs residue: in the excess method you keep the filtrate; in precipitation you keep the residue.
• Rough vs concordant titre: only concordant titres go into the mean.
• Universal indicator vs single indicator: universal indicator measures pH; a single indicator marks a titration end point.
• Hydrated vs anhydrous copper(II) sulfate: blue crystals with water of crystallisation vs white powder without.

Quick self-test

1. State the colour of phenolphthalein in a solution of pH 2 and in a solution of pH 12.
2. Classify a solution of pH 4 and a solution of pH 11.
3. Name the ion that alkalis provide in aqueous solution.
4. A student records titres of 19.10 (rough), 18.45, 18.80, 18.50 and 18.40 cm³. Calculate the mean titre.
5. Classify as soluble or insoluble: silver nitrate, calcium sulfate, sodium carbonate, iron(II) hydroxide.
6. In CuO + 2HCl → CuCl₂ + H₂O, identify the proton donor.
7. Write the balanced equation for potassium carbonate reacting with dilute sulfuric acid.
8. State the best method to make (a) silver chloride, (b) sodium sulfate, (c) magnesium nitrate.
9. Suggest two soluble reactants to make lead(II) chloride.
10. Calculate the mass of lead(II) sulfate formed when 6.62 g of lead(II) nitrate reacts completely with excess sulfuric acid. (Ar: Pb = 207, N = 14, O = 16, S = 32)
11. Explain why the precipitate is washed with distilled water.
12. Explain why magnesium hydroxide is a base but not an alkali.

Answers

1. Colourless at pH 2; pink at pH 12.
2. pH 4 is weakly acidic; pH 11 is strongly alkaline.
3. Hydroxide ion, OH⁻.
4. Concordant: 18.45, 18.50, 18.40 (18.80 is too far off). Mean = (18.45 + 18.50 + 18.40) ÷ 3 = 18.45 cm³.
5. Silver nitrate soluble; calcium sulfate insoluble; sodium carbonate soluble; iron(II) hydroxide insoluble.
6. HCl donates protons; CuO (the oxide ion) accepts them.
7. K₂CO₃(aq) + H₂SO₄(aq) → K₂SO₄(aq) + H₂O(l) + CO₂(g)
8. (a) Precipitation, e.g. silver nitrate + sodium chloride solutions. (b) Titration of sodium hydroxide with sulfuric acid. (c) Excess magnesium oxide or magnesium carbonate with nitric acid, then filter and crystallise.
9. Lead(II) nitrate solution and sodium chloride solution (or dilute hydrochloric acid).
10. Mr Pb(NO₃)₂ = 331; Mr PbSO₄ = 303. Moles = 6.62 ÷ 331 = 0.0200 mol. Mass = 0.0200 × 303 = 6.06 g.
11. To remove the soluble by-product (and leftover solution) so the salt is pure.
12. It accepts protons (reacts with acids), so it is a base, but it is insoluble in water.

Where marks are usually lost

• Giving “orange” for methyl orange in alkali, or “purple” for litmus in alkali.
• Stating a pH number when the question asks you to classify, or putting pH 7 in a weak band.
• Including the rough titre, or a non-concordant one, in the mean.
• Choosing universal indicator for a titration.
• Calling calcium sulfate or lead(II) chloride soluble.
• Writing “hydrogen” or “H” instead of “proton” or “H⁺” in a proton-transfer answer.
• Reacting nitric acid with a metal, or adding the metal oxide without saying “until in excess”.
• Forgetting to filter off the excess solid before crystallising.
• In 2.40C, forgetting to repeat the titration without indicator.
• Leaving out washing in a precipitation method, or heating hydrated crystals to dryness.

Official syllabus

Pearson Edexcel International GCSE in Chemistry (4CH1) specification, Issue 3, September 2024, published by Pearson Education Limited – Topic 2 Inorganic chemistry, sub-topics (f) Acids, alkalis and titrations and (g) Acids, bases and salt preparations, points 2.28–2.43.