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

This guide teaches 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). It covers points 2.28–2.43 in depth. The qualification is untiered. Points 2.33C, 2.40C, 2.41C and 2.43C are bold with a C reference, so they are assessed on Paper 2 only, and this guide labels them Paper 2 only. Other points can appear on either paper.

Links: 4CH1 course hub, printable checklist, revision notes for this unit, practice on Groups 1 and 7, acids, salts and tests, the wider Reactivity and group trends guide and the free 4CH1 all-topics diagnostic.

What these sub-topics cover

Spec point	What you must be able to do	Paper
2.28	Use litmus, phenolphthalein and methyl orange to tell acidic from alkaline solutions	1 and 2
2.29	Classify solutions on the pH scale 0–14	1 and 2
2.30	Use universal indicator to measure approximate pH	1 and 2
2.31	Know acids give H⁺ ions and alkalis give OH⁻ ions in water	1 and 2
2.32	Know alkalis neutralise acids	1 and 2
2.33C	Describe how to carry out an acid–alkali titration	Paper 2 only
2.34	Apply the solubility rules for ionic compounds	1 and 2
2.35–2.36	Explain acids and bases as proton donors and acceptors	1 and 2
2.37	Describe reactions of HCl, H₂SO₄ and HNO₃ with metals, bases and carbonates	1 and 2
2.38	Know which substances act as bases, and what an alkali is	1 and 2
2.39	Make a pure, dry soluble salt from an insoluble reactant	1 and 2
2.40C	Make a pure, dry soluble salt from an acid and an alkali	Paper 2 only
2.41C	Make a pure, dry insoluble salt from two soluble reactants	Paper 2 only
2.42	Practical: hydrated copper(II) sulfate from copper(II) oxide	1 and 2
2.43C	Practical: pure, dry lead(II) sulfate	Paper 2 only

Indicators (2.28)

An indicator changes colour depending on whether a solution is acidic or alkaline. You need three single indicators.

Indicator	Colour in acid	Colour in alkali
Litmus	red	blue
Phenolphthalein	colourless	pink
Methyl orange	red	yellow

Phenolphthalein is the odd one out: it has no colour in acid. Methyl orange is orange only near the change point, so “orange” is not its alkali colour.

The pH scale and universal indicator (2.29–2.30)

The specification fixes five bands. Learn the numbers exactly.

pH	Classification
0–3	strongly acidic
4–6	weakly acidic
7	neutral
8–10	weakly alkaline
11–14	strongly alkaline

Universal indicator is a mixture of indicators, so it changes colour gradually across the whole range. Add a few drops to the solution (or dip the paper in it), then match the colour against a pH chart. Typical colours run from red (strong acid) through orange and yellow, green (neutral), blue, to purple (strong alkali). The result is an approximate pH.

Worked example. Universal indicator turns four solutions these colours: P, yellow; Q, blue; R, purple; S, green. Reading each against the chart gives P about pH 5 (weakly acidic), Q about pH 9 (weakly alkaline), R about pH 13 (strongly alkaline) and S pH 7 (neutral).

Hydrogen ions, hydroxide ions and neutralisation (2.31–2.32)

In aqueous solution:

• an acid is a source of hydrogen ions, H⁺(aq)
• an alkali is a source of hydroxide ions, OH⁻(aq)

When an alkali is added to an acid, the hydroxide ions remove the hydrogen ions. This is neutralisation:

H⁺(aq) + OH⁻(aq) → H₂O(l)

The other ions stay in solution and form the salt.

Carrying out an acid–alkali titration (2.33C, Paper 2 only)

A titration finds the exact volume of acid that neutralises a fixed volume of alkali.

1. Use a pipette and filler to put 25.0 cm³ of alkali into a conical flask.
2. Add a few drops of a single indicator, such as phenolphthalein or methyl orange.
3. Fill a burette with the acid and record the starting reading at eye level.
4. Add acid while swirling the flask. Near the end point, add it drop by drop.
5. Stop when the indicator just changes colour permanently. Record the final reading. The titre is final minus initial.
6. Do a rough titration first, then repeat accurately until you have concordant titres (usually within 0.10 cm³ of each other).

A white tile under the flask shows the colour change clearly. Do not use universal indicator: its gradual change gives no sharp end point.

Worked example: choosing concordant titres. A student records 23.20 (rough), 22.60, 22.95, 22.70 and 22.65 cm³.

Ignore the rough titre (23.20).
22.60, 22.70 and 22.65 lie within 0.10 cm³ of each other.
22.95 is 0.25 cm³ above 22.70, so leave it out.
Mean titre = (22.60 + 22.70 + 22.65) ÷ 3 = 22.65 cm³

Concentration calculations from titres belong to point 1.34C; see the chemical formulae, equations and calculations guide.

Solubility rules (2.34)

Whether a salt dissolves decides how it is made.

Compounds	Rule
Sodium, potassium, ammonium	common ones are all soluble
Nitrates	all soluble
Chlorides	soluble, except silver and lead(II)
Sulfates	soluble, except barium, calcium and lead(II)
Carbonates	insoluble, except sodium, potassium and ammonium
Hydroxides	insoluble, except sodium, potassium and calcium (calcium hydroxide is slightly soluble)

Worked example. Silver chloride is insoluble (chloride exception). Magnesium sulfate is soluble (not a sulfate exception). Ammonium carbonate is soluble (the ammonium rule overrides the carbonate rule).

Acids and bases as proton transfer (2.35–2.36)

A hydrogen ion, H⁺, is a hydrogen atom that has lost its electron, so it is just a proton. That gives a wider definition:

• an acid is a proton donor
• a base is a proton acceptor

With an oxide, the oxide ion accepts protons:

O²⁻ + 2H⁺ → H₂O

With a carbonate, the carbonate ion accepts protons, and carbonic acid breaks down:

CO₃²⁻ + 2H⁺ → H₂O + CO₂

Ammonia accepts a proton to become the ammonium ion:

NH₃ + H⁺ → NH₄⁺

Which substances are bases? (2.38)

Metal oxides, metal hydroxides and ammonia can all act as bases. An alkali is a base that is soluble in water. So every alkali is a base, but not every base is an alkali.

Copper(II) oxide is a base but not an alkali; sodium hydroxide and ammonia solution are alkalis.

Reactions of acids that make salts (2.37)

Hydrochloric acid makes chlorides, sulfuric acid makes sulfates and nitric acid makes nitrates. Learn three general equations.

acid + metal     → salt + hydrogen
acid + base      → salt + water
acid + carbonate → salt + water + carbon dioxide

Examples with state symbols:

Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)
MgO(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l)
Ca(OH)₂(aq) + 2HCl(aq) → CaCl₂(aq) + 2H₂O(l)
2NH₃(aq) + H₂SO₄(aq) → (NH₄)₂SO₄(aq)
ZnCO₃(s) + 2HNO₃(aq) → Zn(NO₃)₂(aq) + H₂O(l) + CO₂(g)

Ammonia gives an ammonium salt and no water. Nitric acid with metals is excluded from the specification.

Choosing a method for preparing a salt

Decide using two questions.

Is the salt soluble?
  NO  → precipitation: mix two soluble reactants (2.41C)
  YES → Is one reactant insoluble (metal, oxide, carbonate, insoluble hydroxide)?
          YES → add excess solid to acid, filter (2.39, 2.42)
          NO  → acid + alkali by titration (2.40C)

Sodium, potassium and ammonium salts need titration: their bases are soluble, so excess could not be filtered off.

Soluble salt from an insoluble reactant (2.39)

Worked example: zinc sulfate crystals from zinc carbonate.

1. Warm dilute sulfuric acid gently in a beaker.
2. Add zinc carbonate a spatula at a time, stirring, until fizzing stops and some solid stays at the bottom. The carbonate is now in excess, so all the acid has reacted.
3. Filter. The residue is unreacted zinc carbonate; the filtrate is zinc sulfate solution.
4. Heat the filtrate in an evaporating basin to remove some water, until it is saturated (crystals form on a cold glass rod dipped in it).
5. Leave it to cool so crystals form. Filter them off and dry them between filter papers or in a warm place.

ZnCO₃(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂O(l) + CO₂(g)

Practical: hydrated copper(II) sulfate (2.42)

This is the required practical for 2.39. Black copper(II) oxide is added to warm dilute sulfuric acid until some black solid remains. The solution turns blue. Filter off the excess oxide, partly evaporate the blue filtrate, then cool to give blue crystals of hydrated copper(II) sulfate, CuSO₄·5H₂O. Dry them on filter paper.

CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l)

Do not heat to dryness: that removes the water of crystallisation.

Worked example: yield. A student starts with 2.00 g of copper(II) oxide and obtains 5.02 g of dry crystals. Use Ar: Cu = 63.5, S = 32, O = 16, H = 1.

Mr CuO = 63.5 + 16 = 79.5
Mr CuSO₄·5H₂O = 63.5 + 32 + 64 + 5 × 18 = 249.5
moles CuO = 2.00 ÷ 79.5 = 0.02516 mol
1 mol CuO gives 1 mol CuSO₄·5H₂O
theoretical mass = 0.02516 × 249.5 = 6.277 g
percentage yield = 5.02 ÷ 6.277 × 100 = 80.0%

Some salt stays dissolved after crystallisation, so the yield is below 100%.

Soluble salt from an acid and an alkali (2.40C, Paper 2 only)

Worked example: potassium nitrate. Both reactants are soluble, so excess cannot be filtered off.

1. Titrate 25.0 cm³ of potassium hydroxide solution with dilute nitric acid, using a few drops of indicator, to find the exact volume of acid needed.
2. Repeat with fresh alkali and exactly that volume of acid, but no indicator, so the salt is not contaminated.
3. Evaporate some water, leave to crystallise, then filter and dry the crystals.

Insoluble salt from two soluble reactants (2.41C, Paper 2 only)

An insoluble salt forms as a precipitate when two solutions are mixed. Choose one solution containing the metal ion and one containing the anion. Nitrates and sodium or potassium compounds are good choices, because they are soluble.

Worked example: barium sulfate from barium chloride and sodium sulfate solutions.

BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)

1. Mix the two solutions in a beaker. A white precipitate forms.
2. Filter to collect the precipitate.
3. Wash it on the filter paper with distilled water to remove the soluble sodium chloride.
4. Dry it in a warm oven or between filter papers.

Practical: lead(II) sulfate (2.43C, Paper 2 only)

Mix lead(II) nitrate solution with dilute sulfuric acid (or any soluble sulfate). White lead(II) sulfate precipitates. Filter, wash with distilled water, and dry in a warm oven.

Pb(NO₃)₂(aq) + H₂SO₄(aq) → PbSO₄(s) + 2HNO₃(aq)
Pb²⁺(aq) + SO₄²⁻(aq) → PbSO₄(s)

Lead compounds are toxic: wear gloves and eye protection.

Common errors

• Writing methyl orange as “orange” in alkali. It is yellow.
• Using universal indicator for a titration.
• Including the rough titre, or a non-concordant one, in the mean.
• Calling calcium sulfate soluble. It is one of the three sulfate exceptions.
• Saying an acid “donates hydrogen” instead of a proton.
• Calling copper(II) oxide an alkali. It is an insoluble base.
• Forgetting to add the insoluble reactant in excess, or forgetting to filter it off.
• Heating hydrated crystals to dryness.
• Leaving out the washing step for a precipitate.

Next steps

Condense this with the revision notes, then try questions 6–10 of the Groups 1 and 7, acids, salts and tests practice set.

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.