Practising IGCSE Chemistry past papers by topic is one of the most effective ways to identify weak areas and build confidence before exam day. Rather than working through entire papers, topic-focused practice lets you master one concept at a time — from balancing equations in Stoichiometry to drawing structural formulae in Organic Chemistry.

This guide organises key IGCSE Chemistry topics with the types of questions you’ll encounter, worked examples, essential equations, and exam tips. Whether you’re preparing for Cambridge IGCSE Chemistry (0620) or Edexcel International GCSE Chemistry (4CH1), these topic breakdowns align with both syllabuses.

Looking for full past papers instead? Visit our complete IGCSE Chemistry Past Papers guide.

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Why Practise Past Papers by Topic?

Working through past papers topic by topic offers several advantages over doing full papers alone:

• Targeted revision — spend more time on topics you find difficult
• Pattern recognition — examiners reuse question styles within each topic
• Confidence building — master one area before moving to the next
• Efficient use of time — no wasted effort on topics you’ve already nailed

The most successful IGCSE Chemistry students combine topic-based practice early in revision with full timed papers closer to the exam.

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Topic 1: Atomic Structure and the Periodic Table

What Examiners Ask

• Describe the structure of an atom (protons, neutrons, electrons)
• Explain isotopes and calculate relative atomic mass
• Deduce electron configurations for the first 20 elements
• Explain trends in the Periodic Table (group and period patterns)
• Describe properties of alkali metals, halogens, and noble gases

Key Concepts to Know

Relative atomic mass calculation:

$$A_r = \frac{\sum(\text{isotope mass} \times \text{abundance})}{100}$$

For example, chlorine has two isotopes: ³⁵Cl (75%) and ³⁷Cl (25%).

$$A_r = \frac{(35 \times 75) + (37 \times 25)}{100} = \frac{2625 + 925}{100} = 35.5$$

Electron configuration: Remember the 2, 8, 8 pattern for the first 20 elements. Potassium (19) is 2, 8, 8, 1 — not 2, 8, 9.

Worked Example

Q: Element X has the electron configuration 2, 8, 7. Identify the element and predict its chemical behaviour.

A: Element X has 17 electrons, so it is chlorine (Cl). It is in Group VII (halogens) and Period 3. It needs one more electron to achieve a stable octet, so it will readily gain one electron to form a Cl⁻ ion. It is a reactive non-metal that forms ionic bonds with metals and covalent bonds with other non-metals.

Exam Tip

When asked about Periodic Table trends, always link the trend to atomic structure. For example, reactivity increases down Group I because the outer electron is further from the nucleus and more easily lost — don’t just state the trend without an explanation.

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Topic 2: Chemical Bonding and Structure

What Examiners Ask

• Draw dot-and-cross diagrams for ionic and covalent compounds
• Explain the properties of ionic, covalent, and metallic substances
• Distinguish between simple molecular and giant covalent structures
• Explain why NaCl has a high melting point but CO₂ does not

Key Concepts to Know

Bonding Type	Particles	Forces	Melting Point	Conducts Electricity?
Ionic	Ions	Strong electrostatic	High	When molten/dissolved
Simple covalent	Molecules	Weak intermolecular	Low	No
Giant covalent	Atoms	Strong covalent bonds	Very high	No (except graphite)
Metallic	Cations + electrons	Metallic bonding	High	Yes (always)

Worked Example

Q: Draw a dot-and-cross diagram for magnesium chloride (MgCl₂) and explain why it conducts electricity when molten but not when solid.

A: Magnesium (2, 8, 2) transfers one electron to each chlorine atom (2, 8, 7), forming Mg²⁺ and two Cl⁻ ions. In the solid state, ions are held in a fixed lattice and cannot move, so no conduction occurs. When molten, the ions are free to move and can carry charge, allowing electrical conduction.

Exam Tip

A common mistake is saying “molecules” when describing ionic compounds. Ionic substances form lattices, not molecules. Examiners deduct marks for this error.

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Topic 3: Stoichiometry and the Mole Concept

What Examiners Ask

• Balance chemical equations
• Calculate moles, mass, and molar mass
• Perform reacting mass calculations
• Calculate percentage yield and atom economy
• Determine empirical and molecular formulae

Essential Equations

$$\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar mass (g/mol)}}$$

$$\text{Moles of gas} = \frac{\text{Volume (dm}^3\text{)}}{24} \quad \text{(at RTP)}$$

$$\text{Concentration (mol/dm}^3\text{)} = \frac{\text{Moles}}{\text{Volume (dm}^3\text{)}}$$

$$\text{Percentage yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100$$

Worked Example

Q: 4.8 g of magnesium reacts with excess hydrochloric acid. Calculate the volume of hydrogen gas produced at RTP. (Aᵣ: Mg = 24, H = 1, Cl = 35.5)

A:

Step 1 — Write the balanced equation: Mg + 2HCl → MgCl₂ + H₂

Step 2 — Calculate moles of Mg: Moles = 4.8 ÷ 24 = 0.2 mol

Step 3 — Use the mole ratio (1:1): Moles of H₂ = 0.2 mol

Step 4 — Calculate volume: Volume = 0.2 × 24 = 4.8 dm³

Exam Tip

Always show your working in calculation questions — even if your final answer is wrong, you can earn method marks for correct steps. Write the balanced equation first, then extract the mole ratio.

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Topic 4: Acids, Bases, and Salts

What Examiners Ask

• Define acids and bases in terms of proton transfer
• Describe reactions of acids with metals, bases, carbonates
• Explain the pH scale and use of indicators
• Describe methods for preparing soluble and insoluble salts
• Perform titration calculations

Key Reactions

Acid + Metal → Salt + Hydrogen $$Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2$$

Acid + Base → Salt + Water $$HCl + NaOH \rightarrow NaCl + H_2O$$

Acid + Carbonate → Salt + Water + Carbon dioxide $$2HCl + CaCO_3 \rightarrow CaCl_2 + H_2O + CO_2$$

Worked Example

Q: 25.0 cm³ of 0.10 mol/dm³ NaOH is neutralised by 20.0 cm³ of HCl. Calculate the concentration of the HCl.

A:

Moles of NaOH = 0.10 × (25.0 ÷ 1000) = 0.0025 mol

From the equation NaOH + HCl → NaCl + H₂O, the ratio is 1:1.

Moles of HCl = 0.0025 mol

Concentration of HCl = 0.0025 ÷ (20.0 ÷ 1000) = 0.125 mol/dm³

Exam Tip

For salt preparation questions, remember: soluble salts are made by titration or adding excess solid to acid, then filtering. Insoluble salts are made by precipitation — mixing two solutions and filtering the precipitate.

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Topic 5: Organic Chemistry

What Examiners Ask

• Name and draw structural formulae for alkanes, alkenes, and alcohols
• Describe addition and substitution reactions
• Explain fractional distillation and cracking of crude oil
• Describe polymerisation (addition and condensation)
• Test for alkenes using bromine water

Homologous Series Summary

Series	General Formula	Functional Group	Example
Alkanes	CₙH₂ₙ₊₂	None (C–C single bonds)	CH₄ (methane)
Alkenes	CₙH₂ₙ	C=C double bond	C₂H₄ (ethene)
Alcohols	CₙH₂ₙ₊₁OH	–OH	C₂H₅OH (ethanol)
Carboxylic acids	CₙH₂ₙ₊₁COOH	–COOH	CH₃COOH (ethanoic acid)

Worked Example

Q: Ethene reacts with bromine. Write the equation, name the product, and state what you would observe.

A:

$$C_2H_4 + Br_2 \rightarrow C_2H_4Br_2$$

The product is 1,2-dibromoethane. This is an addition reaction. You would observe the orange/brown bromine water decolourising (turning colourless), confirming the presence of a C=C double bond.

Exam Tip

When drawing structural formulae, always show all bonds clearly. Simply writing “C₂H₅OH” as a molecular formula will not earn full marks if the question asks for a displayed or structural formula.

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Topic 6: Energetics and Energy Changes

What Examiners Ask

• Define exothermic and endothermic reactions with examples
• Draw and interpret energy profile diagrams
• Explain activation energy and the effect of catalysts
• Calculate energy changes from bond energy data

Key Concept: Bond Energy Calculations

$$\Delta H = \sum(\text{bonds broken}) - \sum(\text{bonds formed})$$

• If ΔH is negative → exothermic (energy released)
• If ΔH is positive → endothermic (energy absorbed)

Worked Example

Q: Use the bond energies below to calculate the energy change for the combustion of methane.

CH₄ + 2O₂ → CO₂ + 2H₂O

Bond energies: C–H = 413 kJ/mol, O=O = 498 kJ/mol, C=O = 805 kJ/mol, O–H = 464 kJ/mol

A:

Bonds broken: 4 × C–H + 2 × O=O = (4 × 413) + (2 × 498) = 1652 + 996 = 2648 kJ

Bonds formed: 2 × C=O + 4 × O–H = (2 × 805) + (4 × 464) = 1610 + 1856 = 3466 kJ

ΔH = 2648 − 3466 = −818 kJ/mol

The reaction is exothermic (negative ΔH), which makes sense for combustion.

Exam Tip

Remember: breaking bonds requires energy (endothermic) and forming bonds releases energy (exothermic). The formula is always bonds broken minus bonds formed.

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Topic 7: Rates of Reaction and Equilibrium

What Examiners Ask

• Describe factors affecting rate (temperature, concentration, surface area, catalysts)
• Interpret rate graphs (volume of gas vs time, mass loss vs time)
• Explain collision theory
• Describe reversible reactions and dynamic equilibrium (Extended only)

Key Concept: Collision Theory

For a reaction to occur, particles must collide with sufficient energy (≥ activation energy) and the correct orientation. Increasing temperature makes particles move faster, leading to more frequent and more energetic collisions.

Worked Example

Q: Explain why powdered calcium carbonate reacts faster with hydrochloric acid than lumps of the same mass.

A: Powdered CaCO₃ has a larger surface area than lumps. This means more particles of CaCO₃ are exposed to the acid at any given time, resulting in a greater frequency of successful collisions per second. The rate of reaction is therefore higher. The total amount of product formed is the same in both cases — the powder just reaches completion faster.

Exam Tip

Rate questions often include graphs. When comparing two experiments on the same graph, look at: (1) the steepness of the initial slope (faster = steeper), and (2) whether the final amount of product is the same or different.

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Topic 8: Electrolysis

What Examiners Ask

• Define electrolysis and identify products at each electrode
• Explain the electrolysis of molten compounds and aqueous solutions
• Describe the electroplating process
• Write half-equations for electrode reactions (Extended)

Key Rules for Predicting Products

Molten compounds: The metal forms at the cathode (negative), the non-metal forms at the anode (positive).

Aqueous solutions:

• At the cathode: hydrogen is produced unless the metal is less reactive than hydrogen
• At the anode: if a halide ion is present, the halogen is produced; otherwise, oxygen is produced

Worked Example

Q: Predict the products of electrolysis of concentrated sodium chloride solution (brine). Write half-equations for each electrode.

A:

At the cathode (−): Hydrogen gas is produced (sodium is more reactive than hydrogen).

$$2H^+ + 2e^- \rightarrow H_2$$

At the anode (+): Chlorine gas is produced (chloride ions are present).

$$2Cl^- \rightarrow Cl_2 + 2e^-$$

Sodium hydroxide solution remains in the electrolyte.

Exam Tip

For half-equations, remember: reduction happens at the cathode (gain of electrons) and oxidation at the anode (loss of electrons). Use the mnemonic OILRIG — Oxidation Is Loss, Reduction Is Gain.

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How to Use Topic-Based Practice Effectively

Follow this strategy to maximise your revision:

1. Identify weak topics — take a diagnostic test or review past mock results
2. Study the theory first — read your textbook or notes before attempting questions
3. Attempt questions without notes — simulate exam conditions
4. Mark using the mark scheme — pay attention to the exact wording examiners expect
5. Track your scores — aim for improvement each time you revisit a topic
6. Move to full papers — once all topics are strong, practise under timed conditions

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Struggling with specific topics? A specialist IGCSE Chemistry tutor can identify your weak areas and provide targeted practice with past paper questions.

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Frequently Asked Questions

Which topics appear most often in IGCSE Chemistry exams?

Stoichiometry (mole calculations), Acids & Bases, and Organic Chemistry are consistently the most heavily weighted topics across both Cambridge and Edexcel papers.

Should I practise by topic or do full papers?

Both. Start with topic-based practice to build foundational understanding, then switch to full timed papers 4–6 weeks before the exam to build stamina and time management skills.

Where can I find more IGCSE Chemistry past papers?

Visit our complete IGCSE Chemistry Past Papers guide for links to Cambridge and Edexcel resources, mark schemes, and examiner reports.

How many past papers should I complete per topic?

Aim for at least 3–5 past paper questions per topic from different exam sessions. This exposes you to the variety of ways examiners can test the same concept.

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Looking for full IGCSE Chemistry past papers? Check out our comprehensive IGCSE Chemistry Past Papers guide for complete papers, mark schemes, and revision strategies.

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