IGCSE Chemistry (0620) – Complete Revision Notes
How to Use These Notes
- Read each unit’s Core Points (AO1 – factual knowledge) first.
- Study the Explanation (AO2 – understanding) and then the Evaluation (AO3 – applying, analysing, evaluating).
- Use the Supplement sections for the extended (non‑core) syllabus content that often appears in higher‑tier papers.
- Colour‑code the sections you find most challenging and revisit them regularly.
Unit 1 – States of Matter
Core Points (AO1)
- Three states of matter: solid, liquid, gas.
- Particle‑model description:
- Solids – particles vibrate in fixed positions; strong intermolecular forces.
- Liquids – particles slide past each other; moderate forces.
- Gases – particles move freely; negligible forces.
- Key terms: melting point, boiling point, evaporation, condensation, sublimation, diffusion.
Explanation (AO2)
Increasing temperature raises kinetic energy, weakening intermolecular forces and causing state changes. Diffusion is the random movement of particles; it is fastest in gases because the particles are far apart.
Supplement (Extended Content)
- Effect of molecular mass on diffusion: Lighter gases (e.g., H₂) diffuse more rapidly than heavier gases (e.g., CO₂) because the average speed of particles is inversely proportional to the square root of their molar mass ( Graham’s law).
Evaluation (AO3)
Discuss why water expands on freezing (open hydrogen‑bonded lattice) and the practical implications for pipes, containers and natural environments.
Unit 2 – Atoms, Elements & Compounds
Core Points (AO1)
- Structure of an atom: protons (+), neutrons (neutral), electrons (–).
- Atomic number (Z) = number of protons; mass number (A) = protons + neutrons.
- Isotopes: same Z, different A.
- Ions: loss of electrons → cation; gain of electrons → anion.
- Bonding types:
- Ionic – transfer of electrons, giant lattice, high melting points.
- Covalent – sharing of electrons; can be molecular or giant covalent.
- Metallic – sea of delocalised electrons; ductile, conductive.
Explanation (AO2)
Use the octet rule to predict formulas of simple ionic compounds (e.g., NaCl) and covalent molecules (e.g., CH₄). Explain why giant covalent structures such as SiO₂ have very high melting points – a continuous network of strong Si–O bonds.
Supplement (Extended Content)
- Calculating average atomic mass: Multiply the relative atomic mass of each isotope by its natural abundance (as a decimal) and add the results. Example: Average Ar of chlorine = (35.45 × 0.757) + (37.45 × 0.243) = 35.45 u.
- Electronic configuration (1–20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶. Use this to rationalise the position of transition metals.
- Giant covalent SiO₂: Each Si atom is tetrahedrally bonded to four O atoms; each O bridges two Si atoms, forming a three‑dimensional network.
Evaluation (AO3)
Compare properties of ionic, covalent (molecular), covalent (giant), and metallic solids and relate them to their structures.
Unit 3 – Stoichiometry
Core Points (AO1)
- Relative atomic mass (Ar) and relative formula mass (Mr) are used to calculate mass relationships.
- Law of conservation of mass – mass of reactants = mass of products.
- Simple mass‑ratio calculations (e.g., 10 g CaO + excess H₂CO₃ → mass of CaCO₃).
Explanation (AO2)
Step‑by‑step conversion: mass → moles (using Mr) → stoichiometric ratio from the balanced equation → mass of product.
Supplement (Extended Content)
- The mole concept: 1 mol = 6.022 × 10²³ particles (Avogadro’s constant). Molar mass (g mol⁻¹) = relative formula mass (Mr) in grams.
- Molar volume of a gas: At r.t.p. (25 °C, 1 atm) 1 mol of any gas occupies ≈ 24.5 cm³; at STP (0 °C, 1 atm) it is 22.4 cm³.
- Limiting‑reactant calculations: Identify the reactant that produces the smallest amount of product; use it to calculate theoretical yield.
- Percentage yield: (actual yield ÷ theoretical yield) × 100 %.
- Empirical and molecular formulae: Determine the simplest whole‑number ratio (empirical) then use the molar mass to find the molecular formula.
- Percentage composition & purity: (mass of element in 100 g of compound ÷ 100) × 100 %.
Evaluation (AO3)
Identify common sources of error (e.g., incomplete reaction, impurity, measurement uncertainty) and discuss how they affect calculated yields and percentages.
Unit 4 – Electrochemistry
Core Points (AO1)
- Electrolysis – passage of electricity through a molten ionic compound or aqueous solution.
- Electrolytic cell components: anode (‑), cathode (+), electrolyte, power source.
- Typical products:
- Molten NaCl → Na (cathode) + Cl₂ (anode).
- Aqueous CuSO₄ → Cu (cathode) + O₂/H₂ (anode depending on conditions).
- Electroplating and metal extraction (e.g., aluminium – Hall‑Héroult process).
Explanation (AO2)
Write half‑reactions for oxidation and reduction, balance electrons, then combine to give the overall equation.
Supplement (Extended Content)
- Half‑equations for aqueous CuSO₄:
- Cathode (reduction): Cu²⁺ + 2 e⁻ → Cu(s)
- Anode (oxidation of water): 2 H₂O → O₂(g) + 4 H⁺ + 4 e⁻
- Hydrogen‑oxygen fuel cell:
| Aspect | Advantage | Disadvantage |
|---|
| Energy efficiency | Higher than combustion engines (≈ 60 % vs 25 %) | Requires pure H₂ and O₂, costly catalysts |
| Emissions | Only water vapour produced | Hydrogen production often from fossil fuels |
| Noise | Very quiet operation | Complex fuel‑storage systems |
Evaluation (AO3)
Discuss the energy efficiency of electrolysis versus thermochemical reduction and the environmental impact of large‑scale chlorine production.
Unit 5 – Chemical Energetics
Core Points (AO1)
- Exothermic reactions – energy released; temperature rise; ΔH < 0.
- Endothermic reactions – energy absorbed; temperature fall; ΔH > 0.
- Activation energy – minimum energy barrier that must be overcome for a reaction to proceed.
- Bond‑energy diagram – reactants → transition state → products.
Explanation (AO2)
Use enthalpy change (ΔH) to predict whether a reaction feels hot or cold; relate ΔH to the balance between energy required to break bonds and energy released on forming new bonds.
Supplement (Extended Content)
Evaluation (AO3)
Analyse why catalysts lower activation energy without being consumed, giving examples such as Pt in catalytic converters and enzymes in biological systems.
Unit 6 – Chemical Reactions
Core Points (AO1)
- Physical change vs chemical change.
- Reaction types: combination, decomposition, single‑displacement, double‑displacement, combustion.
- Factors affecting rate: concentration, temperature, surface area, catalysts.
- Reversible reactions and dynamic equilibrium (qualitative description only).
- Redox: oxidation = loss of electrons; reduction = gain of electrons.
Explanation (AO2)
Write balanced equations, identify oxidising and reducing agents, and use oxidation numbers where required.
Supplement (Extended Content)
- Collision theory: Reactions occur when particles collide with sufficient energy (activation energy) and proper orientation.
- Assigning oxidation numbers: Rules – element in its free state = 0; alkali metals = +1; halogens = –1 (unless combined with more electronegative element); oxygen = –2 (except in peroxides); hydrogen = +1 (except with metals).
- Half‑equations for redox: Example for the reaction of magnesium with hydrochloric acid:
- Oxidation: Mg → Mg²⁺ + 2 e⁻
- Reduction: 2 H⁺ + 2 e⁻ → H₂(g)
- Equilibrium & Le Chatelier (useful for higher‑tier):
| Change | Effect on equilibrium |
|---|
| Increase concentration of reactants | Shift right (more products) |
| Increase temperature for exothermic reaction | Shift left (absorbs heat) |
| Increase pressure (gaseous system) | Shift towards side with fewer moles of gas |
| Add catalyst | No shift; rate of both forward and reverse reactions increases equally |
Evaluation (AO3)
Discuss how changing conditions (concentration, temperature, pressure, catalysts) can be used to control industrial processes such as the Haber‑Bosch synthesis of NH₃.
Unit 7 – Acids, Bases & Salts
Core Points (AO1)
- Acid – produces H⁺ in water; characteristic reactions (e.g., with metals → H₂).
- Base – produces OH⁻ in water; characteristic reactions (e.g., with acids → salt + water).
- Strong acids/bases dissociate completely (HCl, H₂SO₄, NaOH, KOH).
- Weak acids/bases dissociate only partially (acetic acid, ammonia).
- Neutralisation: acid + base → salt + water (exothermic).
- Solubility rules for common salts.
Explanation (AO2)
Use the pH scale (0‑14) to describe acidity/basicity; pH = –log₁₀[H⁺]. Relate pH to [H⁺] concentration and to the strength of the acid or base.
Supplement (Extended Content)
- Calculating pH of a strong acid: For 0.01 M HCl, [H⁺] = 0.01 M → pH = 2.
- Weak acid example (acetic acid, Ka = 1.8 × 10⁻⁵): Use ICE table to find [H⁺] and then pH.
- Buffer concept (optional for higher tier): A mixture of a weak acid and its conjugate base resists pH change.
Evaluation (AO3)
Compare the effectiveness of different antacids (e.g., Mg(OH)₂ vs CaCO₃) in neutralising stomach acid, considering factors such as reaction speed, side‑effects, and dosage.
Unit 8 – The Periodic Table
Core Points (AO1)
- Groups (vertical) and periods (horizontal).
- Trends:
- Atomic radius – decreases across a period, increases down a group.
- Ionisation energy – increases across a period, decreases down a group.
- Electronegativity – similar trend to ionisation energy.
- Key groups for the syllabus:
- Group 1 – Alkali metals (highly reactive, +1 ions).
- Group 7 – Halogens (highly reactive non‑metals, –1 ions).
- Transition metals – variable oxidation states, coloured compounds.
- Group 18 – Noble gases (very low reactivity).
Explanation (AO2)
Explain why metallic character increases down a group (electron shielding reduces effective nuclear charge) and decreases across a period (increasing nuclear charge holds electrons more tightly).
Evaluation (AO3)
Assess how periodic trends help predict the properties of unknown elements and guide the selection of materials for specific applications.
Unit 9 – Metals (including Alloys)
9.1 General Properties of Metals
- Good conductors of heat and electricity (sea of delocalised electrons).
- Malleable, ductile, lustrous, generally high melting and boiling points (except alkali metals).
- Typical reactions:
- With acids → H₂ + salt.
- With oxygen → metal oxide (often basic).
- With water (alkali metals) → metal hydroxide + H₂.
9.2 Extraction of Metals
| Metal | Principal Ore | Extraction Method | Key Reaction(s) |
|---|
| Iron | Hematite (Fe₂O₃) | Blast furnace (reduction with coke) | Fe₂O₃ + 3 C → 2 Fe + 3 CO₂ |
| Aluminium | Bauxite (Al₂O₃·2H₂O) | Hall‑Héroult (electrolysis of molten cryolite‑alumina) | 2 Al₂O₃ → 4 Al + 3 O₂ (electrolysis) |
| Copper | Cu₂S, Cu₂O, CuSO₄·5H₂O | Fire‑refining or electrolytic refining | Cu²⁺ + 2 e⁻ → Cu (electrolysis) |
9.3 Corrosion and Protection
- Corrosion = oxidation of metal in the presence of water/oxygen (e.g., Fe → Fe₂O₃·nH₂O).
- Prevention methods:
- Coating – paint, galvanising (Zn coating).
- Alloying – stainless steel (Cr forms a passive Cr₂O₃ layer).
- Cathodic protection – sacrificial anode (e.g., Zn attached to iron).
9.4 Alloys – Why They Are Harder & Stronger
An alloy is a mixture of two or more metals, or a metal with a non‑metal, that retains metallic properties. Adding a second element disturbs the regular crystal lattice of the base metal, making dislocation movement more difficult and thereby increasing hardness and strength.
Mechanisms of Strengthening (AO2)
- Solid‑solution strengthening – solute atoms replace (substitutional) or occupy (interstitial) sites, causing lattice distortion.
- Precipitation (age) hardening – fine particles of a second phase form during heat treatment and block dislocations.
- Grain‑size reduction – more grain boundaries act as barriers to slip.
- Intermetallic compounds – ordered structures with very high hardness (e.g., Ni₃Al).
Common Alloys
| Alloy | Principal Components | Key Properties | Typical Uses |
|---|
| Brass | Cu + Zn | Corrosion‑resistant, low friction, gold‑like colour | Musical instruments, fittings, decorative hardware |
| Bronze | Cu + Sn (± Al, P, Si) | High strength, excellent wear resistance, low friction | Statues, ship propellers, bearings |
| Stainless steel | Fe + Cr (≥10.5 %) + Ni (optional) | Very high corrosion resistance, good strength, can be hardened | Cutlery, medical instruments, construction, food‑processing equipment |
| Aluminium alloy (e.g., 6061) | Al + Mg + Si | Lightweight, good strength‑to‑weight ratio, corrosion resistant | Aerospace components, bicycle frames, automotive parts |
| Carbon steel | Fe + C (0.2–2 %) | High tensile strength, hardenable by quenching & tempering | Construction beams, tools, machinery parts |
Comparing Pure Metals with Their Alloys
| Property | Pure Metal | Alloy |
|---|
| Hardness | Usually soft (e.g., pure copper is very ductile) | Significantly higher (e.g., brass > copper) |
| Tensile strength | Relatively low | Higher due to solid‑solution & precipitation hardening |
| Corrosion resistance | Varies; many corrode rapidly | Often improved (e.g., stainless steel resists rust) |
| Melting point | Fixed for each element | Can be lower or higher depending on composition (e.g., solder alloys melt below pure tin) |
Key Points to Remember (AO1)
- Alloys retain metallic properties but are generally harder and stronger than the constituent pure metals.
- Strengthening arises because alloying atoms disturb the regular lattice, hindering dislocation motion.
- Improved mechanical and chemical properties make alloys far more useful in industry and everyday life.
- Choosing an alloy involves balancing hardness, strength, corrosion resistance, cost, and required physical properties.
Unit 10 – Chemistry of the Environment
Core Points (AO1)
- Key atmospheric gases: N₂, O₂, CO₂, CH₄, N₂O.
- Greenhouse effect – trapping of infrared radiation by greenhouse gases.
- Acid rain – formation of H₂SO₄ and HNO₃ from SO₂ and NOₓ emissions.
- Ozone layer – depletion by CFCs; role of UV‑absorbing ozone (O₃).
- Water pollution – eutrophication, oil spills, heavy‑metal contamination.
- Waste management – recycling, incineration, landfill, composting.
Explanation (AO2)
Explain how combustion of fossil fuels releases CO₂ and CH₄, enhancing the natural greenhouse effect. Show the chemical equations for the formation of sulfuric and nitric acids in the atmosphere.
Evaluation (AO3)
Discuss the advantages and disadvantages of different waste‑treatment methods (e.g., recycling reduces resource extraction but may be energy‑intensive; incineration reduces volume but can emit pollutants).
Quick Reference Tables
Common Ion‑Formation Rules (AO1)
| Element | Typical Ion | Charge |
|---|
| Alkali metals (Group 1) | +1 | + |
| Alkaline earth metals (Group 2) | +2 | + |
| Halogens (Group 7) | ‑1 (except with oxygen) | ‑ |
| Oxygen | ‑2 (except in peroxides) | ‑ |
| Hydrogen | +1 (with non‑metals) or ‑1 (with metals) | ± |
Useful Constants (AO1)
- Avogadro’s constant: 6.022 × 10²³ particles mol⁻¹
- Standard molar volume of a gas at r.t.p.: 24.5 cm³ mol⁻¹
- Standard molar volume of a gas at STP: 22.4 cm³ mol⁻¹
- ΔH (standard enthalpy change) sign convention: exothermic < 0, endothermic > 0
Exam‑Style Practice Tips
- Always start a calculation with the given masses, convert to moles using Mr, apply the stoichiometric ratio, then convert back to the required unit.
- When writing half‑equations, balance mass first, then charge, and finally combine to check that electrons cancel.
- For alloy questions, identify the strengthening mechanism being asked about and link it to the observed property (e.g., “precipitation hardening → increased hardness”).
- In environmental chemistry, remember to link a chemical process to its real‑world impact (e.g., “SO₂ → H₂SO₄ → acid rain → damage to buildings”).