IGCSE Chemistry (0620) – Complete Syllabus Notes
1. States of Matter
1.1 General features
- Solids – particles tightly packed in a regular arrangement; only vibrate.
- Liquids – particles close together but can move past one another; flow.
- Gases – particles far apart; move freely and fill the container.
1.2 Kinetic‑Particle Theory (KPT)
- Temperature = average kinetic energy of particles.
- Pressure is caused by collisions of particles with the walls of the container.
- Increasing temperature → faster particles → higher pressure (if volume fixed).
1.3 Diffusion & Effusion
Rate of diffusion ∝ 1/√M (M = molar mass). Faster in gases than in liquids.
1.4 Phase changes & heating/cooling curves
- Melting, boiling, sublimation – energy is absorbed (endothermic) without a temperature change.
- Heating/cooling curves show plateaus at the melting point (solid → liquid) and boiling point (liquid → gas).
1.5 Pressure‑volume work (Supplement)
For a gas at constant temperature (isothermal): pV = constant. Work done = pΔV (positive when the gas expands).
2. Atoms, Elements & Compounds
2.1 Elements, compounds and mixtures
- Elements – made of one type of atom.
- Compounds – atoms of different elements chemically combined in fixed ratios.
- Mixtures – physical combinations; can be separated by physical methods.
2.2 Atomic structure
- Protons (+), neutrons (neutral), electrons (–).
- Mass number A = protons + neutrons; atomic number Z = protons.
- Isotopes – same Z, different A (e.g. ¹²C, ¹³C, ¹⁴C).
- Electronic configuration follows the Aufbau principle; valence electrons determine reactivity.
2.3 Ions
- Cations – loss of electrons (e.g. Na⁺, Fe²⁺).
- Anions – gain of electrons (e.g. Cl⁻, SO₄²⁻).
- Charge = number of electrons lost or gained.
2.4 Bonding
2.4.1 Ionic bonding
- Transfer of electrons from a metal to a non‑metal.
- Forms a giant lattice of oppositely charged ions (e.g. NaCl, MgO).
- High melting points, soluble in water, conduct electricity when molten or in solution.
2.4.2 Covalent bonding
- Sharing of electrons between non‑metals.
- Molecular covalent – discrete molecules (e.g. H₂O, CO₂). Low melting points, poor conductors.
- Giant covalent – continuous network (e.g. diamond, SiO₂). Very high melting points, hard, do not conduct electricity.
2.4.3 Metallic bonding
- Positive metal ions in a sea of delocalised electrons.
- Explains conductivity, malleability and ductility of metals.
2.4.4 Dot‑and‑cross diagrams
Show how valence electrons are shared or transferred. Example for H₂O:
H • • O • • • H
2.5 Writing chemical formulae
- Use oxidation numbers or valency to balance total charge (e.g. CaSO₄, Al₂O₃).
- Polyatomic ions retain their charge when incorporated (e.g. NaNO₃, CaCO₃).
3. Stoichiometry
3.1 Relative masses
- Relative atomic mass (Ar) – from the periodic table.
- Relative molecular mass (Mr) – sum of Ar of all atoms in a molecule.
3.2 Mole concept
1 mol = 6.02 × 10²³ particles (Avogadro’s number).
- Mass (g) = moles × Mr.
- Moles = mass ÷ Mr.
- Number of particles = moles × 6.02 × 10²³.
3.3 Common calculations
| Task | Formula | Example |
|---|
| Mass ↔ moles |
n = m / Mr |
12 g C → n = 12 g / 12 g mol⁻¹ = 1 mol |
| Moles ↔ molecules |
Number = n × Nₐ |
0.5 mol H₂O → 0.5 × 6.02×10²³ = 3.01×10²³ molecules |
| Mass ↔ mass (using a balanced equation) |
Use stoichiometric coefficients |
2 Mg + O₂ → 2 MgO; 24 g Mg produce 32 g MgO |
| Limiting reactant & theoretical yield |
Compare moles available with stoichiometric ratios |
40 g C + 40 g O₂ → CO₂; C is limiting, theoretical CO₂ = 40 g C × (44 g CO₂ / 12 g C) = 147 g |
| Percentage yield |
(actual yield ÷ theoretical yield) × 100 % |
Actual 130 g CO₂ → % yield = (130 ÷ 147) × 100 % = 88 % |
| Empirical & molecular formulae |
Convert % → mass → moles, simplify to smallest whole‑number ratio (empirical); then use Mr to find molecular formula. |
C₈H₁₈ (octane) – empirical C₄H₉, Mr = 114 g mol⁻¹ → (114 ÷ 57) = 2 → C₈H₁₈ |
4. Chemical Energetics
4.1 Exothermic & endothermic reactions
- Exothermic – ΔH < 0; heat released to surroundings (e.g. combustion of methane).
- Endothermic – ΔH > 0; heat absorbed (e.g. thermal decomposition of CaCO₃).
4.2 Bond energy
Energy required to break a bond (kJ mol⁻¹). Approximate enthalpy change:
ΔH ≈ Σ (bond energies broken) – Σ (bond energies formed)
4.3 Activation energy (Eₐ) and catalysts
- Eₐ = minimum energy that reacting particles must have.
- Catalysts provide an alternative pathway with lower Eₐ, increasing the rate without being consumed.
5. Chemical Reactions
5.1 Types of reactions (Core)
- Synthesis – A + B → AB
- Decomposition – AB → A + B
- Single‑replacement – A + BC → AC + B
- Double‑replacement – AB + CD → AD + CB
- Combustion – fuel + O₂ → CO₂ + H₂O (usually exothermic)
5.2 Reaction rates (Core)
- Factors: concentration, temperature, surface area, catalyst.
- Collision theory – particles must collide with sufficient energy and proper orientation.
5.3 Reversible reactions & Le Chatelier’s principle (Core)
- Dynamic equilibrium: forward and reverse rates equal.
- Changing concentration, pressure, temperature or adding a catalyst shifts the equilibrium position.
5.4 Redox reactions – simultaneous oxidation and reduction (Core + Supplement)
Definition
A redox (reduction‑oxidation) reaction involves the loss of electrons by one species (oxidation) and the gain of electrons by another (reduction) occurring at the same time.
Key concepts
- Oxidation – loss of electrons or increase in oxidation number.
- Reduction – gain of electrons or decrease in oxidation number.
- Oxidising agent – substance that is reduced (accepts electrons).
- Reducing agent – substance that is oxidised (donates electrons).
Oxidation‑number rules (Supplement)
- Elements in their elemental form have ON = 0.
- For mono‑atomic ions, ON = ionic charge.
- Hydrogen is +1 (except in metal hydrides where it is –1).
- Oxygen is –2 (except in peroxides –1, and in OF₂ where it is +2).
- The sum of ONs in a neutral compound is 0; in an ion it equals the overall charge.
| Element / Ion | Typical oxidation number | Example |
|---|
| H | +1 (–1 in metal hydrides) | H₂O, NaH |
| O | –2 (–1 in peroxides) | H₂O, H₂O₂ |
| Al | +3 | Al³⁺ |
| Cl | –1 (can be +1, +5, +7 in oxy‑anions) | Cl⁻, ClO₃⁻ |
| Fe | +2 or +3 | Fe²⁺, Fe³⁺ |
Identifying a redox reaction
- Write the formula of each reactant and product.
- Assign oxidation numbers to every atom.
- Atoms whose oxidation number increases are oxidised; those that decrease are reduced.
- Check that total increase = total decrease (electrons conserved).
Worked examples
- Combustion of magnesium
$$\mathrm{2Mg(s) + O_2(g) \rightarrow 2MgO(s)}$$
Mg: 0 → +2 (oxidation) O: 0 → –2 (reduction)
- Zinc reacting with copper(II) sulphate
$$\mathrm{Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)}$$
Zn: 0 → +2 (oxidation) Cu: +2 → 0 (reduction)
- Acidic dissolution of iron
$$\mathrm{Fe(s) + 2H^{+}(aq) \rightarrow Fe^{2+}(aq) + H_2(g)}$$
Fe: 0 → +2 (oxidation) H: +1 → 0 (reduction)
Balancing redox equations – Half‑reaction method (Supplement)
- Separate the overall reaction into oxidation and reduction half‑reactions.
- Balance all atoms except O and H.
- Balance O by adding H₂O; balance H by adding H⁺ (acidic) or OH⁻ (basic).
- Balance charge by adding electrons (e⁻).
- Multiply the half‑reactions so that the number of electrons cancelled is the same.
- Add the half‑reactions and cancel species that appear on both sides.
Example (acidic medium):
$$\mathrm{MnO_4^- + Fe^{2+} \rightarrow Mn^{2+} + Fe^{3+}}$$
- Oxidation: $\mathrm{Fe^{2+} \rightarrow Fe^{3+} + e^-}$
- Reduction: $\mathrm{MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O}$
- Multiply oxidation by 5, then add.
- Balanced equation: $\mathrm{5Fe^{2+} + MnO_4^- + 8H^+ \rightarrow 5Fe^{3+} + Mn^{2+} + 4H_2O}$
6. Acids, Bases & Salts
6.1 Acids
- Taste sour, turn blue litmus red, produce H⁺ in water.
- Strong acids (complete dissociation): HCl, H₂SO₄, HNO₃, HBr, HI.
- Weak acids (partial dissociation): CH₃COOH, H₂CO₃, H₂S.
6.2 Bases
- Feel slippery, turn red litmus blue, produce OH⁻ in water.
- Strong bases: NaOH, KOH, Ca(OH)₂ (soluble).
- Weak bases: NH₃, Al(OH)₃ (insoluble).
6.3 pH scale
$$pH = -\log[H^+]$$
- Acidic: pH < 7; Neutral: pH = 7; Basic: pH > 7.
6.4 Neutralisation
Acid + base → salt + water.
Example: $$\mathrm{HCl + NaOH \rightarrow NaCl + H_2O}$$
6.5 Indicators
- Natural: litmus, phenolphthalein, methyl orange.
- Colour change occurs at characteristic pH ranges.
6.6 Salt preparation
- By neutralisation (acid + base).
- By precipitation (mixing two aqueous solutions to form an insoluble salt).
- By reaction of an acid with a metal or metal oxide (e.g. 2HCl + Zn → ZnCl₂ + H₂).
7. Electrochemistry
7.1 Electrolysis of molten salts
- Cations travel to the cathode and gain electrons (reduction).
- Anions travel to the anode and lose electrons (oxidation).
- Example: $$\mathrm{2NaCl(l) \rightarrow 2Na(l) + Cl_2(g)}$$
7.2 Electrolysis of aqueous solutions
- Water can be oxidised (O₂) or reduced (H₂); the ion that is more easily discharged takes precedence.
- e.g. Electrolysis of NaCl solution → H₂ at cathode, Cl₂ at anode (because Cl⁻ is discharged more readily than OH⁻).
7.3 Electrochemical cells
- Galvanic (voltaic) cell – spontaneous redox reaction produces electricity.
- Electrolytic cell – non‑spontaneous reaction driven by external electricity.
7.4 Fuel cell (hydrogen‑oxygen)
Overall reaction: $$\mathrm{2H_2 + O_2 \rightarrow 2H_2O}$$
Produces electricity, water and heat; no combustion gases.
8. Metals
8.1 Reactivity series
K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > (H) > Cu > Ag > Au.
8.2 Extraction methods
- Highly reactive metals (e.g. Al, Na) – extracted by electrolysis of molten ore.
- Less reactive metals (e.g. Fe, Cu) – extracted by reduction with carbon or carbon monoxide.
8.3 Corrosion
- Oxidation of metals, usually in presence of water and oxygen (e.g. Fe → Fe₂O₃·nH₂O).
- Prevention: coating, galvanisation (zinc coating), cathodic protection.
8.4 Alloys
- Mixtures of two or more metals (or metal + non‑metal) with improved properties.
- Examples: Brass (Cu + Zn), Bronze (Cu + Sn), Steel (Fe + C).
9. The Periodic Table
| Group |
Typical properties |
Trend down a group |
Trend across a period (left → right) |
| 1 (alkali metals) |
Soft, low melting points, very reactive |
Increase in atomic radius, decrease in ionisation energy, increase in metallic character |
Decrease in atomic radius, increase in ionisation energy, decrease in metallic character |
| 2 (alkaline earth metals) |
Harder than alkali metals, form +2 ions |
Same trends as Group 1 |
Same trends as Group 1 |
| 13–16 (p‑block) |
Varied: metals, metalloids, non‑metals |
Metallic character increases down the group |
Metallic character decreases across the period; ionisation energy and electronegativity increase |
| 17 (halogens) |
Non‑metals, diatomic gases, form –1 ions |
Increase in atomic radius, decrease in electronegativity |
Decrease in atomic radius, increase in electronegativity, increase in reactivity up to fluorine |
| 18 (noble gases) |
Inert, complete valence shells |
Increase in atomic radius, decrease in ionisation energy |
All have full shells; no obvious trend in chemical reactivity (all very low) |
9.1 Periodic trends (key concepts)
- Atomic radius – decreases across a period, increases down a group.
- Ionisation energy – energy required to remove the outermost electron; opposite trend to atomic radius.
- Electronegativity – tendency to attract electrons; increases across a period, decreases down a group.
- Metallic ↔ non‑metallic character – metals on the left/bottom, non‑metals on the right/top.