IGCSE Chemistry (0620) – Complete Revision Notes
Learning Objective
Describe the change from metallic to non‑metallic character across a period and understand how this trend integrates with the wider Cambridge IGCSE Chemistry syllabus.
1. States of Matter
- Solid, liquid, gas – particle arrangement, kinetic energy and intermolecular forces.
- Changes of state – melting, boiling, sublimation; always endothermic (energy absorbed).
- Gas laws (core) – pV = nRT. Qualitative trends:
- Increasing temperature → volume ↑ (at constant pressure).
- Increasing pressure → volume ↓ (at constant temperature).
- Diffusion (core & supplement) – spontaneous mixing of gases or vapours; rate increases with temperature and decreases with molecular mass.
- Key formulae –
p = F/A, V = nRT/p, ρ = m/V.
2. Atoms, Elements & Compounds
2.1 Atomic Structure (core)
- Protons (+), neutrons (0), electrons (–). Atomic number (Z) = number of protons.
- Isotopes – same Z, different mass number (A). Example: 12C, 13C.
- Effective nuclear charge (Zeff) = Z – S (shielding). Increases across a period.
2.2 Electronic Configuration (core, 1–20)
| Element | Z | Configuration |
|---|
| H | 1 | 1s¹ |
| He | 2 | 1s² |
| Li | 3 | 1s² 2s¹ |
| Be | 4 | 1s² 2s² |
| B | 5 | 1s² 2s² 2p¹ |
| C | 6 | 1s² 2s² 2p² |
| N | 7 | 1s² 2s² 2p³ |
| O | 8 | 1s² 2s² 2p⁴ |
| F | 9 | 1s² 2s² 2p⁵ |
| Ne | 10 | 1s² 2s² 2p⁶ |
2.3 Mixtures, Elements & Compounds (core)
- Elements – made of one type of atom; cannot be broken down chemically.
- Compounds – two or more elements chemically combined in fixed ratios; can be broken down by chemical reactions.
- Mixtures – physical combinations of two or more substances; components retain their own properties and can be separated physically.
2.4 Ions and Bonding (core)
- Ion formation – metals lose electrons → cations; non‑metals gain electrons → anions.
- Ionic bonding – electrostatic attraction between oppositely charged ions (e.g., NaCl).
- Covalent bonding – sharing of electron pairs; represented by dot‑and‑cross diagrams (e.g., H₂O).
- Metallic bonding – delocalised “sea of electrons” surrounding a lattice of metal cations (giant‑metallic structure). Explains conductivity, malleability and high melting points.
- Alloys (core) – mixtures of two or more metals (or a metal and a non‑metal) that form a single phase, e.g., brass (Cu + Zn), steel (Fe + C).
- Formula writing – empirical vs. molecular formulas; e.g., glucose: empirical C₆H₁₂O₆, molecular C₆H₁₂O₆ (same in this case).
3. Stoichiometry
- Relative atomic mass (Ar) and relative molecular mass (Mr).
- Mole concept (core) – 1 mol = 6.02 × 10²³ particles; n = m/Mr.
- Mass‑mass, mass‑mole, mole‑mole calculations; limiting‑reactant and percentage‑yield (supplement).
- Empirical & molecular formulas (supplement) – derived from percent composition or combustion analysis.
- “Mole‑free” shortcuts for common compounds (e.g., 1 g H ≈ 8 g O in water).
4. Electrochemistry
5. Chemical Energetics
- Exothermic vs. endothermic – energy released (ΔH < 0) or absorbed (ΔH > 0).
- Enthalpy change (ΔH) – measured by calorimetry; sign convention must be stated explicitly.
- Bond‑energy diagram – breaking bonds requires energy; forming bonds releases energy.
- Approximate calculation:
ΔH ≈ Σ(bond energies broken) – Σ(bond energies formed).
- Activation energy (supplement) – minimum energy required for a reaction to proceed; explains effect of temperature on rate.
6. Chemical Reactions
6.1 Physical vs. Chemical Change (core)
- Physical change – state or form changes, composition unchanged (e.g., melting ice).
- Chemical change – new substances formed, composition altered (e.g., rusting of iron).
6.2 Rate of Reaction (core)
- Factors: concentration, temperature, surface area, catalyst.
- Collision theory – effective collisions required; activation energy is the barrier.
6.3 Equilibrium (core)
- Reversible reactions: A + B ⇌ C + D.
- Dynamic equilibrium – forward and reverse rates equal.
- Le Chatelier’s principle – predicts shift when concentration, temperature, pressure or catalyst changes.
- Typical reversible examples: CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O, NH₄Cl(s) ⇌ NH₃(g) + HCl(g).
6.4 Redox (core & supplement)
- Oxidation – loss of electrons, increase in oxidation number.
- Reduction – gain of electrons, decrease in oxidation number.
- Redox can be identified by:
- Change in oxidation numbers.
- Transfer of electrons in half‑equations.
- Example:
2 Mg + O₂ → 2 MgO
Mg: 0 → +2 (oxidised); O: 0 → –2 (reduced).
7. Acids, Bases & Salts
- Acids – sour taste, turn blue litmus red, produce H⁺ in water (e.g., HCl, H₂SO₄).
- Strong acids – dissociate completely (HCl, H₂SO₄, HNO₃, HBr, HI).
- Weak acids – only partially dissociate (CH₃COOH, H₂CO₃).
- Bases – slippery feel, turn red litmus blue, produce OH⁻ (e.g., NaOH, Ca(OH)₂).
- Strong bases – soluble metal hydroxides (NaOH, KOH, Ca(OH)₂).
- Weak bases – ammonia (NH₃) and amines.
- Neutralisation – acid + base → salt + water (e.g., H₂SO₄ + 2 NaOH → Na₂SO₄ + 2 H₂O).
- pH scale (0–14) – indicator colour chart (core):
| pH | Colour (Universal Indicator) |
|---|
| 0‑1 | Red |
| 2‑3 | Orange |
| 4‑5 | Yellow |
| 6‑7 | Green |
| 8‑9 | Blue‑green |
| 10‑11 | Blue |
| 12‑14 | Purple |
8. The Periodic Table – Layout & General Trends
- Groups (vertical) – elements in the same group have similar chemical properties.
- Periods (horizontal) – properties change progressively from left to right.
8.1 Trends Across a Period
- Atomic radius ↓ – electrons are pulled closer by increasing Zeff.
- Ionisation energy ↑ – more energy required to remove the outer electron.
- Electronegativity ↑ – stronger ability to attract bonding electrons.
- Metallic character ↓ – elements become less willing to lose electrons.
8.2 Trends Down a Group
- Atomic radius ↑ – additional electron shells.
- Ionisation energy ↓ – outer electrons are farther from the nucleus.
- Electronegativity ↓.
- Metallic character ↑ – greater tendency to lose electrons.
8.3 Group‑Specific Patterns (core)
- Group 1 (alkali metals) – always form +1 ions; very reactive, low IE and EN.
- Group 2 (alkaline earth metals) – form +2 ions; slightly less reactive than Group 1.
- Group 7 (halogens) – form –1 ions; high EN, strong oxidising agents.
- Group 18 (noble gases) – full valence shells, essentially inert.
- Transition metals (Groups 3‑12) – variable oxidation states, form coloured compounds, often act as catalysts.
9. Metallic → Non‑Metallic Character Across a Period
9.1 Why the Trend Occurs
- Electrons are added to the same principal energy level; Zeff felt by valence electrons increases.
- Higher Zeff pulls the electron cloud closer → atomic radius decreases.
- Smaller atoms hold outer electrons more tightly → first ionisation energy rises.
- Greater ability to attract electrons in a bond → electronegativity rises.
- Result: elements on the left readily lose electrons (metallic, reducing agents); those on the right readily gain electrons (non‑metallic, oxidising agents).
9.2 Illustrative Data – Period 2
| Group | Element | Metallic Character | Electronegativity (Pauling) | First IE (kJ mol⁻¹) | Typical Oxidation State(s) |
| 1 | Li | High | 0.98 | 520 | +1 |
| 2 | Be | Moderate | 1.57 | 900 | +2 |
| 13 | B | Low | 2.04 | 800 | +3, –3 |
| 14 | C | Very Low | 2.55 | 1086 | +4, –4 |
| 15 | N | Negligible | 3.04 | 1402 | –3, +5 |
| 16 | O | None | 3.44 | 1314 | –2 |
| 17 | F | None | 3.98 | 1681 | –1 |
9.3 Illustrative Data – Period 3 (selected)
| Group | Element | Metallic Character | Electronegativity | First IE (kJ mol⁻¹) | Typical Oxidation State(s) |
| 1 | Na | High | 0.93 | 496 | +1 |
| 2 | Mg | Moderate | 1.31 | 738 | +2 |
| 13 | Al | Low | 1.61 | 578 | +3 |
| 14 | Si | Very Low | 1.90 | 787 | +4, –4 |
| 15 | P | Negligible | 2.19 | 1012 | –3, +5 |
| 16 | S | None | 2.58 | 1000 | –2, +6 |
| 17 | Cl | None | 3.16 | 1251 | –1 |
9.4 Summary of the Trend
- Left‑hand elements (e.g., Li, Na) are strong reducing agents, form basic oxides and ionic compounds.
- Right‑hand elements (e.g., F, Cl) are strong oxidising agents, form acidic oxides and covalent compounds.
- Both ionisation energy and electronegativity increase together, reflecting the shift from electron‑donating to electron‑accepting behaviour.
9.5 Chemical Behaviour Implications
- Metals → metallic bonds, malleable, conductive, form cations; typical reactions:
- Metal + acid → salt + H₂ (e.g., Zn + 2 HCl → ZnCl₂ + H₂).
- Metal + non‑metal oxide → salt (e.g., 2 Mg + O₂ → 2 MgO).
- Non‑metals → covalent bonds, poor conductors, form anions or share electrons; typical reactions:
- Non‑metal + metal oxide → acid (e.g., C + 2 MgO → 2 MgCO₃ → CO₂ + H₂O).
- Halogen + metal → ionic halide (e.g., 2 Na + Cl₂ → 2 NaCl).
10. Metals – Properties, Extraction & Reactivity
- Physical properties – shiny, ductile, malleable, good conductors of heat and electricity.
- Reactivity series (core) – K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > (less reactive) Cu > Ag > Au.
- Extraction methods
- Electrolysis – e.g., Al from Al₂O₃ (Hall‑Héroult process).
- Reduction with carbon – e.g., Fe from Fe₂O₃ in a blast furnace.
- Distillation – e.g., Zn from ZnO (volatilisation and condensation).
- Corrosion – oxidation of metals, especially iron; prevention by coating, alloying, cathodic protection.
- Alloys (core) – single‑phase mixtures of metals (or metal + non‑metal) with improved properties; examples: brass (Cu + Zn), stainless steel (Fe + Cr + Ni).
11. The Environment
- Green chemistry principles – waste minimisation, atom economy, safer solvents, energy efficiency.
- Key pollutants – CO₂ (global warming), NOₓ (acid rain), SO₂ (acid rain), particulate matter.
- Water‑testing methods (core)
- Dry‑cobalt(II) chloride test – colour change (blue → pink) indicates presence of water vapour.
- Copper(II) sulphate test – blue crystals dissolve in water, indicating moisture.
- Conductivity test – pure water is a poor conductor; presence of ions (from dissolved salts) increases conductivity.
- Acid rain – formation of H₂SO₄ and HNO₃ from SO₂ and NOₓ; impacts on buildings, soils and aquatic life.
- Ozone depletion – catalytic destruction by CFCs; relevance to the chemistry of halogens.