IGCSE Chemistry – Core Topics (Cambridge 0620)
Learning Objectives
- Describe the states of matter and the kinetic‑particle theory.
- Explain atomic structure, isotopes, ions and the main types of bonding.
- Carry out stoichiometric calculations using formulae, relative masses and the mole concept.
- Understand electrolysis, half‑equations and the principle of fuel cells.
- Distinguish exothermic and endothermic processes and use bond‑energy ideas.
- Analyse factors that affect reaction rate, equilibrium and redox behaviour.
- Identify properties of acids, bases and salts and apply pH concepts.
- Use the position of an element in the Periodic Table to predict its physical and chemical properties.
- Describe the properties, uses, extraction and corrosion of metals.
- Explain key environmental chemistry topics (water treatment, fertilizers, air quality, climate change).
1. States of Matter
Key Concepts
- Solids – fixed shape and volume; particles vibrate in fixed positions; usually high density.
- Liquids – fixed volume but take the shape of their container; particles slide past each other; moderate density.
- Gases – no fixed shape or volume; particles move rapidly in all directions; low density.
- Diffusion – spontaneous mixing of gases or liquids; faster at higher temperature and lower particle mass.
Kinetic‑Particle Theory (KPT)
| State | Particle Arrangement | Particle Motion | Inter‑particle Forces |
|---|
| Solid | Close‑packed, ordered | Vibrations about fixed points | Strong |
| Liquid | Close‑packed, unordered | Sliding past one another | Moderate |
| Gas | Widely spaced, unordered | Free, rapid motion | Very weak |
Practical Points
- Melting point (solid → liquid) and boiling point (liquid → gas) are characteristic of each substance.
- Density (ρ) = mass/volume; solids > liquids > gases (generally).
2. Atoms, Elements & Compounds
2.1 Atomic Structure
- Atoms consist of a nucleus (protons + neutrons) surrounded by electrons in shells.
- Atomic number (Z) = number of protons = number of electrons in a neutral atom.
- Mass number (A) = protons + neutrons.
- Isotopes: atoms of the same element (same Z) with different A.
2.2 Ions
- Cations: loss of electrons → positive charge (e.g., Na⁺, Al³⁺).
- Anions: gain of electrons → negative charge (e.g., Cl⁻, O²⁻).
- Ion formation follows the “octet rule” (main‑group elements) and the group‑charge relationship (see Periodic Table section).
2.3 Types of Bonding
Ionic Bonding
- Transfer of electrons from a metal to a non‑metal.
- Produces a lattice of oppositely charged ions.
- High melting/boiling points, soluble in water, conduct electricity when molten or in solution.
Covalent (Molecular) Bonding
- Sharing of electron pairs between non‑metals.
- Simple molecules (e.g., H₂O, CO₂) have low melting/boiling points; do not conduct electricity.
Giant Covalent Structures
- Continuous network of covalent bonds (e.g., diamond, SiO₂, graphite).
- Very high melting points; hardness varies (diamond hard, graphite slippery).
Metallic Bonding
- Delocalised “sea of electrons” around positively charged metal ions.
- Properties: conductivity, malleability, ductility, luster, variable melting points.
2.4 Representative Formulas
| Compound Type | General Formula | Example |
|---|
| Ionic | Metal + Non‑metal (charges balanced) | NaCl, CaO |
| Molecular | Non‑metal + Non‑metal (often with prefixes) | CO₂, NH₃ |
| Giant Covalent | Elemental network | SiO₂, C (diamond) |
| Metallic | Metal only | Fe, Al |
3. Stoichiometry
3.1 Relative Atomic Mass (Ar) & Relative Molecular Mass (Mr)
- Ar = weighted average of isotopic masses (relative to ¹²C = 12.00).
- Mr = sum of Ar values for all atoms in a formula.
3.2 The Mole Concept
- 1 mol = 6.022 × 10²³ particles (Avogadro’s number).
- Mass of 1 mol = molar mass (g mol⁻¹) = Mr (g mol⁻¹) for compounds.
3.3 Calculations
- Convert mass ↔ moles using molar mass.
- Use balanced equations to relate moles of reactants ↔ products.
- Convert moles ↔ number of particles if required.
Example
Calculate the mass of CaO produced when 5.0 g of CaCO₃ decomposes:
CaCO₃(s) → CaO(s) + CO₂(g)
M(CaCO₃) = 100.1 g mol⁻¹
M(CaO) = 56.1 g mol⁻¹
5.0 g CaCO₃ × (1 mol / 100.1 g) = 0.050 mol CaCO₃
0.050 mol CaCO₃ → 0.050 mol CaO
0.050 mol × 56.1 g mol⁻¹ = 2.8 g CaO
3.4 Limiting Reactant & Percent Yield
- Identify the reactant that produces the fewest moles of product (limiting).
- Percent yield = (actual yield / theoretical yield) × 100 %.
4. Electrochemistry
4.1 Electrolytic Cells
- Non‑spontaneous reactions driven by an external electric current.
- Components: anode (positive, oxidation), cathode (negative, reduction), electrolyte, power source.
- Half‑equations are written for each electrode.
Example – Electrolysis of Molten NaCl
Anode: 2Cl⁻ → Cl₂(g) + 2e⁻ (oxidation)
Cathode: Na⁺ + e⁻ → Na(l) (reduction)
Overall: 2NaCl(l) → 2Na(l) + Cl₂(g)
4.2 Predicting Products
- Identify the cation – it is reduced at the cathode (metal or H⁺).
- Identify the anion – it is oxidised at the anode (non‑metal or O²⁻).
- Consider the nature of the electrolyte (aqueous vs. molten) – water can be reduced/oxidised if the ion is inert.
4.3 Fuel Cells
- Generate electricity from a spontaneous redox reaction (e.g., H₂ + ½O₂ → H₂O).
- Two half‑cells separated by a porous membrane; electrons travel through an external circuit.
5. Chemical Energetics
5.1 Exothermic & Endothermic Reactions
- Exothermic: energy released (ΔH < 0); temperature of surroundings rises.
- Endothermic: energy absorbed (ΔH > 0); temperature of surroundings falls.
5.2 Activation Energy (Ea)
- Minimum energy required for reactants to form an activated complex.
- Catalysts lower Ea, increasing the rate without being consumed.
5.3 Bond Energy Approach
ΔH ≈ Σ (bond energies broken) – Σ (bond energies formed).
- Breaking bonds requires energy; forming bonds releases energy.
- Useful for estimating the enthalpy change of a reaction.
6. Chemical Reactions
6.1 Reaction Rate
- Rate ∝ frequency of effective collisions.
- Factors: concentration, temperature, surface area, catalysts, nature of reactants.
6.2 Reversible Reactions & Equilibrium
- Both forward and reverse reactions occur simultaneously.
- Dynamic equilibrium: rates equal, concentrations remain constant.
- Le Chatelier’s principle predicts the shift when conditions change (concentration, pressure, temperature).
6.3 Redox Reactions
- Oxidation = loss of electrons; reduction = gain of electrons.
- Oxidising agent gains electrons; reducing agent loses electrons.
- Use oxidation numbers to identify electron transfer.
Example – Reaction of Magnesium with Hydrochloric Acid
Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Oxidation: Mg → Mg²⁺ + 2e⁻
Reduction: 2H⁺ + 2e⁻ → H₂(g)
7. Acids, Bases & Salts
7.1 Acid and Base Definitions
- Arrhenius: Acid → H⁺ in water; Base → OH⁻ in water.
- Bronsted‑Lowry: Acid donates a proton; Base accepts a proton.
7.2 pH Scale
- pH = –log[H⁺].
- pH < 7 = acidic; pH = 7 = neutral; pH > 7 = basic.
7.3 Neutralisation
Acid + Base → Salt + Water
Example: H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O
7.4 Salt Preparation & Solubility Rules
| Rule | Typical Outcome |
|---|
| All nitrates (NO₃⁻) are soluble | e.g., NaNO₃, KNO₃ |
| All alkali‑metal salts are soluble | e.g., NaCl, K₂SO₄ |
| Most sulphates are soluble except CaSO₄, SrSO₄, BaSO₄ | |
| Most carbonates, phosphates, hydroxides are insoluble except with alkali metals | |
7.5 Oxides
- Acidic oxides (e.g., SO₂, CO₂) react with water to give acids.
- Basic oxides (e.g., CaO, MgO) react with water to give bases.
8. The Periodic Table – Predicting Properties from Position
8.1 Structure of the Table
- Ordered by increasing atomic number (Z).
- Periods – 7 horizontal rows (Period 1 → 7).
- Groups – 18 vertical columns (Group 1 → 18). Elements in a group have the same number of valence electrons.
- Blocks – s‑block (Groups 1‑2, He), p‑block (Groups 13‑18), d‑block (Transition metals, Groups 3‑12), f‑block (Lanthanides & Actinides).
8.2 Valence‑Electron Configuration
| Group | Valence‑electron configuration |
|---|
| 1 (alkali metals) | ns¹ |
| 2 (alkaline earth) | ns² |
| 13 | ns²np¹ |
| 14 | ns²np² |
| 15 | ns²np³ |
| 16 | ns²np⁴ |
| 17 (halogens) | ns²np⁵ |
| 18 (noble gases) | ns²np⁶ (except He: 1s²) |
8.3 Trends Across a Period (Left → Right)
| Property | Trend | Reason |
|---|
| Atomic radius | Decreases | Increasing nuclear charge pulls electrons closer. |
| Ionisation energy | Increases | Electrons are held more tightly. |
| Electronegativity | Increases | Greater ability to attract bonding electrons. |
| Metallic character | Decreases | Loss of electrons becomes less favourable. |
8.4 Trends Down a Group (Top → Bottom)
| Property | Trend | Reason |
|---|
| Atomic radius | Increases | Additional electron shells are added. |
| Ionisation energy | Decreases | Outer electrons are farther from the nucleus and shielded. |
| Electronegativity | Decreases | Weaker pull on bonding electrons. |
| Metallic character | Increases | Atoms more readily lose electrons as size grows. |
8.5 Group‑Charge Relationship (Main‑Group Elements)
| Group | Typical Ion Charge | Examples |
|---|
| 1 (alkali) | +1 | Na⁺, K⁺ |
| 2 (alkaline earth) | +2 | Mg²⁺, Ca²⁺ |
| 13 | +3 (or +1 for some metals) | Al³⁺ |
| 14 | ±4, ±2, 0 | Si⁴⁺, C⁴⁻ |
| 15 | −3 (or +5, +3) | P³⁻, N³⁻ |
| 16 | −2 | O²⁻, S²⁻ |
| 17 (halogens) | −1 | Cl⁻, Br⁻ |
| 18 (noble gases) | Generally no charge | He, Ne, Ar |
8.6 Transition Metals (d‑Block)
- Variable oxidation states (e.g., Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺).
- Often form coloured ions and complex ions.
- Metallic character is high; generally good conductors and malleable.
- Trends are less regular than for the main‑group; however, atomic radius still increases down the group.
8.7 Noble Gases
- Full valence shells → very low reactivity.
- High ionisation energies; used as inert atmospheres (e.g., Ar in welding).
- He has the highest ionisation energy of all elements.
8.8 Predicting Properties – Quick Checklist
- Identify the element’s period → gives an idea of size and shielding.
- Identify the group → gives valence‑electron count, typical ion charge, metallic/non‑metallic character.
- Determine the block → predicts possible oxidation states (especially for transition metals).
- Apply trends:
- Left‑most groups → large radius, low IE, low EN, strong metallic character.
- Right‑most groups → small radius, high IE, high EN, non‑metallic.
- Moving down → larger radius, lower IE/EN, more metallic.
Example Prediction
Predict the properties of bromine (Br, Group 17, Period 4):
- Non‑metal, diatomic (X₂) at room temperature.
- Relatively high electronegativity (≈2.8) and ionisation energy (≈114 kJ mol⁻¹), but lower than chlorine (above it).
- Forms –1 anion (Br⁻) in ionic compounds.
- Reacts vigorously with alkali metals and alkaline earth metals.
9. Metals
9.1 General Properties
- Shiny, ductile, malleable, good conductors of heat and electricity.
- High melting/boiling points (except alkali metals).
- Typically form basic oxides and hydroxides.
9.2 Reactivity Series
Ordered from most to least reactive (common series):
K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > (H) > Cu > Ag > Au
- Metals above hydrogen displace H⁺ from acids; those above a metal in the series displace that metal from its salts.
9.3 Extraction of Metals
| Metal | Typical Extraction Method |
|---|
| Aluminium (Group 13) | Electrolysis of molten Al₂O₃ (Hall–Héroult process) |
| Iron (Group 8) | Reduction of Fe₂O₃ with CO in a blast furnace |
| Copper (Group 11) | Electrolytic refining; also smelting of Cu₂S with O₂ |
| Gold (Group 11) | Crude ore is treated with cyanide solution (cyanidation) and then electrolysis |
9.4 Corrosion & Protection
- Corrosion = oxidation of a metal, most commonly iron → rust (Fe₂O₃·nH₂O).
- Prevention methods: painting, galvanising (zinc coating), cathodic protection, alloying (e.g., stainless steel).
9.5 Alloys
- Mixture of two or more metals (or a metal and a non‑metal) that has useful properties.
- Examples: brass (Cu + Zn), bronze (Cu + Sn), steel (Fe + C), stainless steel (Fe + Cr + Ni).
10. Chemistry of the Environment
10.1 Water Chemistry
- Hard water contains Ca²⁺ and Mg²⁺; softened by ion‑exchange or adding washing soda (Na₂CO₃).
- Water testing: pH, conductivity, dissolved oxygen, turbidity.
- Treatment methods – coagulation & flocculation, filtration, chlorination, UV disinfection.
10.2 Fertilisers
- Sources of N, P, K (NPK). Common compounds: ammonium nitrate (NH₄NO₃), superphosphate (Ca(H₂PO₄)₂), potassium chloride (KCl).
- Over‑use leads to eutrophication (excess nutrients → algal blooms → oxygen depletion).
10.3 Air Quality
- Major pollutants: sulphur dioxide (SO₂), nitrogen oxides (NOₓ), carbon monoxide (CO), particulate matter (PM).
- Acid rain formation: SO₂ + H₂O → H₂SO₃; NOₓ + H₂O → HNO₃.
- Control measures: flue‑gas desulphurisation, catalytic converters, low‑sulphur fuels.
10.4 Climate Change
- Greenhouse gases (CO₂, CH₄, N₂O) trap infrared radiation.
- Human activities (burning fossil fuels, deforestation) increase atmospheric CO₂.
- Mitigation: renewable energy, energy efficiency, carbon capture, reforestation.
Practice Questions
- Predict which of the following elements will have the highest first ionisation energy: Na, Mg, Al, Si.
Answer: Si (right‑most in the same period).
- Write the balanced equation for the electrolysis of aqueous NaCl and list the gases produced at each electrode.
Answer: 2NaCl(aq) → 2Na(l) + Cl₂(g) (anode). At the cathode, water is reduced: 2H₂O + 2e⁻ → H₂(g) + 2OH⁻.
- Calculate the mass of CO₂ formed when 10.0 g of CH₄ combusts completely.
Solution: CH₄ + 2O₂ → CO₂ + 2H₂O; M(CH₄)=16.0 g mol⁻¹, M(CO₂)=44.0 g mol⁻¹. 10.0 g CH₄ × (1 mol/16.0 g)=0.625 mol CH₄ → 0.625 mol CO₂ × 44.0 g mol⁻¹ = 27.5 g CO₂.
- Explain why the reactivity of alkali metals increases down Group 1.
Answer: Atomic radius increases, outer electron is farther from the nucleus and more shielded, so ionisation energy decreases, making electron loss easier.
- Identify the oxidation numbers of Fe in Fe₂O₃ and in FeSO₄ and state the change in oxidation state when Fe₂O₃ is reduced to FeSO₄.
Answer: In Fe₂O₃, Fe = +3 (since O = –2). In FeSO₄, Fe = +2 (SO₄²⁻ has –2 overall). Change: +3 → +2 (gain of one electron per Fe atom).
Summary
The Periodic Table is a powerful predictive tool. By locating an element’s period (row) and group (column) you can anticipate its:
- Atomic size and shielding
- Ionisation energy and electronegativity
- Metallic or non‑metallic character
- Typical ion charge and common types of compounds
- Behaviour in redox reactions and electrochemical processes
These predictions underpin the understanding of bonding, reactivity, extraction, environmental impact and the many practical applications covered throughout the IGCSE Chemistry syllabus.