Cambridge IGCSE Chemistry (0620) – Full Syllabus Notes

1 States of Matter (Core 1.1‑1.4)

1.1 Three States of Matter

• Solids – definite shape & volume; particles vibrate in fixed positions; usually high density.
• Liquids – definite volume only; particles slide past one another; moderate density.
• Gases – no fixed shape or volume; particles move freely and are far apart; low density.

1.2 Changes of State

1.3 Diffusion (Core 1.2)

• Diffusion = movement of particles from an area of high concentration to low concentration.
• Rate increases with higher temperature, lower particle mass and greater concentration gradient.
• Example: Perfume spreading in a room.

1.4 Gas Laws (Core 1.4)

• Ideal‑gas equation: PV = nRT
• Boyle’s law (T constant): P ∝ 1/V
• Charles’s law (P constant): V ∝ T
• Gay‑Lussac’s law (V constant): P ∝ T

These relationships can be derived from the kinetic‑particle theory, which assumes that gas particles are in constant random motion and that collisions are perfectly elastic.

2 Atoms, Elements & Compounds (Core 2.1‑2.7)

2.1 Atomic Structure

• Protons (+) and neutrons (0) in the nucleus; electrons (–) in shells.
• Atomic number (Z) = number of protons = number of electrons in a neutral atom.
• Mass number (A) = protons + neutrons.

2.2 Isotopes

• Atoms of the same element with different numbers of neutrons.
• Example: ¹²C (6 p, 6 n) and ¹⁴C (6 p, 8 n).

2.3 Ions

• Cations – loss of electrons, positive charge.
• Anions – gain of electrons, negative charge.
• Ion formation is governed by the desire to achieve a noble‑gas electron configuration.

2.4 Bonding

Ionic Bonding

• Transfer of electrons from a metal to a non‑metal.
• Forms oppositely charged ions that attract each other (e.g., NaCl).

Covalent Bonding

• Sharing of electrons between non‑metals.
• Non‑polar: equal sharing (e.g., H₂).
• Polar: unequal sharing (e.g., H₂O).

Metallic Bonding (expanded)

• Valence electrons are delocalised over a lattice of positively charged metal ions – the “sea of electrons”.
• Explains conductivity, malleability, ductility and luster of metals.

Giant Covalent Structures

• Diamond – each C atom tetrahedrally bonded to four others; very hard, high melting point.
• Graphite – layers of C atoms in hexagonal sheets; layers held by weak forces, giving softness and good conductivity along the sheets.
• Silicon dioxide (SiO₂) – each Si atom bonded to four O atoms in a three‑dimensional network; high melting point, poor conductor.

3 Stoichiometry (Core 3.1‑3.2)

3.1 Mole Concept

• 1 mol = 6.022 × 10²³ particles (Avogadro’s constant).
• Molar mass (M) = atomic/molecular mass in g mol⁻¹.
• Number of moles: n = m / M.

3.2 Molar Volume of a Gas

At r.t.p. (25 °C, 1 atm) 1 mol of any gas occupies ≈ 24 dm³.

3.3 Writing Chemical Formulas

• Balance total positive and negative charges using oxidation numbers/valency.
• Example: Al³⁺ + O²⁻ → Al₂O₃ (2 Al³⁺ give 6 +, 3 O²⁻ give 6 –).

3.4 Calculations

Limiting‑Reactant & Theoretical Yield

1. Convert masses of reactants to moles (n = m/M).
2. Use the stoichiometric coefficients to compare mole ratios and identify the limiting reactant.
3. Calculate moles of product from the limiting reactant, then convert to mass.

Empirical & Molecular Formulae

• Convert % composition (or mass) to moles, divide by the smallest number of moles, and round to the nearest whole number – gives the empirical formula.
• Determine the molecular formula by comparing the empirical‑formula mass with the molar mass (M): n = M / M_empirical, then multiply the empirical subscripts by n.

Percentage Yield & Purity

Percentage yield = (actual yield ÷ theoretical yield) × 100 %.

Purity = (mass of pure substance ÷ total mass) × 100 %.

Worked Example (Limiting Reactant)

Reaction: 2 Mg + O₂ → 2 MgO

Given: 3.0 g Mg and 2.0 g O₂.

M(Mg)=24.3 g mol⁻¹, M(O₂)=32.0 g mol⁻¹.

Moles Mg = 3.0 ÷ 24.3 = 0.124 mol.

Moles O₂ = 2.0 ÷ 32.0 = 0.0625 mol.

Stoichiometry requires 2 mol Mg per 1 mol O₂ → required Mg = 2 × 0.0625 = 0.125 mol.

Mg is slightly short, so Mg is the limiting reactant.

Theoretical MgO = 0.124 mol × (2 mol MgO / 2 mol Mg) = 0.124 mol → mass = 0.124 × 40.3 = 5.0 g MgO.

4 Electrochemistry (Core 4.1‑4.2)

4.1 Electrolysis

• Passage of electricity through a molten or aqueous electrolyte to drive a non‑spontaneous redox reaction.
• Electrodes:

  • Anode – positive, oxidation occurs.
  • Cathode – negative, reduction occurs.

Predicting Products

Half‑Equation Method (step‑by‑step)

1. Write separate oxidation and reduction half‑equations.
2. Balance each half‑equation for mass and charge (add H₂O, H⁺, OH⁻ as needed).
3. Equalise the number of electrons transferred.
4. Add the half‑equations to obtain the overall reaction.

4.2 Fuel Cells

• Spontaneous redox reaction that produces electricity.
• Typical example: H₂ + ½ O₂ → H₂O (ΔE° > 0).

5 Chemical Energetics (Core 5.1)

5.1 Exothermic & Endothermic Reactions

• Exothermic – energy released; ΔH < 0; surroundings become warmer.
• Endothermic – energy absorbed; ΔH > 0; surroundings become cooler.

5.2 Enthalpy Change Symbols

• ΔH_f – heat of formation (elements → 1 mol of compound).
• ΔH_c – heat of combustion (1 mol of fuel reacts with O₂).

5.3 Activation Energy (Ea)

• Minimum energy required for reactant particles to collide successfully.
• Higher Ea → slower reaction (all else equal).

5.4 Bond‑Energy Calculations (simplified)

ΔH ≈ Σ(bond energies broken) – Σ(bond energies formed).

Example: H₂ + Cl₂ → 2 HCl

Bonds broken: H–H (436 kJ mol⁻¹) + Cl–Cl (243 kJ mol⁻¹) = 679 kJ.

Bonds formed: 2 × H–Cl (431 kJ mol⁻¹) = 862 kJ.

ΔH ≈ 679 – 862 = –183 kJ (exothermic).

5.5 Reaction‑Pathway Diagram

The diagram shows reactants → transition state (Ea) → products; ΔH is the vertical difference between reactants and products.

6 Chemical Reactions (Core 6.1‑6.4)

6.1 Rate of Reaction

• Factors that increase rate: higher temperature, higher concentration, larger surface area, presence of a catalyst.
• Collision theory – reactions occur when particles collide with sufficient energy (≥ Ea) and proper orientation.

6.2 Reversible Reactions & Chemical Equilibrium

• When forward and reverse rates are equal, the system is at dynamic equilibrium.
• Equilibrium constant (K_c) expresses the ratio of product concentrations to reactant concentrations at equilibrium (for gases, K_p may be used).
• Le Chatelier’s principle predicts the shift when conditions change (concentration, pressure, temperature, catalysts).

Le Chatelier Example

For the Haber process: N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g) + heat.

Increasing pressure favours the side with fewer gas moles (right side).

Removing NH₃ continuously drives the reaction to the right, increasing yield.

6.3 Industrial Processes (Core 6.4)

6.4 Redox Notation

• Oxidation – loss of electrons (increase in oxidation number).
• Reduction – gain of electrons (decrease in oxidation number).
• Balancing redox equations using the half‑reaction method (see Section 4.1).

7 Acids, Bases & Salts (Core 7.1‑7.3)

7.1 Properties

• Acids – produce H⁺ (or H₃O⁺) in water; taste sour; turn blue litmus red; react with metals → H₂.
• Bases – produce OH⁻ in water; feel slippery; turn red litmus blue; neutralise acids.

7.2 pH Scale

• pH = –log[H⁺].
• Strong acids (e.g., HCl) dissociate completely → pH ≈ 0–1 (1 M).
• Weak acids (e.g., CH₃COOH) only partially dissociate → higher pH for the same concentration.
• Strong bases (e.g., NaOH) dissociate completely; weak bases (e.g., NH₃) do not.

7.3 Neutralisation

Acid + Base → Salt + Water.

Example: H₂SO₄ + 2 NaOH → Na₂SO₄ + 2 H₂O.

7.4 Preparation of Salts

• Acid + metal oxide → salt + water.
• Acid + metal carbonate → salt + CO₂ + water.
• Acid + metal → salt + hydrogen gas.
• Acid + base (neutralisation) → salt + water.

7.5 Solubility Rules (Core 7.2)

7.6 Strong vs. Weak Acids & Bases

• Strong acids: HCl, HBr, HI, H₂SO₄, HNO₃ – complete ionisation.
• Weak acids: CH₃COOH, H₂CO₃, H₃PO₄ – partial ionisation.
• Strong bases: NaOH, KOH, Ca(OH)₂ – complete ionisation.
• Weak bases: NH₃, Al(OH)₃ – partial ionisation.

8 The Periodic Table (Core 8.1‑8.5)

8.1 General Layout

• Groups = vertical columns (same number of valence electrons).
• Periods = horizontal rows (increasing principal quantum number).
• Blocks: s‑block (Groups 1‑2, He), p‑block (Groups 13‑18), d‑block (transition metals), f‑block (lanthanides & actinides).

8.2 Trends Across a Period

8.3 Trends Down a Group

8.4 Transition Metals (Core 8.4)

• Variable oxidation states (e.g., Fe²⁺ / Fe³⁺, Cu⁺ / Cu²⁺).
• Typical properties: coloured compounds, good conductors, high melting points.
• Form complex ions such as [Fe(CN)₆]⁴⁻.

8.5 Relationship Between Group Number and Ion Charge (Main‑Group Elements)

Elements tend to lose or gain electrons to achieve a noble‑gas configuration. The number of electrons transferred is directly linked to the group (valence‑electron) number.

How to Predict the Ion Charge

1. Identify the element’s group number.
2. Metals (Groups 1, 2, 13): lose all valence electrons → positive charge equals the number of electrons lost.
3. Non‑metals (Groups 15‑17): gain enough electrons to reach eight valence electrons → negative charge = 8 – group‑valence‑electrons.
4. Write the ion symbol with the appropriate superscript.

Worked Examples

Na (Group 1): loses 1 e⁻ → Na⁺ (charge +1).

Cl (Group 17): gains 1 e⁻ → Cl⁻ (charge –1).

Al (Group 13): loses 3 e⁻ → Al³⁺ (charge +3).

O (Group 16): gains 2 e⁻ → O²⁻ (charge –2).

Common Misconceptions

• “All elements in a group form the same ion charge.” – Transition metals can exhibit several oxidation states; the simple pattern applies only to the main‑group elements listed above.
• “Higher group number always means a more negative ion.” – True for the non‑metal groups (15‑17) but opposite for the metal groups (1, 2, 13).

9 Metals (Core 9.1‑9.4)

9.1 Physical Properties

• Good conductors of heat and electricity (due to metallic bonding).
• Malleable and ductile – can be hammered or drawn into wires.
• Shiny (metallic luster) and usually solid at r.t.p.

9.2 Reactivity Series

Ordered from most to least reactive (typical series):

K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > ( H ) > Cu > Ag > Au

• More reactive metals displace less reactive ones from aqueous solutions of their salts.
• Reactivity decreases down a group for the s‑block metals.

9.3 Extraction of Metals

• Highly reactive metals (e.g., Na, Mg) – extracted by electrolysis of molten salts.
• Moderately reactive metals (e.g., Fe, Zn) – extracted by reduction with carbon (in a blast furnace) or by electrolysis of aqueous solutions.
• Less reactive metals (e.g., Cu, Ag, Au) – extracted by refining processes such as electro‑refining or by using less reactive reagents.

9.4 Alloys & Uses

• Alloys are mixtures of two or more metals (or a metal and a non‑metal) that have improved properties (e.g., steel = Fe + C, brass = Cu + Zn).
• Common uses: wiring (Cu), construction (steel), jewellery (Ag, Au), batteries (Zn, Cd).