Cambridge Syllabus Notes

Chemistry – Core IGCSE (0620) Revision Notes

Learning Objectives

Identify and describe the main features of the Cambridge IGCSE Chemistry syllabus.
Interpret reaction‑pathway (energy‑profile) diagrams and determine whether a reaction is exothermic or endothermic.
Explain the role of bond breaking and bond making in the energetics of a reaction.
Carry out simple quantitative calculations using bond‑energy data.
Recall key facts from the other core topics (states of matter, atomic structure, stoichiometry, etc.) so that the energetics section can be linked to the whole curriculum.

1. States of Matter

Key Concepts

Solids – fixed shape and volume; particles vibrate in fixed positions.
Liquids – fixed volume but take the shape of the container; particles slide past one another.
Gases – no fixed shape or volume; particles move rapidly and are far apart.
Diffusion & Effusion – spontaneous spreading of particles; rate increases with temperature and decreases with particle mass.
Kinetic‑Particle Theory – explains pressure, temperature and changes of state.

Practical Connections

Observing the temperature change when a gas condenses or a solid melts provides a simple illustration of ΔH (latent heat).

2. Atoms, Elements & Compounds

Atomic Structure (Core)

Protons (+), neutrons (0), electrons (‑).
Atomic number = number of protons; determines the element.
Relative atomic mass (Ar) – weighted average of isotopes.
Electronic configuration for the first 20 elements (1s, 2s, 2p, 3s, 3p).

Ions & Isotopes (Core)

Ions form by loss (cations) or gain (anions) of electrons.
Isotopes have the same Z but different neutron numbers; important for calculating relative atomic mass.

Bonding (Core)

Ionic bonding – transfer of electrons; giant lattice; high melting points; soluble in water.
Covalent bonding – sharing of electrons; represented by dot‑and‑cross diagrams; molecular compounds have low melting points.
Metallic bonding – sea of delocalised electrons; explains conductivity, malleability and ductility.

Supplementary Box – Giant Structures

Brief notes on silica (SiO₂), alumina (Al₂O₃) and diamond (C) for A‑Level extension.

3. Stoichiometry (Core)

Writing Formulas

Use valency/charge balance for ionic compounds.
Use the “Covalent‑bond” method for molecular compounds (e.g., CO₂, NH₃).

Relative Masses

Calculate relative formula mass (M) by adding Ar values of constituent atoms.
Use M to convert between mass and number of formula units (no mole concept required for core).

Supplement – The Mole (A‑Level)

Avogadro’s constant (6.02 × 10²³) and the relationship 1 mol = M g for optional extension.

4. Electrochemistry (Core)

Electrolysis

Definition: using electricity to drive a non‑spontaneous redox reaction.
Key terms: anode (oxidation, +ve), cathode (reduction, ‑ve), electrolyte.
Typical products:
- Molten NaCl → Na (cathode) + Cl₂ (anode)
- Molten PbBr₂ → Pb (cathode) + Br₂ (anode)
- Dilute H₂SO₄ (aq) → H₂ (cathode) + O₂ (anode)

Fuel Cells (Optional)

Hydrogen‑oxygen fuel cell: 2 H₂ + O₂ → 2 H₂O (electrical energy produced, ΔH < 0).

5. Chemical Energetics – Exothermic & Endothermic Reactions

Key Definitions

Enthalpy change, ΔH – heat change at constant pressure (kJ mol⁻¹).
Exothermic reaction – ΔH < 0; heat released to surroundings.
Endothermic reaction – ΔH > 0; heat absorbed from surroundings.
Activation energy, E_a – minimum energy needed to reach the transition state.
Transition state (activated complex) – highest‑energy arrangement along the reaction pathway.

Bond‑Breaking & Bond‑Making

Breaking a bond: endothermic (energy absorbed).
Forming a bond: exothermic (energy released).

Approximate enthalpy change:

ΔH ≈ Σ E(bonds broken) – Σ E(bonds formed)

Values are average bond energies (kJ mol⁻¹) from a standard table.

Components of a Reaction‑Pathway (Energy‑Profile) Diagram

Reaction coordinate – horizontal axis representing progress from reactants to products.
Energy (potential) – vertical axis.
Reactants – energy level at the left‑hand end.
Products – energy level at the right‑hand end.
Peak (Transition State) – highest point; vertical distance from reactants = E_a.
ΔH – vertical difference between reactant and product levels (negative = exothermic, positive = endothermic).
Catalyst effect (optional) – lowers the peak height (E_a) without changing ΔH.

Exothermic vs. Endothermic – Quick Comparison

Feature	Exothermic	Endothermic
Sign of ΔH	ΔH < 0 (heat released)	ΔH > 0 (heat absorbed)
Energy level of products	Lower than reactants	Higher than reactants
Typical example	Combustion of methane CH₄ + 2 O₂ → CO₂ + 2 H₂O	Thermal decomposition of calcium carbonate CaCO₃ → CaO + CO₂
Effect on surroundings	Temperature rises	Temperature falls

Step‑by‑Step Guide to Interpreting a Diagram

Locate the reactant energy level (left side).
Identify the highest point – the transition state. Measure the vertical distance from the reactants; this is E_a.
Locate the product energy level (right side).
Determine ΔH:
- If product level < reactant level → ΔH < 0 → exothermic.
- If product level > reactant level → ΔH > 0 → endothermic.
Look for additional features:
- Multiple peaks = multi‑step reaction.
- A lowered peak in the presence of a catalyst = reduced E_a (rate ↑, ΔH unchanged).

Quantitative Example – Combustion of Methane (Bond Energies)

Reaction (simplified): CH₄ + 2 O₂ → CO₂ + 2 H₂O

Bond	Number	Energy (kJ mol⁻¹)
C–H	4	413
O=O	2	498
C=O (in CO₂)	2	799
O–H (in H₂O)	4	463

Calculation:

Energy required to break bonds:
   4 × C–H = 4 × 413 = 1652 kJ
   2 × O=O = 2 × 498 =  996 kJ
   Total broken = 2648 kJ

Energy released on forming bonds:
   2 × C=O = 2 × 799 = 1598 kJ
   4 × O–H = 4 × 463 = 1852 kJ
   Total formed = 3450 kJ

ΔH ≈ Σ broken – Σ formed = 2648 – 3450 = –802 kJ mol⁻¹

The negative sign confirms an exothermic reaction.

6. Chemical Reactions – Rate, Equilibrium & Redox (Core)

Rate of Reaction

Factors that increase rate: higher temperature, higher concentration, larger surface area, presence of a catalyst.
Rate expressed qualitatively (fast/slow) or by a simple proportionality: rate ∝ concentration.

Reversible Reactions & Equilibrium

Represented by a double arrow (⇌).
Dynamic equilibrium – forward and reverse rates are equal.
Le Chatelier’s principle (A‑Level extension) – not required for core but useful for understanding.

Redox Identification (Core)

Oxidation = loss of electrons (or increase in oxidation number).
Reduction = gain of electrons (or decrease in oxidation number).
Typical redox reactions in the syllabus:
- Combustion (C, H, S oxidised to CO₂, H₂O, SO₂).
- Metal + acid → H₂ (metal oxidised, H⁺ reduced).

7. Acids, Bases & Salts (Core)

Definitions

Acid – substance that produces H⁺ ions in aqueous solution.
Base – substance that produces OH⁻ ions in aqueous solution.
Salt – product of an acid–base neutralisation.

Characteristic Reactions

Acid + metal → salt + H₂ (e.g., Zn + 2 HCl → ZnCl₂ + H₂).
Acid + carbonate → salt + CO₂ + H₂O (e.g., CaCO₃ + 2 HCl → CaCl₂ + CO₂ + H₂O).
Base + acid → salt + H₂O (neutralisation).

Indicators & pH (Supplement)

Phenolphthalein (colourless → pink in basic solution) and litmus (red ↔ blue). The pH scale (0–14) is an A‑Level addition.

8. The Periodic Table (Core)

Organisation

Groups (vertical) – elements have similar chemical properties.
Periods (horizontal) – properties change progressively.

Key Groups for IGCSE

Group	Typical Properties
1 (alkali metals)	Very reactive, form +1 ions, react vigorously with water.
7 (halogens)	Very reactive non‑metals, form –1 ions, diatomic gases (F₂, Cl₂, Br₂, I₂).
18 (noble gases)	Very low reactivity, complete outer shell.
Transition metals (Groups 3–12)	Variable oxidation states, form coloured compounds, good conductors.

Periodic Trends (Supplement)

Atomic radius, ionisation energy and electronegativity – useful for predicting bond type and reaction behaviour.

9. Metals (Core)

Physical & Chemical Properties

Conduct heat and electricity, malleable, ductile, shiny.
React with acids (produce H₂), with oxygen (form oxides), and with water (alkali metals only).

Extracting Metals

From ores – reduction with carbon (e.g., Fe₂O₃ + 3 C → 2 Fe + 3 CO₂) or electrolysis (e.g., aluminium).
Re‑refining – using electrolysis to purify (e.g., Cu²⁺ + 2 e⁻ → Cu).

10. The Environment (Core)

Key Topics

Acid rain – formation from SO₂ and NOₓ; effect on stone and water bodies.
Ozone depletion – CFCs, UV radiation.
Greenhouse gases – CO₂, CH₄, water vapour; global warming.
Waste treatment – incineration, recycling, landfill.

11. Organic Chemistry (Core)

Simple Hydrocarbons

Alkanes – general formula CₙH₂ₙ₊₂, single bonds only, combustion produces CO₂ + H₂O.
Alkenes – CₙH₂ₙ, at least one C=C double bond; addition reactions (e.g., H₂ addition).
Alcohols – contain –OH group; general formula CₙH₂ₙ₊₂O.

Functional‑Group Identification (A‑Level supplement)

Use of structural formulas and IUPAC naming basics.

12. Experimental Techniques (Core)

Common Skills

Measuring mass, volume and temperature accurately.
Using a Bunsen burner safely.
Filtration, crystallisation, distillation (basic).
Collecting gases over water and applying the correction for water vapour pressure.

Data Handling

Drawing clear, labelled diagrams (including energy‑profile diagrams).
Presenting results in tables and using appropriate units.
Identifying sources of error and suggesting improvements.

13. Interpreting Reaction‑Pathway Diagrams – Integrated Practice

Practice Question 1

A diagram shows reactants at 120 kJ mol⁻¹, the transition‑state peak at 190 kJ mol⁻¹, and products at 80 kJ mol⁻¹.

Calculate the activation energy E_a.
Calculate ΔH.
State whether the reaction is exothermic or endothermic.

Answers:

E_a = 190 – 120 = 70 kJ mol⁻¹.
ΔH = 80 – 120 = –40 kJ mol⁻¹.
ΔH < 0 → exothermic.

Practice Question 2

Explain how the diagram in Question 1 would change if a catalyst were added.

The height of the peak (E_a) would be reduced (e.g., from 70 kJ mol⁻¹ to a lower value), but the reactant and product energy levels remain unchanged, so ΔH stays –40 kJ mol⁻¹.

Practice Question 3

The dissolution of ammonium nitrate in water feels cold. Predict the sign of ΔH and sketch a brief energy‑profile description.

ΔH > 0 (endothermic) – energy is taken from the surroundings, causing a temperature drop.
Sketch: reactants at a lower level, a modest peak (E_a), products at a higher level.

Practice Question 4 – Bond‑Energy Calculation

Calculate ΔH for H₂ + Cl₂ → 2 HCl using:

H–H = 436 kJ mol⁻¹
Cl–Cl = 243 kJ mol⁻¹
H–Cl = 431 kJ mol⁻¹

Solution:

Energy to break bonds:
   H–H  = 436 kJ
   Cl–Cl = 243 kJ
   Total broken = 679 kJ

Energy released on forming bonds:
   2 × H–Cl = 2 × 431 = 862 kJ

ΔH = Σ broken – Σ formed = 679 – 862 = –183 kJ mol⁻¹

The negative value confirms an exothermic reaction.

Summary

These notes cover every core topic of the Cambridge IGCSE Chemistry (0620) syllabus, with a particular focus on interpreting reaction‑pathway diagrams. By linking the visual energy‑profile to bond‑breaking/forming, ΔH and E_a, students can:

Distinguish exothermic from endothermic processes.
Predict the temperature change of the surroundings.
Estimate enthalpy changes using bond‑energy data.
Understand how a catalyst influences rate without altering ΔH.
Integrate this knowledge with the broader chemistry curriculum (states of matter, atomic structure, stoichiometry, etc.).

Mastery of these ideas prepares learners for the IGCSE examination and provides a solid foundation for further A‑Level study.

Interpret reaction pathway diagrams showing exothermic and endothermic reactions