Edexcel IGCSE Chemistry · Topics 1–4 — Revision Pack

1. Outline

Topic 1 Principles of chemistry builds the language of chemistry: particles, atoms, bonding and calculations.

2. States, elements & mixtures

Particle model

• Solid: particles tightly packed in a regular lattice; vibrate about fixed positions; fixed shape and volume.
• Liquid: particles close but randomly arranged; can slide past each other; fixed volume, takes shape of container.
• Gas: particles far apart and move randomly at high speed; fill the container; easily compressed.

• Melting / freezing at melting point; boiling / condensing at boiling point — temperature stays constant during a pure change of state while energy goes into breaking (or forming) forces between particles.
• Diffusion: particles spread from high to low concentration; faster in gases than liquids; faster for lighter particles and at higher temperature.

Elements, compounds, mixtures

• Element: substance made of only one type of atom (e.g. O₂, Fe).
• Compound: two or more elements chemically bonded in fixed proportions (e.g. H₂O, NaCl).
• Mixture: two or more substances not chemically combined; composition can vary; can be separated by physical methods.

Separation techniques

3. Atomic structure & Periodic Table

Nucleus = protons + neutrons. Electrons in shells. Lithium electronic configuration 2,1.

Subatomic particles

• Atomic number (Z) = number of protons (= number of electrons in a neutral atom).
• Mass number (A) = protons + neutrons.
• Isotopes: same number of protons, different numbers of neutrons (same chemical properties; different mass / physical properties).

Electronic configuration

• Shells hold up to 2, 8, 8… (for IGCSE first 20 elements).
• Group number ≈ number of outer-shell electrons; period number = number of shells.
• Examples: Na = 2,8,1; Cl = 2,8,7; O = 2,6.

Periodic Table trends

• Metals (left / centre) → lose electrons → positive ions; non-metals (right) → gain / share electrons.
• Down a group: atoms get larger; outer electron further from nucleus → Group 1 more reactive; Group 7 less reactive.
• Across a period: metallic → non-metallic character.

4. Formulae, equations & moles

Writing and balancing equations

• Word equation → chemical equation with formulae → balance atoms on both sides (change coefficients, never subscripts).
• State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous.

Relative formula mass (M_r)

Sum of relative atomic masses (A_r) of all atoms in the formula. e.g. H₂O: 1+1+16 = 18; CaCO₃: 40+12+48 = 100.

The mole

One mole contains Avogadro’s number of particles (6.02 × 10²³). Concentrations:

Mass

g

M_r

g/mol

Amount n

mol

Concentration

mol/dm³

Empirical formula

1. Divide each mass (or %) by A_r → moles of each element.
2. Divide by the smallest number of moles → simplest whole-number ratio.

Percentage yield & molar gas volume (C)

• (C) At room temperature and pressure, 1 mole of any gas occupies about 24 dm³.
• (C) n = V / 24 (V in dm³) for gases at r.t.p.

5. Bonding

Ionic = transfer of electrons; covalent = sharing; metallic = lattice of positive ions in a sea of delocalised electrons.

Ionic bonding

• Metal + non-metal: electrons transferred → cations and anions → giant ionic lattice held by strong electrostatic attraction.
• High melting/boiling points; conduct when molten or aqueous (ions free to move); brittle; often soluble in water.

Covalent bonding

• Non-metals share pairs of electrons. Simple molecules (H₂O, CO₂, CH₄): strong bonds inside molecule, weak forces between → low m.p./b.p.; do not conduct.
• Giant covalent (diamond, graphite, SiO₂): very high m.p.; diamond hard insulator; graphite soft layers, conducts (delocalised electrons).

Metallic bonding

• Positive metal ions in a sea of delocalised electrons → strong attraction.
• Good conductors of heat and electricity; malleable and ductile; generally high m.p.

6. Electrolysis

• Electrolysis: splitting an ionic compound using electricity. Needs molten or aqueous ions (mobile charge).
• Anode (+): anions lose electrons (oxidation). Cathode (−): cations gain electrons (reduction).
• Remember OIL RIG: Oxidation Is Loss; Reduction Is Gain (of electrons).

Molten ionic compounds

e.g. molten PbBr₂: cathode → Pb; anode → Br₂.

Aqueous solutions (C)

• Water also provides H⁺ and OH⁻. Preferential discharge rules apply.
• At cathode: less reactive metal OR hydrogen if metal is more reactive than H.
• At anode: halide ions → halogen if concentrated; otherwise often O₂ from OH⁻.
• (C) Write ionic half-equations, e.g. Cu2+ + 2e− → Cu; 2Cl− → Cl2 + 2e−.

Uses

• Extraction of reactive metals (e.g. aluminium from molten Al₂O₃).
• Electroplating; purifying copper; producing chlorine and hydrogen from brine.

7. Self-check questions

Try these yourself first. Answers are in the next section.

1. How do particle arrangement and motion differ in a solid and a gas?
2. An atom has atomic number 11 and mass number 23. How many protons, neutrons and electrons does it have?
3. Write the electronic configuration of chlorine (Z = 17).
4. Calculate the M_r of Na₂CO₃ (Na=23, C=12, O=16).
5. How many moles are in 8.0 g of CH₄ (M_r = 16)?
6. Explain why ionic compounds conduct when molten but not when solid.
7. (C) A reaction’s theoretical yield is 12.0 g but only 9.0 g is obtained. Calculate % yield.
8. (C) During electrolysis of concentrated aqueous NaCl, what is produced at each electrode?

8. Self-check answers

1. Outline

2. Group 1 & Group 7

Group 1 — alkali metals (Li, Na, K…)

• Soft, low density, stored under oil; one outer electron → form M⁺ ions.
• React with water: 2M + 2H2O → 2MOH + H2 — alkaline solution + hydrogen (pops with lighted splint).
• Reactivity increases down the group (outer electron further / more shielded → lost more easily).
• Flame colours: Li crimson/red; Na yellow/orange; K lilac.

Group 7 — halogens (F, Cl, Br, I…)

• Diatomic molecules X₂; gain one electron → X⁻.
• Reactivity decreases down the group.
• Displacement: a more reactive halogen displaces a less reactive one from its salt solution, e.g. Cl₂ + 2KBr → 2KCl + Br₂ (solution turns orange/brown).

3. Gases in the atmosphere

• Approximate composition of dry air: ~78% nitrogen, ~21% oxygen, ~0.04% carbon dioxide, plus noble gases.
• Oxygen: relights a glowing splint.
• Carbon dioxide: turns limewater milky (cloudy white).
• Hydrogen: lighted splint → squeaky pop.
• Chlorine: bleaches damp litmus (blue → red then white).

(C) Know how oxygen and nitrogen can be separated from liquid air by fractional distillation (different boiling points).

4. Metals, reactivity & extraction

Use the series for displacement, extraction method, and reaction with water/acid.

Reactions linked to reactivity

• With water / steam / dilute acid: more reactive metals react more vigorously (produce H₂ with acid if above H).
• Displacement: more reactive metal displaces less reactive from a salt solution.

Extraction

• Metals more reactive than carbon → extracted by electrolysis of molten ore (e.g. Al from Al₂O₃).
• Metals less reactive than carbon → reduced by heating with carbon / carbon monoxide (e.g. Fe in blast furnace).
• Unreactive metals may occur native (e.g. Au).

Rusting & uses

• Iron rusts when both water and oxygen are present: hydrated iron(III) oxide.
• Prevention: barrier methods (paint, oil, plating), sacrificial protection (Zn), galvanising, alloys (stainless steel).
• Common uses linked to properties: Al (low density, oxide layer); Cu (conducts, ductile); Fe/steel (strong, cheap).

5. Acids, alkalis & salts

Acids and alkalis

• Acids in water → H⁺ (aq). Alkalis → OH⁻ (aq).
• pH scale 0–14: acid < 7, neutral = 7, alkali > 7. Universal indicator / pH meter.
• Neutralisation: H+ + OH− → H2O.

• Acid + metal → salt + hydrogen (if metal above H).
• Acid + base (metal oxide/hydroxide) → salt + water.
• Acid + carbonate → salt + water + carbon dioxide.

Titration (C)

• Accurate volumes of acid and alkali using pipette and burette; indicator (e.g. phenolphthalein / methyl orange).
• Use mean titre (concordant results) then n = c × V and mole ratios from the equation.

Preparing salts

• Soluble salt from insoluble base: add excess base to acid → filter → evaporate / crystallise.
• Soluble salt by titration (C): acid + alkali → evaporate.
• Insoluble salt: precipitation — mix two soluble salts → filter, wash, dry.

6. Chemical tests

Cation flame tests

Cations with NaOH(aq)

• Cu²⁺ → blue precipitate.
• Fe²⁺ → green ppt (goes brown on standing); Fe³⁺ → brown ppt.
• Al³⁺, Zn²⁺, Ca²⁺ → white ppt; Al³⁺ and Zn²⁺ dissolve in excess NaOH.
• NH₄⁺ → warm with NaOH → NH₃ (turns damp red litmus blue).

Anion tests

• Carbonate: acid → CO₂ (limewater milky).
• Halide: acidified AgNO₃ → Cl⁻ white ppt; Br⁻ cream; I⁻ yellow (check solubility in NH₃).
• Sulfate: acidified BaCl₂ → white BaSO₄ ppt.

7. Self-check questions

1. Why does potassium react more vigorously with water than lithium?
2. Write the ionic equation for chlorine displacing bromide ions.
3. State the approximate percentages of N₂ and O₂ in air.
4. Why is aluminium extracted by electrolysis rather than with carbon?
5. Name the salt formed from sulfuric acid and copper(II) oxide.
6. How do you test for a chloride ion?
7. (C) In a titration, 25.0 cm³ of 0.100 mol/dm³ NaOH needs 20.0 cm³ of HCl. Find the concentration of HCl.
8. What observations show Fe²⁺ with NaOH(aq)?

8. Self-check answers

1. Outline

2. Energetics

• Exothermic: heat given out to surroundings → temperature of surroundings rises; ΔH negative. Examples: combustion, neutralisation, respiration.
• Endothermic: heat taken in from surroundings → temperature falls; ΔH positive. Examples: thermal decomposition, photosynthesis, some dissolving.

• If more energy is released making bonds than is used breaking them → exothermic overall.
• (C) Use average bond energies from a table to calculate ΔH for a reaction.
• Energy level diagrams: reactants higher than products for exothermic; arrow down labelled −ΔH; activation energy from reactants up to peak.

3. Rates of reaction

Rate can be followed by: volume of gas vs time; mass loss vs time; colour change / turbidity (precipitation); concentration change.

Collision theory

• Particles must collide with enough energy (≥ activation energy) and correct orientation.
• Anything that increases successful collision frequency increases rate.

Typical practical: marble chips + dilute HCl → measure CO₂ volume or mass loss; compare chip size, concentration, temperature.

4. Reversible reactions & equilibria

• Reversible: ⇌ — products can reform reactants.
• Dynamic equilibrium (closed system): forward and reverse rates equal; concentrations constant (not necessarily equal).

Le Chatelier’s principle (qualitative)

• ↑ temperature → equilibrium shifts in the endothermic direction.
• ↑ pressure (gases) → shifts to side with fewer gas moles.
• ↑ concentration of a species → shifts to remove that species.
• Catalyst: reaches equilibrium faster; does not change position of equilibrium.

Industrial examples (C)

• Haber process: N2 + 3H2 ⇌ 2NH3 (exothermic). Compromise: moderate T (~450 °C), high P (~200 atm), iron catalyst.
• Contact process: SO₂ → SO₃ with V₂O₅ catalyst for sulfuric acid manufacture.

5. Self-check questions

1. Is combustion of methane exothermic or endothermic? What is the sign of ΔH?
2. Explain, using collision theory, why raising temperature increases rate.
3. Why does powdered Zn react faster with acid than a large lump of the same mass?
4. What does a catalyst do to activation energy and to the position of equilibrium?
5. For N2 + 3H2 ⇌ 2NH3 (exothermic), what happens to the yield of NH₃ if temperature is increased?
6. (C) Why is a compromise temperature used in the Haber process?
7. (C) Breaking bonds is endothermic; making bonds is exothermic. When is a reaction overall endothermic?

6. Self-check answers

1. Outline

2. Crude oil & alkanes

Fractions condense at different heights — shorter chains (lower b.p.) rise higher.

• Crude oil: mixture of hydrocarbons; separated by fractional distillation.
• Shorter chains: lower boiling point, more volatile, more flammable, runnier.
• Cracking: large alkanes → smaller alkanes + alkenes (heat + catalyst) to meet demand for fuels and monomers.

Alkanes

• General formula CnH2n+2; saturated (only C–C single bonds).
• Names: methane, ethane, propane, butane…
• Complete combustion → CO₂ + H₂O; incomplete → CO / C (soot) — CO is toxic.
• React with halogens in UV light (substitution) — one H replaced at a time.

3. Alkenes

• General formula CnH2n; unsaturated — contain C=C double bond.
• Test: bromine water orange/brown → colourless (addition across C=C). Alkanes do not decolourise rapidly in the dark.
• Addition reactions: with H₂ (Ni catalyst) → alkane; with H₂O (steam, catalyst) → alcohol; with halogens → dihaloalkane; polymerisation → poly(alkene).

4. Alcohols, carboxylic acids & esters (C)

Alcohols

• Functional group −OH; general formula C_nH_2n+1OH (e.g. ethanol C₂H₅OH).
• Ethanol: fermentation of sugars (yeast, ~30 °C, anaerobic) or hydration of ethene.
• Uses: fuels, solvents, drinks; complete combustion → CO₂ + H₂O.
• Oxidised to carboxylic acids (e.g. ethanol → ethanoic acid).

Carboxylic acids

• Functional group −COOH; e.g. ethanoic acid CH₃COOH (vinegar).
• Weak acids: pH typically 3–4 in dilute solution; react with metals, bases, carbonates like other acids.

Esters

• Alcohol + carboxylic acid ⇌ ester + water (acid catalyst, heat). Sweet / fruity smells.
• Naming idea: alkyl from alcohol + …oate from acid (e.g. ethanol + ethanoic acid → ethyl ethanoate).

5. Synthetic polymers

• Addition polymerisation: many alkene monomers join; C=C opens. e.g. ethene → poly(ethene). Repeat unit shown in brackets with n.
• Condensation polymerisation (C): monomers with two functional groups join, losing a small molecule (often H₂O). e.g. polyester, nylon.
• Uses: plastics packaging, fibres, ropes — linked to properties (flexible, strong, inert…).
• Disposal problems: non-biodegradable → landfill / litter; incineration → CO₂ / toxic gases; recycling preferred where possible.

6. Self-check questions

1. Why do different fractions separate in a fractionating column?
2. Give the general formulae of alkanes and alkenes.
3. How can you distinguish ethene from ethane using bromine water?
4. Write a word equation for the complete combustion of propane.
5. What is cracking and why is it useful?
6. (C) How is ethanol made by fermentation? State two conditions.
7. (C) Name the ester from ethanol and ethanoic acid.
8. What is the difference between addition and condensation polymerisation?

7. Self-check answers