CIE Cambridge O-Level · Paper 5070

Chemistry
Quick Study Notes

All 11 chapters · Balanced equations · State symbols · Reaction conditions · Diagrams · 30-question self-assessment quiz

⚗️ States of Matter ⚛️ Atoms & Bonding 🔢 Stoichiometry ⚡ Electrochemistry 🔥 Energetics 💥 Reactions & Rates 🧪 Acids & Bases 📊 Periodic Table 🔩 Metals 🌍 Environment 🔬 Practical Skills 📝 Take the Quiz
01
States of Matter
Particles, properties and changes of state

Solid

  • Fixed shape & volume
  • Particles vibrate in fixed positions
  • Strong forces between particles
  • Cannot be compressed

Liquid

  • Fixed volume, no fixed shape
  • Particles slide past each other
  • Weaker forces than solid
  • Flows; cannot be compressed

Gas

  • No fixed shape or volume
  • Particles move freely & rapidly
  • Very weak forces
  • Easily compressed

Particle Arrangement Diagrams

SOLID Fixed positions, vibrate only LIQUID Slide past each other, no fixed shape GAS Move freely, far apart, fill container ✦ Regular arrangement ✦ Irregular, close together ✦ Widely spaced, random

Changes of State

SOLID LIQUID GAS Melting ▲ Freezing ▼ Evaporation/Boiling ▲ Condensation ▼ Sublimation (e.g. iodine, dry ice) Deposition

Heating (energy absorbed)

  • Melting: solid → liquid (at melting point)
  • Boiling/Evaporation: liquid → gas (at boiling point)
  • Sublimation: solid → gas directly (e.g. iodine, CO₂)

Cooling (energy released)

  • Freezing: liquid → solid
  • Condensation: gas → liquid
  • Deposition: gas → solid directly
💡 Memory tip: During melting and boiling, temperature stays constant — energy breaks intermolecular bonds, not raising temperature.

Diffusion

Diffusion is the net movement of particles from a region of high concentration to low concentration. Gases diffuse faster than liquids. Lighter/smaller molecules diffuse faster (e.g. NH₃ diffuses faster than HCl).

Ammonia vs Hydrogen Chloride Diffusion
Experiment — white ring of NH₄Cl forms closer to HCl end
+
Room temperatureIn a long glass tube
NH₃(g) + HCl(g) → NH₄Cl(s)
White smoke/solid ammonium chloride forms where gases meet
  • NH₃ has Mr = 17; HCl has Mr = 36.5 → NH₃ travels further
  • White ring forms closer to HCl end of tube
  • Demonstrates relative speeds of diffusion

Diagram — NH₃ vs HCl Diffusion Tube

NH₃ HCl NH₃ diffuses (Mr = 17, faster) HCl diffuses (Mr = 36.5, slower) NH₄Cl (white) longer distance (NH₃ travels more) shorter distance (HCl travels less) Cotton wool soaked in NH₃ Cotton wool soaked in HCl open open White ring of NH₄Cl(s) forms closer to HCl end

Random, zigzag movement of small particles (e.g. smoke) caused by continuous collisions with invisible gas molecules. Evidence for the particle theory of matter.

02
Atoms, Elements & Compounds
Atomic structure, bonding and formulae
Key Definitions

Atom, Element, Compound, Mixture

Atom: smallest particle of an element that can exist.
Element: substance made of only one type of atom (e.g. O, Fe, Na).
Compound: two or more elements chemically bonded together (e.g. H₂O, NaCl).
Mixture: two or more substances mixed together but not chemically combined.

Atomic Structure

ParticleLocationRelative MassRelative Charge
ProtonNucleus1+1
NeutronNucleus10
ElectronShells (orbitals)negligible (1/1840)−1

Key Numbers

  • Atomic number (Z) = number of protons
  • Mass number (A) = protons + neutrons
  • Number of electrons = number of protons (neutral atom)
  • Isotopes: same Z, different A (different number of neutrons)

Electron Shells (2,8,8 rule)

  • Shell 1: max 2 electrons
  • Shell 2: max 8 electrons
  • Shell 3: max 8 electrons (O-level)
  • Na (Z=11): 2, 8, 1
  • Cl (Z=17): 2, 8, 7

Types of Bonding

Ionic Bonding
Metal + Non-metal → transfer of electrons
+
Na(s) → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
Na⁺Cl⁻ = NaCl(s) ionic compound in giant lattice
  • Metals lose electrons to form positive ions (cations)
  • Non-metals gain electrons to form negative ions (anions)
  • High melting points; conduct electricity when melted or dissolved
  • Soluble in water (usually)
Covalent Bonding
Non-metal + Non-metal → shared electrons
+
H₂O — 2 shared pairs; CH₄ — 4 shared pairs; CO₂ — 2 double bonds
  • Simple molecular: low m.p./b.p.; poor conductors
  • Giant covalent (e.g. diamond, SiO₂): very high m.p.; poor conductors
  • Diamond: every C bonded to 4 C atoms — hardest natural substance
  • Graphite: layers of hexagonal C rings; delocalised electrons → conducts
Metallic Bonding
Metal atoms → sea of delocalised electrons
+
  • Positive metal ions in sea of free electrons
  • Good conductors of heat and electricity
  • Malleable and ductile — layers of ions slide
  • High melting points (strong electrostatic attraction)

Common Ion Charges to Remember

IonFormulaIonFormula
AmmoniumNH₄⁺HydroxideOH⁻
NitrateNO₃⁻SulfateSO₄²⁻
CarbonateCO₃²⁻ChlorideCl⁻
PhosphatePO₄³⁻OxideO²⁻
ZincZn²⁺Iron(II)Fe²⁺
Iron(III)Fe³⁺Copper(II)Cu²⁺
03
Stoichiometry
The mole, calculations, empirical and molecular formulae
Core Formula

The Mole

1 mole = 6.02 × 10²³ particles (Avogadro's number)
Molar mass = mass of 1 mole in grams = relative atomic/formula mass in g/mol

Key Equations

  • n = m ÷ Mr
  • n = V(dm³) ÷ 24 (at r.t.p.)
  • n = c × V(dm³)
  • c = n ÷ V

Symbols

  • n = moles
  • m = mass (g)
  • Mr = molar mass
  • V = volume
  • c = concentration (mol/dm³)

Gas Volume

  • At r.t.p.: 1 mol gas = 24 dm³
  • At s.t.p.: 1 mol gas = 22.4 dm³
  • Convert cm³ ÷ 1000 = dm³

Empirical & Molecular Formula

Empirical formula: simplest whole number ratio of atoms.
Molecular formula: actual number of atoms in one molecule.
Molecular formula = n × empirical formula, where n = Mr(molecular) ÷ Mr(empirical)

ExampleGlucose: Empirical = CH₂O (Mr=30); Molecular = C₆H₁₂O₆ (Mr=180); n = 180/30 = 6

Percentage Yield & Purity

% Yield% yield = (actual yield ÷ theoretical yield) × 100
% Purity% purity = (mass of pure substance ÷ total mass) × 100

Mole Ratio from Equations

Using a Balanced Equation for Calculations
Step-by-step method
+
2H₂(g) + O₂(g) → 2H₂O(l)
2 mol H₂ reacts with 1 mol O₂ to give 2 mol H₂O
  • Step 1: Write balanced equation
  • Step 2: Find moles of known substance (n = m/Mr)
  • Step 3: Use ratio to find moles of unknown
  • Step 4: Convert moles to mass/volume as needed
💡 Concentration tip: mol/dm³ is the same as M (molar). Always convert cm³ to dm³ by dividing by 1000 before using c = n/V.

Titration Calculations

Titration formula pathc₁V₁ ÷ n₁ = c₂V₂ ÷ n₂  (where n₁,n₂ are the stoichiometric coefficients)
  • Use indicator to find end-point (e.g. phenolphthalein: colourless in acid, pink in alkali)
  • Methyl orange: red in acid, yellow in alkali
  • Average concordant titres (within 0.1 cm³ of each other)
04
Electrochemistry
Electrolysis, electrode reactions and applications
Key Concept

Electrolysis

Electrolysis is the decomposition of an ionic compound (when molten or dissolved) using electricity.
Cathode (−): cations (positive ions) gain electrons → reduced.
Anode (+): anions (negative ions) lose electrons → oxidised.

💡 Memory: OILRIG — Oxidation Is Loss, Reduction Is Gain (of electrons).  |  PANIC: Positive Anode, Negative Is Cathode.

Electrolysis of Molten Lead(II) Bromide

Molten PbBr₂
Simple ionic electrolysis
+
Heat until moltenDirect current (d.c.)
At cathode: Pb²⁺(l) + 2e⁻ → Pb(l)
At anode: 2Br⁻(l) → Br₂(g) + 2e⁻
Overall: PbBr₂(l) → Pb(l) + Br₂(g)
  • Grey liquid lead forms at cathode
  • Brown bromine gas produced at anode

Electrolysis of Dilute Sulfuric Acid

Dilute H₂SO₄(aq) — Inert electrodes
Aqueous electrolysis
+
Inert electrodes (platinum/graphite)
At cathode: 2H⁺(aq) + 2e⁻ → H₂(g)
At anode: 4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻
Overall: 2H₂O(l) → 2H₂(g) + O₂(g)
  • Volume of H₂ at cathode is TWICE volume of O₂ at anode (ratio 2:1)
  • H₂SO₄ acts as electrolyte only — not decomposed

Electrolysis of Concentrated NaCl(aq) — Brine

Brine (concentrated NaCl solution)
Industrial — Chlor-alkali process
+
Inert electrodes (graphite)Concentrated NaCl(aq)
At cathode (−): 2H⁺(aq) + 2e⁻ → H₂(g)
At anode (+): 2Cl⁻(aq) → Cl₂(g) + 2e⁻
Remaining in solution: Na⁺(aq) + OH⁻(aq) → NaOH(aq)
  • Three useful products: Cl₂ (anode), H₂ (cathode), NaOH (remaining solution)
  • Cl₂ is discharged at the anode (not O₂) because Cl⁻ concentration is high — high concentration favours Cl⁻ discharge over OH⁻
  • H⁺ is discharged at the cathode (not Na) because Na⁺ is above H in the reactivity series
  • Uses: Cl₂ → bleach, PVC; H₂ → margarine, fuel cells; NaOH → soap, paper
💡 Remember: In dilute NaCl, O₂ would be produced at the anode instead of Cl₂ — concentration matters!

Electroplating

Copper Electroplating
Application of electrolysis
+
CuSO₄(aq) electrolyte
Cathode (object to plate): Cu²⁺(aq) + 2e⁻ → Cu(s)
Anode (copper): Cu(s) → Cu²⁺(aq) + 2e⁻
  • Object to be plated = cathode
  • Metal to plate with = anode
  • Electrolyte = solution of the metal salt being used
  • Anode dissolves; cathode gains mass

Discharge Order in Aqueous Solutions

At Cathode (reduction)At Anode (oxidation — inert electrode)
Ions below H₂ in reactivity: Cu²⁺, Ag⁺ → discharged preferentiallyHalide ions (Cl⁻, Br⁻, I⁻) if concentrated → discharged preferentially
H⁺ discharged if metal ion is above H₂ (Na⁺, Mg²⁺, etc.)Otherwise OH⁻ → O₂ gas produced
05
Chemical Energetics
Exothermic, endothermic reactions and bond energies
Exothermic

Heat Released

Products have less energy than reactants. Temperature increases. ΔH is negative.
Examples: combustion, neutralisation, respiration, oxidation of metals

Endothermic

Heat Absorbed

Products have more energy than reactants. Temperature decreases. ΔH is positive.
Examples: thermal decomposition, photosynthesis, dissolving ammonium nitrate

Key Reactions with ΔH

Combustion of Methane
Exothermic — ΔH negative
+
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l)   ΔH = −890 kJ/mol
  • Complete combustion requires excess oxygen
  • Incomplete combustion gives CO(g) and C(s) (soot)
Thermal Decomposition of Calcium Carbonate
Endothermic — ΔH positive
+
Strong heat (>840°C)
CaCO₃(s) → CaO(s) + CO₂(g)   ΔH = +178 kJ/mol

Bond Energies

Breaking bonds requires energy (endothermic).
Making bonds releases energy (exothermic).

ΔH CalculationΔH = Σ(bonds broken) − Σ(bonds formed)
If positive → endothermic; if negative → exothermic

Activation Energy

The minimum energy required for a reaction to occur. A catalyst lowers the activation energy, providing an alternative pathway. The overall ΔH remains the same with or without a catalyst.

Energy Profile Diagrams

Exothermic Reaction (ΔH < 0) Reaction Progress → Energy Reactants Products Ea Ea (catalyst) −ΔH Transition state Endothermic Reaction (ΔH > 0) Reaction Progress → Energy Reactants Products Ea Ea (catalyst) +ΔH Transition state
💡 Catalyst reminder: The dashed green line shows how a catalyst lowers Ea without changing the energy levels of reactants or products — ΔH stays the same.
06
Chemical Reactions
Rate of reaction, reversible reactions and equilibrium

Factors Affecting Rate of Reaction

Temperature ↑

  • More energy → more collisions
  • Higher proportion of particles exceed activation energy

Concentration ↑

  • More particles per volume
  • More frequent collisions

Surface Area ↑

  • Smaller particles → bigger surface
  • More particles exposed to react

Catalyst

  • Provides alternative pathway
  • Lowers activation energy
  • Not used up in reaction

Light (some reactions)

  • Photochemical reactions
  • e.g. Cl₂ + CH₄ in UV light

Pressure (gases)

  • ↑ pressure → closer together
  • More frequent collisions

Important Reaction Examples

Reaction of Marble with HCl
Rate investigation — gas collection
+
CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
  • Measure CO₂ gas volume vs time
  • Steeper graph = faster rate
  • Powder reacts faster than lumps (greater surface area)
Decomposition of H₂O₂
Catalytic decomposition
+
MnO₂ catalyst
2H₂O₂(aq) → 2H₂O(l) + O₂(g)
  • MnO₂ is the catalyst — not consumed
  • O₂ relights a glowing splint (test)

Reversible Reactions & Equilibrium

Le Chatelier's Principle

Dynamic Equilibrium

In a closed system at equilibrium, forward and reverse reactions occur at equal rates. If conditions change, the equilibrium shifts to oppose the change.

Haber Process — Making Ammonia
Industrial reversible reaction
+
450°C200 atmIron catalyst
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)   ΔH = −92 kJ/mol
  • High pressure favours forward reaction (4 moles gas → 2 moles)
  • Low temperature favours more NH₃ but is too slow — 450°C is compromise
  • Iron catalyst speeds equilibrium attainment (does not shift position)
  • Yield ≈ 15% at 450°C, 200 atm (unreacted gases recycled)
Contact Process — Making Sulfuric Acid (Step 2)
Industrial process
+
450°C~2 atmV₂O₅ catalyst
2SO₂(g) + O₂(g) ⇌ 2SO₃(g)
SO₃(g) + H₂SO₄(l) → H₂S₂O₇(l) → diluted with water → 2H₂SO₄(l)
  • SO₃ NOT dissolved directly in water (produces acid mist) — dissolved in conc. H₂SO₄ first (oleum)
  • Step 1: S(s) + O₂(g) → SO₂(g)
07
Acids, Bases and Salts
pH, reactions and salt preparation
Acids

Proton Donors

Produce H⁺(aq) ions in water. pH < 7.
HCl — hydrochloric acid
H₂SO₄ — sulfuric acid
HNO₃ — nitric acid
CH₃COOH — ethanoic acid (weak)

Bases/Alkalis

Proton Acceptors

Alkalis produce OH⁻(aq) ions in water. pH > 7.
NaOH — sodium hydroxide
Ca(OH)₂ — calcium hydroxide
NH₃ — ammonia (weak)
CuO — base (not alkali)

💡 pH scale: 0–6 = acidic; 7 = neutral; 8–14 = alkaline. Each unit change in pH = 10× change in H⁺ concentration.

Acid Reactions — Key Equations

Acid + Metal
Salt + Hydrogen gas
+
Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)
  • H₂ gas — burns with a squeaky pop (lighted splint test)
  • Not all metals react — Cu, Ag, Au are below H₂ in reactivity series
Acid + Metal Oxide (Base)
Salt + Water (neutralisation)
+
CuO(s) + 2HCl(aq) → CuCl₂(aq) + H₂O(l)
MgO(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l)
Fe₂O₃(s) + 3H₂SO₄(aq) → Fe₂(SO₄)₃(aq) + 3H₂O(l)
Acid + Alkali (Neutralisation)
Salt + Water
+
HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l)
HNO₃(aq) + NaOH(aq) → NaNO₃(aq) + H₂O(l)
Ionic equation: H⁺(aq) + OH⁻(aq) → H₂O(l)
Acid + Carbonate
Salt + Water + Carbon Dioxide
+
CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
Na₂CO₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g)
NaHCO₃(aq) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g)
  • CO₂ turns limewater milky: Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)
Acid + Ammonia
Ammonium salt — no water
+
NH₃(g) + HCl(g) → NH₄Cl(s)
NH₃(aq) + HNO₃(aq) → NH₄NO₃(aq)
2NH₃(aq) + H₂SO₄(aq) → (NH₄)₂SO₄(aq)
  • NH₄NO₃ and (NH₄)₂SO₄ are important fertilisers

Salt Preparation Methods

Salt TypeMethodExample
Soluble (not Na/K/NH₄)Excess metal/metal oxide/carbonate + acid → filter, evaporateCuSO₄ from CuO + H₂SO₄
Na, K, NH₄ saltsTitration (alkali + acid, no excess)NaCl from NaOH + HCl
Insoluble saltsPrecipitation — mix two solutionsBaSO₄ from BaCl₂ + Na₂SO₄
Precipitation — Making Insoluble Salts
Mix two solutions
+
BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)
  • Filter → wash with distilled water → dry
  • BaSO₄ test: white ppt. with dilute HCl and BaCl₂ → identifies SO₄²⁻
08
The Periodic Table
Trends, groups and periods
Organisation

Arrangement of Elements

Elements are arranged in order of increasing atomic number. Vertical columns = Groups (same number of outer electrons → similar properties). Horizontal rows = Periods (same number of electron shells).

Periodic Trends (Left → Right across a period)

PropertyTrend across Period 3 (Na → Cl)
Atomic radiusDecreases (more protons pull electrons closer)
Ionisation energyGenerally increases
Metallic characterDecreases (Na, Mg, Al → Si → P, S, Cl)
ElectronegativityIncreases

Group 1 — Alkali Metals (Li, Na, K…)

Group 1 Metals with Water
Vigorous reaction — produces H₂ and alkali
+
2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g)
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)
  • Li — fizzes gently; Na — fizzes, melts; K — fizzes, burns lilac/purple flame
  • Reactivity increases DOWN the group (outer electron further from nucleus)
  • All have 1 outer electron; form M⁺ ions

Group 7 — Halogens (F, Cl, Br, I)

Halogen Displacement Reactions
More reactive halogen displaces less reactive one
+
Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)  [solution turns orange-brown]
Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq)   [solution turns brown/black]
Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq)   [solution turns brown]
  • Reactivity DECREASES down Group 7 (larger atom, outer electrons further/more shielded)
  • Colours: F₂ pale yellow gas; Cl₂ yellow-green gas; Br₂ red-brown liquid; I₂ grey-black solid
  • Halogens form −1 ions (gain 1 electron)

Group 0 — Noble Gases (He, Ne, Ar…)

  • Full outer electron shells → very stable → unreactive (inert)
  • Used in: He (balloons, diving), Ne (neon lights), Ar (light bulbs, welding)
  • All monatomic; boiling points increase down the group

Transition Metals (Period 4 — Ti to Cu)

Properties

  • High m.p. & density
  • Variable oxidation states
  • Coloured compounds
  • Good catalysts

Examples

  • Fe: catalyst (Haber)
  • Mn: MnO₂ catalyst
  • V₂O₅: Contact process
  • Ni: hydrogenation

Coloured Ions

  • Cu²⁺ → blue
  • Fe²⁺ → pale green
  • Fe³⁺ → orange/brown
  • Cr³⁺ → green
09
Metals
Reactivity series, extraction and corrosion
Reactivity Series

Most → Least Reactive

K > Na > Ca > Mg > Al > Zn > Fe > Ni > Sn > Pb > H > Cu > Ag > Au
Mnemonic: King Natasha Can Make A Zinc Fence Not Snapping Poorly, Here Comes Silver And Gold

Metal Extraction Methods

MetalMethodReason
K, Na, Ca, Mg, AlElectrolysis of molten oreToo reactive to reduce with carbon
Zn, Fe, Sn, PbReduction with carbon/COLess reactive than carbon
Cu (impure)Reduction, then electrolytic purificationLow reactivity; purify for conductivity
Ag, AuFound native (uncombined)Very unreactive
Blast Furnace — Extraction of Iron
Reduction of iron ore (haematite)
+
~1500°C
C(s) + O₂(g) → CO₂(g)   [coke burns]
CO₂(g) + C(s) → 2CO(g)   [CO formed]
Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g)   [iron reduced]
CaCO₃(s) → CaO(s) + CO₂(g)   [limestone decomposes]
CaO(s) + SiO₂(s) → CaSiO₃(l)   [slag removes impurity]
  • Raw materials: iron ore (Fe₂O₃), coke (C), limestone (CaCO₃), hot air blast
  • CO is the main reducing agent
  • Slag (CaSiO₃) floats on molten iron and is tapped off separately
  • Iron produced is 96% pure — contains C impurities → brittle

Diagram — Blast Furnace

250°C 600°C 1000°C 1500°C iron ore + coke + limestone (charged at top) CO₂ + C → 2CO Fe₂O₃ + 3CO → 2Fe + 3CO₂ (main reduction zone) C + O₂ → CO₂ coke burns (tuyeres) Hot air blast Hot air blast Slag tap (CaSiO₃ — floats) Molten iron (tapped at bottom) Waste gases (CO, CO₂, N₂ exit here) Fe₂O₃ + coke + CaCO₃ Blast Furnace — Extraction of Iron Raw materials: iron ore (haematite), coke, limestone, hot air
Extraction of Aluminium — Electrolysis of Bauxite (Al₂O₃)
Hall-Héroult Process
+
Molten Al₂O₃ dissolved in cryolite~950°C
At cathode: Al³⁺(l) + 3e⁻ → Al(l)
At anode: 2O²⁻(l) → O₂(g) + 4e⁻
4Al(l) + 3O₂(g) — from carbon anode: C + O₂ → CO₂ (anode burns away)
  • Cryolite (Na₃AlF₆) lowers melting point of Al₂O₃ from 2050°C to ~950°C → saves energy
  • Carbon anodes must be regularly replaced (oxidised by O₂)
  • Al is expensive due to large electricity consumption

Diagram — Hall-Héroult Electrolytic Cell

Steel shell Cathode (−) carbon lining Molten Aluminium — tapped via side tap Molten Al₂O₃ in cryolite (~950 °C) Al³⁺ moves → cathode O²⁻ moves → anode Al³⁺ → ← O²⁻ Anode (+) (Carbon) Anode (+) (Carbon) O₂(g) bubbles off C + O₂ → CO₂ (anode wears away) DC Supply + − − wire to cathode lining Al tap (liquid Al tapped) Cathode (−): Al³⁺(l) + 3e⁻ → Al(l) Anode (+): 2O²⁻(l) → O₂(g) + 4e⁻ (then C + O₂ → CO₂) Cryolite (Na₃AlF₆) lowers m.p. of Al₂O₃ from 2050 °C to ~950 °C — saves energy

Corrosion of Iron (Rusting)

Rusting of Iron
Requires O₂ AND H₂O
+
4Fe(s) + 3O₂(g) + 6H₂O(l) → 4Fe(OH)₃(s) → 2Fe₂O₃·xH₂O(s)
Hydrated iron(III) oxide = rust
  • Both O₂ and H₂O must be present — one alone is not enough
  • Salt water accelerates rusting (better electrolyte)
  • Prevention: painting, oiling/greasing, galvanising (Zn coating), electroplating, sacrificial anode (Zn/Mg bolted on)

Displacement Reactions

Metal Displacement from Salt Solution
More reactive metal displaces less reactive
+
Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
Zn(s) + 2AgNO₃(aq) → Zn(NO₃)₂(aq) + 2Ag(s)
Mg(s) + ZnSO₄(aq) → MgSO₄(aq) + Zn(s)
  • More reactive metal pushes out less reactive metal from its salt solution
  • Cu deposits on iron — reddish-brown solid

Alloys

AlloyCompositionProperties / Uses
Steel (mild)Fe + 0.1–0.3% CStronger than iron; bridges, cars
Stainless steelFe + Cr + NiCorrosion resistant; cutlery, sinks
BronzeCu + SnHarder than copper; statues, coins
BrassCu + ZnHarder, attractive; taps, musical instruments
DuraluminAl + Cu + MgLight & strong; aircraft
10
Chemistry of the Environment
Air, water, pollution and the greenhouse effect

Composition of Clean, Dry Air

Nitrogen

  • ~78%
  • N₂ — unreactive

Oxygen

  • ~21%
  • O₂ — supports combustion

Argon

  • ~0.9%
  • Noble gas — inert

Carbon dioxide

  • ~0.04%
  • CO₂ — greenhouse gas

Air Pollution

PollutantSourceEffectReduction
COIncomplete combustionToxic — binds haemoglobinCatalytic converters
SO₂Burning fossil fuels (S impurities)Acid rain (H₂SO₄ in rain)Remove S from fuel; flue gas desulfurisation
NOxHigh temp combustion (N₂+O₂)Acid rain; smogCatalytic converters
CO₂Burning fossil fuelsEnhanced greenhouse effect → global warmingUse renewables; carbon capture
Particulates (soot)Incomplete combustionRespiratory problems; global dimmingFilters; complete combustion
CFCAerosols, refrigerants (old)Ozone layer depletionBanned by Montreal Protocol
Acid Rain Formation
SO₂ and NOₓ reactions in atmosphere
+
2SO₂(g) + O₂(g) + 2H₂O(l) → 2H₂SO₄(aq)  [sulfurous → sulfuric acid]
4NO₂(g) + O₂(g) + 2H₂O(l) → 4HNO₃(aq)  [nitrogen dioxide → nitric acid]
N₂(g) + O₂(g) → 2NO(g)  [in car engines — high temp]
2NO(g) + O₂(g) → 2NO₂(g)
  • Acid rain: pH 4–5; corrodes limestone buildings, kills fish in lakes
  • CaCO₃(s) + H₂SO₄(aq) → CaSO₄(aq) + H₂O(l) + CO₂(g) — limestone dissolves

Catalytic Converter Reactions

Car Catalytic Converter
Pt/Pd/Rh catalyst — honeycomb structure
+
Pt/Rh/Pd catalystHigh temperature
2CO(g) + 2NO(g) → 2CO₂(g) + N₂(g)
2C₃H₈(g) + 7O₂(g) → 6CO₂(g) + 8H₂O(g)

Water Treatment

  • Sedimentation: large particles settle
  • Filtration: sand and gravel beds remove smaller particles
  • Chlorination: Cl₂ added to kill microorganisms (bacteria)
  • Fluoridation: NaF added in some areas to prevent tooth decay
  • Distilled water used in labs — no ions, very pure
  • Hard water: contains Ca²⁺/Mg²⁺ ions — removed by ion exchange or boiling (temporary hardness) or adding Na₂CO₃ (permanent hardness)
Removal of Temporary Hardness
Boiling or adding Na₂CO₃
+
Ca(HCO₃)₂(aq) → CaCO₃(s) + H₂O(l) + CO₂(g)  [boiling]
Ca²⁺(aq) + CO₃²⁻(aq) → CaCO₃(s)  [adding Na₂CO₃]
11
Experimental Techniques & Chemical Analysis
Separation, identification and safety

Separation Techniques

TechniqueUsed ForExample
FiltrationSolid from liquidSand from water
EvaporationSoluble solid from solutionSalt from salt water
CrystallisationPure crystals from solutionCuSO₄ crystals
DistillationLiquid from solution (by boiling point)Pure water from seawater
Fractional distillationMixtures of liquidsEthanol/water; crude oil fractions
ChromatographyMixtures of dissolved substancesFood dyes, inks, amino acids
CentrifugationInsoluble solid from liquidBlood cell separation

Chromatography — Rf Value

Rf FormulaRf = distance moved by substance ÷ distance moved by solvent front
  • Rf values between 0 and 1 (no units)
  • Same substance has same Rf under identical conditions
  • Used to identify substances by comparison with known standards

Tests for Gases

GasTestPositive Result
H₂Lighted splint near mouth of tubeSqueaky pop
O₂Glowing splint into tubeSplint relights
CO₂Bubble through limewater Ca(OH)₂Milky/cloudy white precipitate
Cl₂Damp litmus paperBleaches paper white
NH₃Damp red litmus paperTurns blue (alkaline)
H₂O (vapour)Anhydrous CuSO₄ (white)Turns blue

Tests for Ions (Flame Tests & Precipitates)

Flame Tests (Metal Cations)

IonFlame Colour
Li⁺Crimson red
Na⁺Yellow/orange
K⁺Lilac/purple
Ca²⁺Brick red
Cu²⁺Blue-green
Ba²⁺Green

NaOH Tests for Metal Ions

IonPrecipitate with NaOH
Cu²⁺Blue ppt. [Cu(OH)₂]
Fe²⁺Green ppt. [Fe(OH)₂]
Fe³⁺Red-brown ppt. [Fe(OH)₃]
Al³⁺White ppt. — dissolves in excess NaOH
Zn²⁺White ppt. — dissolves in excess NaOH
NH₄⁺Warm → NH₃ gas (turns litmus blue)

Tests for Anions

AnionTestResult
Cl⁻Add dilute HNO₃, then AgNO₃(aq)White ppt. AgCl — insoluble in dil. HNO₃
Br⁻Add dilute HNO₃, then AgNO₃(aq)Cream ppt. AgBr
I⁻Add dilute HNO₃, then AgNO₃(aq)Yellow ppt. AgI
SO₄²⁻Add dilute HCl, then BaCl₂(aq)White ppt. BaSO₄ — insoluble in HCl
CO₃²⁻Add dilute HCl — bubble through limewaterCO₂ gas — limewater turns milky
NO₃⁻Add NaOH, then Al foil, warmNH₃ gas evolved (turns damp red litmus blue)
⚠️ Exam tip: Always add dilute acid FIRST when testing for SO₄²⁻ (to remove CO₃²⁻ interference) and for halides (to prevent OH⁻ forming precipitates with Ag⁺).

Safety in the Laboratory

Hazard Symbols

  • ☠ Toxic
  • 🔥 Flammable
  • ⚡ Oxidising
  • ⚗ Corrosive
  • ☣ Harmful/Irritant

General Rules

  • Wear goggles & lab coat
  • No mouth pipetting
  • Label all containers
  • Dispose of chemicals properly

Measuring Accurately

  • Read burette at eye level (bottom of meniscus)
  • Use pipette for exact volumes
  • Use balance to 2 d.p.
📝
Self-Assessment Quiz
30 questions covering all 11 chapters — test your knowledge!

0 of 30 answered