O Level Chemistry 5070 — every topic, free.

Assessment objectives: AO1 Knowledge with understanding, AO2 Handling information and problem-solving, AO3 Experimental skills (20% — Papers 3 and 4 test this alone).

Tests for gases

Tests for aqueous cations (metal ions)

Add the reagent a little at a time, then in excess — the behaviour in excess is what distinguishes several of these.

Tests for anions

Flame tests for metal ions

Test for water: anhydrous copper(II) sulfate turns white → blue; anhydrous cobalt(II) chloride turns blue → pink. To show water is pure, check it boils at exactly 100 °C and freezes at exactly 0 °C at standard atmospheric pressure.

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1.1Solids, liquids and gases

Changes of state and their names: melting (s→l), freezing (l→s), boiling/evaporation (l→g), condensation (g→l), sublimation (s→g directly).

1.2Diffusion

Diffusion is faster at higher temperature (particles have more kinetic energy and move faster) and faster for lighter particles — at the same temperature, a gas with a lower relative molecular mass diffuses more quickly.

The classic demonstration: cotton wool soaked in concentrated ammonia solution at one end of a glass tube and concentrated hydrochloric acid at the other. A white ring of ammonium chloride forms nearer the hydrochloric acid end, because NH₃ (M_r = 17) is lighter than HCl (M_r = 36.5) and so diffuses faster.

2.1Elements, compounds and mixtures

• Element: a substance made of only one type of atom; cannot be broken down chemically.
• Compound: two or more different elements chemically combined in a fixed ratio; has properties different from its elements and can only be separated chemically.
• Mixture: two or more substances not chemically combined; keeps the properties of its components and can be separated by physical means.

2.2Atomic structure and the Periodic Table  2.3Isotopes

Electronic configuration: shells fill 2, 8, 8… (e.g. chlorine, 17 electrons → 2,8,7). In the Periodic Table, the group number equals the number of outer-shell electrons and the period number equals the number of occupied shells. Elements in the same group have similar chemical properties because they have the same number of outer electrons.

2.4Ions and ionic bonds

Metals lose electrons to form positive ions (cations); non-metals gain electrons to form negative ions (anions). Group I → 1+, Group II → 2+, Group VI → 2−, Group VII → 1−.

2.5Simple molecules and covalent bonds  2.6Giant covalent structures

Simple molecular substances (H₂, Cl₂, H₂O, CH₄, CO₂) have low melting and boiling points: the covalent bonds within molecules are strong, but the intermolecular forces between molecules are weak and only these need to be overcome when melting or boiling. They do not conduct electricity — there are no free electrons or ions.

All giant covalent structures have very high melting points because many strong covalent bonds must be broken.

2.7Metallic bonding

• Good conductors of electricity and heat — delocalised electrons are free to move and carry charge/energy.
• Malleable and ductile — the layers of positive ions can slide over each other without breaking the metallic bonding, because the delocalised electrons keep holding the lattice together.
• High melting points — strong attraction between the ions and the delocalised electrons.

3.1Formulae

Write formulae by balancing the charges so the compound is neutral. Common ions to know:

Balancing equations: change only the big numbers in front (coefficients), never the small subscripts inside a formula. Include state symbols where asked: (s) solid, (l) liquid, (g) gas, (aq) aqueous.

3.2Relative masses  3.3The mole and the Avogadro constant

4.1Electrolysis

In dilute sulfuric acid, hydrogen is produced at twice the volume of oxygen — the classic Hofmann voltameter observation.

4.2Hydrogen–oxygen fuel cells

A fuel cell uses hydrogen and oxygen to produce electricity, with water as the only product:

5.1Exothermic and endothermic reactions

6.1Physical and chemical changes  6.2Rate of reaction

Physical change: no new substance formed, usually easily reversible (melting, dissolving). Chemical change: a new substance is formed, usually hard to reverse, often with an energy change, colour change, gas or precipitate.

6.3Reversible reactions and equilibrium

A reversible reaction (⇌) can go both ways. In a closed system it reaches equilibrium: the forward and backward reactions occur at the same rate, so the concentrations of reactants and products stay constant (they are not equal — just unchanging).

The classic reversible reaction is the hydration of anhydrous copper(II) sulfate — white ⇌ blue, used as the test for water.

The Haber process (ammonia): N₂ + 3H₂ ⇌ 2NH₃ — about 450 °C, 200 atm, iron catalyst. The temperature is a compromise: the forward reaction is exothermic so a lower temperature would give more ammonia, but the rate would be uneconomically slow.

6.4Redox

Oxidation numbers can also be used: oxidation is an increase in oxidation number, reduction a decrease. Roman numerals in names give the oxidation number, e.g. iron(III) is Fe³⁺.

7.1Characteristic properties of acids and bases

Strong vs weak: a strong acid is fully dissociated into ions in solution (HCl, HNO₃, H₂SO₄); a weak acid is only partially dissociated (ethanoic acid, carbonic acid). This is different from concentration — a concentrated weak acid is still weak.

Indicators: litmus is red in acid, blue in alkali. Methyl orange is red in acid, yellow in alkali. Thymolphthalein is colourless in acid, blue in alkali.

7.2Oxides

7.3Preparation of salts

8.1Arrangement of elements

Elements are arranged in order of increasing proton number. Groups (columns) contain elements with the same number of outer-shell electrons and hence similar chemical properties; periods (rows) show the number of occupied shells. Metals are on the left, non-metals on the right.

Across a period, elements change from metallic to non-metallic. Down a group, atoms gain more shells, so the atomic radius increases and the outer electrons are further from the nucleus and more shielded.

8.2Group I — the alkali metals  8.3Group VII — the halogens

Group I with water: lithium fizzes steadily; sodium melts into a ball and moves rapidly on the surface; potassium reacts violently and the hydrogen produced ignites with a lilac flame.

8.4Transition elements  8.5Noble gases

Transition elements (the central block) differ from Group I metals: they have high densities, high melting points, form coloured compounds, show variable oxidation states, and are often used as catalysts (iron in the Haber process, nickel in hydrogenation). They are also much harder and less reactive than Group I metals.

Noble gases (Group VIII/0) are unreactive because they already have a full outer electron shell, so they have no tendency to gain, lose or share electrons. Uses follow from that inertness: helium in balloons (low density, non-flammable), argon in lamps and as an inert welding atmosphere.

9.1Properties  9.2Uses  9.3Alloys

General properties of metals: good conductors of heat and electricity, malleable, ductile, shiny when freshly cut, generally high melting points and densities — all explained by metallic bonding (2.7).

Uses linked to properties: copper for electrical wiring (excellent conductor, ductile); aluminium for aircraft and overhead cables (low density, corrosion-resistant due to its oxide layer); mild steel for car bodies and construction (strong, cheap); stainless steel for cutlery and chemical plant (resists corrosion).

9.4Reactivity series  9.5Corrosion

Displacement: a more reactive metal displaces a less reactive one from its compound (e.g. zinc displaces copper from copper(II) sulfate — the blue solution fades and a brown/pink solid forms). These reactions are exothermic and are used to place unknown metals in the series.

Preventing rust: barrier methods (painting, greasing, plastic coating, galvanising with zinc) keep out air and water. Sacrificial protection works differently: a more reactive metal such as zinc or magnesium is attached and corrodes instead of the iron, because it loses electrons more readily. Galvanising does both.

9.6Extraction of metals

Iron in the blast furnace — raw materials: iron ore (haematite, Fe₂O₃), coke, limestone and hot air.

1. Coke burns: C + O₂ → CO₂ (exothermic — provides the heat)
2. Carbon dioxide is reduced: CO₂ + C → 2CO
3. Carbon monoxide reduces the ore: Fe₂O₃ + 3CO → 2Fe + 3CO₂
4. Limestone removes the sand impurity: CaCO₃ → CaO + CO₂, then CaO + SiO₂ → CaSiO₃ (slag, which floats on the molten iron and is tapped off separately)

Aluminium by electrolysis: aluminium oxide is dissolved in molten cryolite to lower the melting point and so reduce energy costs. At the carbon cathode Al³⁺ + 3e⁻ → Al; at the carbon anode 2O²⁻ → O₂ + 4e⁻. The anodes burn away in the oxygen produced and must be replaced regularly.

10.1Water  10.2Fertilisers

Water treatment for drinking: sedimentation/filtration to remove solids, then chlorination to kill bacteria. Note this makes water safe, not chemically pure.

Tests: for water — anhydrous copper(II) sulfate white → blue, or cobalt(II) chloride blue → pink. For pure water — boiling point exactly 100 °C and freezing point exactly 0 °C at standard atmospheric pressure.

Fertilisers supply the elements plants need: nitrogen, phosphorus and potassium (NPK). Nitrogen promotes leaf/stem growth (proteins), phosphorus root growth, potassium flowering and disease resistance. Ammonium salts and nitrates are common nitrogenous fertilisers.

10.3Air quality and climate

Clean dry air is approximately 78% nitrogen, 21% oxygen, with argon, carbon dioxide and other gases making up the rest.

Catalytic converters in cars reduce these pollutants, e.g. 2CO + 2NO → 2CO₂ + N₂, converting toxic gases into harmless ones.

11.1Formulae and terminology  11.2Naming organic compounds

Naming: the stem gives the number of carbon atoms — meth- (1), eth- (2), prop- (3), but- (4) — and the ending gives the family.

11.3Fuels

Fossil fuels are coal, natural gas and petroleum. Petroleum is a mixture of hydrocarbons separated by fractional distillation, which works because different fractions have different boiling points — the column is hottest at the bottom, so smaller molecules with lower boiling points rise higher before condensing.

Combustion: complete combustion (plenty of oxygen) gives carbon dioxide and water; incomplete combustion (limited oxygen) gives carbon monoxide (toxic) and/or carbon (soot). Methane is the main constituent of natural gas: CH₄ + 2O₂ → CO₂ + 2H₂O

11.4Alkanes  11.5Alkenes

Alkanes are saturated hydrocarbons with only single covalent bonds. They are generally unreactive, except for combustion and substitution with chlorine.

Alkenes contain a C=C double bond and are unsaturated, which makes them much more reactive than alkanes.

Addition reactions of alkenes — the double bond opens and only one product is formed:

11.6Alcohols  11.7Carboxylic acids

Uses of ethanol: as a solvent and as a fuel. Alcohols burn in oxygen to give carbon dioxide and water: C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O

Carboxylic acids behave like other acids — they are weak acids, so they react in the usual three ways:

Making ethanoic acid by oxidising ethanol: (a) with acidified aqueous potassium manganate(VII) (purple → colourless), or (b) by bacterial oxidation, which is how vinegar is produced and why wine turns sour in air.

11.8Polymers

Disposal problems: most addition polymers are non-biodegradable, so they persist in landfill for centuries; burning them can release toxic gases (such as HCl from PVC) and carbon dioxide; and litter harms wildlife. Recycling reduces this but requires sorting by polymer type.

12.1Experimental design  12.2Acid–base titrations

Apparatus for measuring volume: a burette (accurate, variable volumes, readings to 0.05 cm³), a pipette (one fixed accurate volume), a measuring cylinder (approximate volumes only). Choose the burette or pipette whenever accuracy matters.

12.3Chromatography  12.4Separation and purification

Paper chromatography separates mixtures of soluble coloured (or, with a locating agent, colourless) substances. The baseline is drawn in pencil — ink would dissolve and run — and the solvent level must start below the baseline, or the spots wash off into the solvent.

R_f values are always between 0 and 1, and are constant for a given substance and solvent, so they identify substances by comparison. A pure substance gives one spot; a mixture gives more than one.

Assessing purity: a pure substance has a sharp, fixed melting and boiling point. Impurities lower the melting point and raise the boiling point, and make the melting occur over a range rather than at one temperature.

12.5Identification of ions and gases

All the required tests are collected in the qualitative analysis tables at the top of this guide — gases, cations, anions and flame tests. Remember: those tables are given to you in Papers 3 and 4 but not in Paper 2.

The experimental contexts you must know

• Quantitative measurement of volumes, masses, temperatures, times and lengths
• Rates of reaction; temperature changes during reactions; heating and cooling curves
• Salt preparation; solubility; melting and boiling points
• Separation and purification: filtration, crystallisation, simple and fractional distillation, chromatography
• Electrolysis; displacement reactions of metals and halogens; conditions for rusting
• Identification of metal ions, non-metal ions and gases; chemical tests for water
• Test-tube reactions of dilute acids, including ethanoic acid; tests for oxidising and reducing agents
• Titrations, and unfamiliar procedures with simple apparatus

Recording observations — what examiners want to see

Graphs and evaluation

• Independent variable on the x-axis; label both axes with quantity and unit.
• Scales must let the points fill more than half the grid in both directions, using sensible ratios.
• Plot as small crosses (×) or encircled dots, accurate to half a small square.
• Draw a single thin best-fit line or smooth curve; ignore clearly anomalous points when drawing it (but circle and mention them).
• For a gradient, use a triangle whose hypotenuse spans at least half the line, and mark it on the graph.

Safety should match the hazard: acids and alkalis → wear eye protection; heating → use tongs and a heatproof mat, point the test tube away from people; toxic or irritant gases (chlorine, sulfur dioxide) → work in a fume cupboard; flammable liquids (ethanol) → keep away from naked flames and heat with a water bath.

Official (free)

• Cambridge International — 5070 subject page: syllabus, specimen papers, past papers, mark schemes and examiner reports.
• Examiner reports list the exact tests and explanations candidates got wrong each series.

Free archives

• GCE Guide · PastPapers.co — full CAIE archives including practical papers.
• Physics & Maths Tutor — topic-sorted questions.

The method that fits how chemistry is marked

1. Learn the tests table cold. It is provided in the practical paper but not in Paper 2, and it appears in almost every theory paper — the cheapest marks in the subject.
2. Drill moles. One mole question a day for a month covers nearly every calculation the exam can ask.
3. Practise explanations out loud. Most explanation marks follow a fixed pattern: particles → forces → what happens to them. Say it until it is automatic.
4. Write equations from words. Take word equations from past papers and convert them to balanced symbol equations with state symbols.
5. Do the practical papers. 20% of your grade, and the question patterns repeat heavily year to year.