A Level Chemistry 9701 — all 37 topics, free.

Three routes. AS Level only = Papers 1, 2, 3. A Level staged over two years = Papers 1, 2, 3 in year 1, then Papers 4 and 5 in year 2. A Level in one series = all five papers.

1.1Particles in the atom and atomic radius

Behaviour in an electric field: protons deflect one way, electrons deflect far more in the opposite direction (same magnitude of charge, tiny mass), neutrons are undeflected.

Trends in atomic radius. Across a period the radius decreases — the nuclear charge rises while shielding stays roughly constant, so electrons are pulled in harder. Down a group it increases — an extra filled shell adds distance and shielding. A cation is smaller than its atom (a whole shell is often lost); an anion is larger (more electron–electron repulsion for the same nuclear charge).

1.2Isotopes

Isotopes have the same number of protons but different numbers of neutrons. They have identical chemical properties (same electronic configuration) but slightly different physical properties such as density and rate of diffusion.

1.3Electrons, energy levels and atomic orbitals

An orbital is a region that can hold up to two electrons of opposite spin. s orbitals are spherical; p orbitals are dumb-bell shaped along the x, y and z axes. Electrons fill degenerate orbitals singly first, with parallel spins, before pairing (Hund's rule).

1.4Ionisation energy

Ionisation energy depends on nuclear charge, atomic radius and shielding. Successive ionisation energies always increase (removing an electron from an increasingly positive ion), and a large jump marks the start of a new, closer, less shielded shell — which is how you deduce the group.

Across Period 3 the general rise is broken twice: Al < Mg because Al's outer electron is in a 3p orbital, higher in energy and better shielded than Mg's 3s; and S < P because S's 3p⁴ has a paired electron in one orbital, and the pair repulsion makes it easier to remove.

2.1Relative masses of atoms and molecules

Relative isotopic, atomic, molecular and formula masses are all defined against 1/12 of the mass of a ¹²C atom. They have no units.

2.2The mole and the Avogadro constant

2.3Formulas

Empirical formula is the simplest whole-number ratio of atoms; molecular formula is the actual number. Get the empirical formula by dividing percentage (or mass) by A_r, then dividing through by the smallest result.

2.4Reacting masses and volumes

The limiting reagent is the one that gives the smaller number of moles of product — never assume it is the one with the smaller mass. Percentage yield = (actual ÷ theoretical) × 100.

3.1Electronegativity and bonding

Electronegativity is the ability of an atom to attract the electron pair in a covalent bond. It increases across a period and decreases down a group; F is the most electronegative element. A large difference gives ionic bonding, a small difference polar covalent, and zero difference pure covalent.

3.2Ionic bonding

Electrostatic attraction between oppositely charged ions in a giant lattice. High melting point, brittle, conducts only when molten or aqueous (ions free to move).

3.3Metallic bonding

Attraction between a lattice of positive ions and a sea of delocalised electrons. Explains electrical and thermal conductivity, malleability (layers slide without breaking the bonding) and high melting points. Strength increases with ionic charge and decreasing ionic radius — which is why Mg melts far above Na.

3.4Covalent bonding and coordinate (dative covalent) bonding

A covalent bond is a shared pair of electrons. In a dative covalent bond both electrons come from the same atom — as in NH₄⁺, H₃O⁺, and every ligand–metal bond in a complex ion. Show it with an arrow from the donor.

3.5Shapes of molecules

3.6Intermolecular forces, electronegativity and bond properties

In increasing strength: instantaneous dipole–induced dipole (id–id, London) forces — present in everything, stronger for larger and more elongated molecules with more electrons; permanent dipole–permanent dipole (pd–pd); hydrogen bonding — only when H is bonded directly to N, O or F and a lone pair is available on the acceptor.

Hydrogen bonding explains water's anomalously high boiling point, ice being less dense than water (an open tetrahedral lattice), the miscibility of alcohols with water, and the higher boiling points of alcohols than of comparable alkanes.

3.7Dot-and-cross diagrams

Draw them for ionic compounds (showing charges and square brackets), simple covalent molecules, molecules with multiple bonds, dative bonds, and species with an expanded octet such as SF₆ and PCl₅. Show all outer-shell electrons — including lone pairs on the central atom.

4.1The gaseous state: ideal and real gases

Ideal gas assumptions: negligible molecular volume, no intermolecular forces, elastic collisions, random motion. Real gases deviate most at high pressure and low temperature, where the molecules are close enough for their own volume and their mutual attractions to matter.

4.2Bonding and structure

Diamond: every carbon bonded to four others tetrahedrally, no free electrons, extremely hard. Graphite: layers of hexagons, three bonds per carbon, the fourth electron delocalised — conducting along the layers, soft because weak forces let layers slide.

5.1Enthalpy change, ΔH

Standard conditions: 298 K and 100 kPa, all substances in their standard states. Know the definitions of ΔH°_f (formation, from elements), ΔH°_c (combustion, complete, in excess oxygen), ΔH°_neut, ΔH°_at and bond energy — each "per mole of" something specific.

5.2Hess's law

The enthalpy change is independent of the route taken. Note the two formulas point in opposite directions — formation arrows point up from the elements, combustion arrows point down to the oxides.

6.1Redox processes: electron transfer and oxidation numbers

Oxidation-number rules: elements = 0; Group 1 = +1, Group 2 = +2; H = +1 (except −1 in metal hydrides); O = −2 (except −1 in peroxides, +2 in OF₂); F = −1 always; the sum equals the overall charge.

The oxidising agent is itself reduced; the reducing agent is itself oxidised. Use oxidation numbers to construct and balance redox half-equations, and to name compounds using Roman numerals — iron(III) chloride, manganate(VII).

7.1Chemical equilibria: reversible reactions, dynamic equilibrium

At dynamic equilibrium the forward and reverse rates are equal and concentrations are constant — in a closed system. Le Chatelier's principle: a system at equilibrium shifts to oppose any change imposed on it.

7.2Brønsted–Lowry theory of acids and bases

An acid is a proton donor, a base a proton acceptor. Every acid has a conjugate base (itself minus H⁺) and every base a conjugate acid. Strong acids are fully dissociated; weak acids are partially dissociated — an equilibrium.

8.1Rate of reaction

Reaction requires collisions with sufficient energy (≥ E_a) and the correct orientation. Rate is increased by higher concentration or pressure (more frequent collisions), larger surface area, higher temperature, and a catalyst.

8.2Effect of temperature and activation energy

The Boltzmann distribution shows the spread of molecular energies. Raising the temperature flattens and broadens the curve and shifts the peak to the right, so a much larger proportion of molecules exceed E_a — which is why a 10 K rise can roughly double the rate even though the mean energy rises only slightly.

8.3Homogeneous and heterogeneous catalysts

A catalyst provides an alternative route of lower activation energy; it is not used up. Homogeneous catalysts are in the same phase as the reactants (they form an intermediate); heterogeneous catalysts are in a different phase and work by adsorption of reactants onto active sites, weakening bonds, then desorption of products.

9.1Periodicity of physical properties in Period 3

Across Na → Ar: atomic radius decreases (rising nuclear charge, similar shielding); ionic radius falls across the cations then jumps up at P³⁻/S²⁻/Cl⁻ (anions are larger); first ionisation energy rises with the Al and S dips; electrical conductivity is high for Na, Mg, Al (metallic, and it rises with the number of delocalised electrons per atom) then negligible from Si onwards.

Melting point is the classic graph: rises Na → Mg → Al (stronger metallic bonding), peaks sharply at Si (giant covalent), then collapses to simple molecular values — P₄, S₈ (higher than P₄ because S₈ is bigger, so stronger id–id forces), Cl₂, and lowest of all Ar (single atoms).

9.2Periodicity of chemical properties in Period 3

Reaction with oxygen gives Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀, SO₂/SO₃. Reaction with chlorine gives NaCl, MgCl₂, Al₂Cl₆, SiCl₄, PCl₅. With water: Na reacts vigorously giving a strongly alkaline solution; Mg reacts very slowly with cold water but readily with steam.

Acid–base character of the oxides shifts across the period: Na₂O and MgO are basic, Al₂O₃ is amphoteric, and SiO₂, P₄O₁₀, SO₂ and SO₃ are acidic. The pattern follows the bonding — ionic oxides are basic, giant/simple covalent oxides are acidic.

Hydrolysis of the chlorides: NaCl and MgCl₂ simply dissolve (roughly neutral, MgCl₂ slightly acidic); Al₂Cl₆, SiCl₄ and PCl₅ are hydrolysed, producing fumes of HCl and strongly acidic solutions.

9.3Chemical periodicity of other elements

You should be able to predict the properties of an unfamiliar element or compound from its position in the Periodic Table — for example, deducing the formula, bonding, acid–base character of the oxide, and reaction with water for an element of Period 2 or Period 4 by analogy with Period 3.

10.1Group 2 metals, magnesium to barium, and their compounds

Down the group: reactivity with water increases (ionisation energies fall as the atomic radius grows), and the metals become stronger reducing agents. Hydroxide solubility increases; sulfate solubility decreases (BaSO₄ is famously insoluble — the basis of the sulfate test and of barium meals).

Thermal stability of carbonates and nitrates increases down the group. The explanation: the cation's charge density (charge/radius ratio) falls, so it polarises the large anion less, weakening the anion less, so more energy is needed to decompose it.

11.1Physical properties of the Group 17 elements

Down the group the boiling point increases — more electrons, stronger id–id forces. Colours: Cl₂ pale green gas, Br₂ red-brown liquid, I₂ grey-black solid subliming to a purple vapour. In aqueous solution: chlorine pale yellow-green, bromine orange, iodine brown; in an organic solvent, iodine is a distinctive violet/purple.

11.2Chemical properties of the halogens and hydrogen halides

Oxidising power decreases down the group — so chlorine displaces bromide and iodide, and bromine displaces iodide only. Use displacement reactions and their colour changes as tests.

Thermal stability of the hydrogen halides decreases down the group (HCl > HBr > HI) because the H–X bond gets longer and weaker. HI decomposes noticeably on gentle heating, giving purple iodine vapour.

11.3Some reactions of the halide ions

The concentrated sulfuric acid results show the increasing reducing power of the halide ions down the group: iodide is a strong enough reducing agent to take sulfur from +6 all the way to −2.

11.4The reactions of chlorine

Both are disproportionation reactions — chlorine is simultaneously oxidised and reduced (0 → −1 and 0 → +1, or 0 → +5). Chlorine in water gives HCl and HClO; the HClO is the bactericide, which is why chlorination is used to treat drinking water despite chlorine's toxicity.

12.1Nitrogen and sulfur

Nitrogen is unreactive because of the very strong N≡N triple bond and the molecule's non-polarity. Ammonia is a base (its lone pair accepts a proton) and a ligand; NH₃ + H⁺ → NH₄⁺, and heating an ammonium salt with NaOH(aq) releases ammonia — the standard test, turning damp red litmus blue.

Environmental chemistry: nitrogen oxides form in internal combustion engines and catalyse the oxidation of SO₂ to SO₃ in the atmosphere; both contribute to acid rain. Catalytic converters remove NO_x, CO and unburnt hydrocarbons. Sulfur dioxide from fossil fuels dissolves to give acidic solutions; flue-gas desulfurisation with CaO or CaCO₃ removes it.

Ammonium fertilisers raise crop yields but leaching causes eutrophication, and adding them to alkaline soils releases ammonia gas, wasting nitrogen.

13.1Formulas, functional groups and naming

Know empirical, molecular, general, structural, displayed and skeletal formulas, and be able to convert between them. Naming follows IUPAC: longest chain → suffix for the principal functional group → substituents alphabetically with lowest possible locants.

13.2Characteristic organic reactions

Classify by type (addition, substitution, elimination, hydrolysis, oxidation, reduction, polymerisation) and by mechanism (free-radical, nucleophilic, electrophilic). A nucleophile donates an electron pair; an electrophile accepts one. Homolytic fission gives two radicals; heterolytic fission gives a cation and an anion.

13.3Shapes of organic molecules; σ and π bonds

A σ bond is end-on overlap, with the electron density along the internuclear axis; a π bond is sideways overlap of p orbitals, above and below that axis. A double bond is one σ + one π. The π bond restricts rotation, which is exactly why cis–trans isomerism exists. Bond angles: 109.5° around a saturated carbon, 120° around a C=C.

13.4Isomerism

Structural isomerism: chain, positional and functional group. Stereoisomerism: cis–trans (needs a C=C or ring and two different groups on each of the two carbons) and, at A2, optical isomerism.

14.1Alkanes

Saturated, largely unreactive (strong non-polar C–C and C–H bonds). Combustion — complete gives CO₂ and H₂O, incomplete gives CO and soot. Free-radical substitution with halogens in UV light, in three stages:

Cracking breaks long-chain alkanes into shorter alkanes and alkenes, meeting demand for petrol and for alkene feedstock.

14.2Alkenes

Mechanism of electrophilic addition (Br₂ + ethene): the π electrons induce a dipole in Br₂; the curly arrow goes from the C=C to the δ+ bromine, a second arrow from the Br–Br bond to the δ− bromine, giving a carbocation and Br⁻; a third arrow from the Br⁻ lone pair to the positive carbon completes the addition.

Carbocation stability: tertiary > secondary > primary, because alkyl groups are electron-releasing and spread the positive charge. That is the reason behind Markovnikov's rule — quote it as the explanation, not the rule itself.

15.1Halogenoalkanes

S_N1 vs S_N2. Tertiary halogenoalkanes go by S_N1: the C–X bond breaks first to give a relatively stable tertiary carbocation, which the nucleophile then attacks. Primary halogenoalkanes go by S_N2: the nucleophile attacks the δ+ carbon from the opposite side to the halogen in one step, through a transition state, with inversion of configuration.

Rate of hydrolysis follows the C–X bond enthalpy: C–I is weakest so iodoalkanes hydrolyse fastest, then bromo, then chloro. Demonstrate it with AgNO₃ in ethanol and compare how quickly each precipitate appears.

16.1Alcohols

Classifying an unknown alcohol is a standard question: react with acidified dichromate. Orange → green with an aldehyde/acid product = primary; orange → green with a ketone = secondary; no colour change = tertiary. The tri-iodomethane test then separates methyl-carbinols from the rest.

17.1Aldehydes and ketones

Reduction with NaBH₄ (or LiAlH₄ in dry ether) gives a primary alcohol from an aldehyde and a secondary alcohol from a ketone — nucleophilic addition of H⁻.

Nucleophilic addition of HCN (with NaCN, trace acid or alkali) gives a hydroxynitrile, adding one carbon. Mechanism: CN⁻ lone pair attacks the δ+ carbonyl carbon, the π electrons move onto the oxygen giving an alkoxide, which is then protonated.

18.1Carboxylic acids

Weak acids: they react with reactive metals (H₂), with carbonates (CO₂ effervescence — a test that distinguishes them from phenols and alcohols), and with alkalis to give salts. Reduced by LiAlH₄ in dry ether to a primary alcohol. Esterified by an alcohol with a concentrated H₂SO₄ catalyst.

18.2Esters

Acid hydrolysis (dilute acid, reflux) is reversible and gives the carboxylic acid and the alcohol. Alkaline hydrolysis (NaOH(aq), reflux) goes to completion and gives the carboxylate salt plus the alcohol — this is saponification, the basis of soap-making.

19.1Primary amines

Amines are bases — the nitrogen lone pair accepts a proton, giving an alkylammonium salt. Prepared by heating a halogenoalkane with excess ethanolic ammonia in a sealed tube (excess, to reduce further substitution to secondary and tertiary amines), or by reducing a nitrile.

19.2Nitriles and hydroxynitriles

Nitriles are the key chain-lengthening intermediate: halogenoalkane + KCN(ethanolic, reflux) → nitrile, then either hydrolysis (dilute acid, reflux) to a carboxylic acid or reduction (LiAlH₄, or H₂/Ni) to a primary amine — one extra carbon either way.

20.1Addition polymerisation

Alkenes polymerise by addition: the C=C opens and monomers join with no other product. Deduce the repeat unit from the monomer (and vice versa) — draw it in square brackets with trailing bonds and a subscript n.

Environmental consequences: addition polymers have inert, non-polar C–C backbones, so they are non-biodegradable. Disposal by landfill wastes space and by incineration can release toxic gases (HCl from PVC); recycling and the development of degradable polymers are the responses.

21.1Organic synthesis (AS)

You will be asked to design two- or three-step routes between AS functional groups, giving reagents and conditions for each step. Work backwards from the target: identify the functional group change and the change in carbon number, then pick the step that achieves it.

22.1Infrared spectroscopy

Bonds absorb IR at characteristic wavenumbers. You are given a data table in the exam — the skill is reading it, not memorising it. The signals that decide most questions:

22.2Mass spectrometry

The molecular ion peak (M) at the highest m/z gives the M_r. The M+1 peak comes from the 1.1% natural abundance of ¹³C, so the number of carbon atoms ≈ (100 × height of M+1) / (1.1 × height of M).

Look for characteristic M+2 patterns: a 1:1 ratio means one chlorine (³⁵Cl:³⁷Cl ≈ 3:1 gives M:M+2 of 3:1 — careful, chlorine is 3:1 and bromine is 1:1); a 1:1 M:M+2 pair means one bromine (⁷⁹Br:⁸¹Br ≈ 1:1).

23.1Lattice energy and Born–Haber cycles

Lattice energy becomes more exothermic with higher ionic charge and smaller ionic radius — MgO (2+/2−, small ions) is far more exothermic than NaCl.

A Born–Haber cycle links ΔH_f to atomisation enthalpies, ionisation energies, electron affinities and lattice energy. Set it out as an energy-level diagram with arrows, and remember the second electron affinity is endothermic (adding an electron to an already negative ion).

23.2Enthalpies of solution and hydration

Hydration enthalpy is exothermic and, like lattice energy, becomes more exothermic with higher charge density. A salt dissolves readily when the hydration enthalpies roughly repay the lattice energy.

23.3Entropy change, ΔS

Entropy is a measure of the number of ways the particles and their energy can be arranged — the disorder of the system. It increases solid → liquid → gas, on dissolving a solid, and when the number of gaseous moles increases. Dissolving an ionic solid can decrease entropy if the ions strongly order the surrounding water.

23.4Gibbs free energy change, ΔG

24.1Electrolysis

24.2Standard electrode potentials and the Nernst equation

Standard conditions: 298 K, 1 mol dm⁻³ solutions, 100 kPa gases, measured against the standard hydrogen electrode (defined as 0.00 V). A platinum electrode is used where no metal is involved.

The more positive E° is the better oxidising agent (more readily reduced). A cell reaction is feasible when E°_cell is positive.

25.1Acids and bases

For a weak acid, [H⁺] = √(K_a × [HA]) using the usual approximations. For a buffer, use the Henderson–Hasselbalch relation pH = pK_a + log([salt]/[acid]).

A buffer resists pH change on addition of small amounts of acid or alkali. An acidic buffer is a weak acid plus its salt: added H⁺ reacts with the large reservoir of A⁻, added OH⁻ reacts with the large reservoir of HA.

Titration curves and indicators. Strong acid–strong base has a long vertical section from about pH 3 to 11 (most indicators work). Weak acid–strong base has its equivalence point above pH 7 — use phenolphthalein. Strong acid–weak base has it below pH 7 — use methyl orange. Weak acid–weak base has no sharp vertical section, so no indicator is suitable. At the half-equivalence point of a weak acid titration, pH = pK_a.

25.2Partition coefficients

The ratio of concentrations of a solute distributed between two immiscible solvents at equilibrium, at a fixed temperature. Valid only if the solute is in the same molecular state in both solvents — association or dissociation invalidates the simple expression. This is the principle behind solvent extraction and chromatography.

26.1Rate equations, orders of reaction and rate constants

Orders are found experimentally, never from the stoichiometric equation. From a concentration–time graph: zero order gives a straight line of constant gradient; first order gives a curve with a constant half-life; second order gives a steeper curve with a half-life that doubles each time. From a rate–concentration graph: zero order is horizontal, first order is a straight line through the origin, second order is an upward curve.

The rate-determining step is the slowest step; the rate equation tells you which species (and how many of each) are involved in it or before it. This is how kinetics distinguishes S_N1 (rate = k[halogenoalkane], first order overall) from S_N2 (rate = k[halogenoalkane][nucleophile], second order).

26.2Homogeneous and heterogeneous catalysts (A2)

Explain catalysis in terms of an alternative route of lower E_a, and — for transition metals — their ability to use variable oxidation states to form intermediates. Two examples to know: Fe in the Haber process (heterogeneous, adsorption on active sites) and Fe²⁺/Fe³⁺ catalysing the S₂O₈²⁻/I⁻ reaction (homogeneous, via changing oxidation state).

27.1Group 2 metals and their compounds (A2 treatment)

At A2 the same trends are explained quantitatively with energetics. Thermal stability of carbonates increases down the group because the lattice energies of MCO₃ and MO both become less exothermic as the cation grows, but the small oxide ion means ΔH_latt(MO) falls off faster, so decomposition becomes less favourable. Solubility trends are explained by comparing lattice energy with the sum of hydration enthalpies.

28.1General physical and chemical properties (Ti to Cu)

A transition element has an incomplete d sub-shell in at least one of its ions. That definition excludes Sc (Sc³⁺ is d⁰) and Zn (Zn²⁺ is d¹⁰) — a favourite one-mark question.

Characteristic properties, all traceable to the partly filled d sub-shell: variable oxidation states, coloured compounds, catalytic activity, and the formation of complex ions. They also have high melting points and densities compared with Group 1 and 2 metals.

28.2Characteristic chemical properties

A ligand is a species with a lone pair that forms a dative covalent bond to the central metal ion. Monodentate: H₂O, NH₃, Cl⁻, CN⁻, OH⁻. Bidentate: ethanedioate, 1,2-diaminoethane. Hexadentate: EDTA⁴⁻. The coordination number is the number of dative bonds — usually 6 (octahedral) with small ligands, 4 (tetrahedral) with larger ones like Cl⁻.

Ligand exchange reactions give sharp colour changes: [Cu(H₂O)₆]²⁺ (pale blue) + excess NH₃ → [Cu(NH₃)₄(H₂O)₂]²⁺ (deep blue); + concentrated HCl → [CuCl₄]²⁻ (yellow-green), which is also a change of shape and coordination number.

28.3Colour of complexes

Ligands split the five degenerate d orbitals into two groups of different energy. An electron absorbs a photon of visible light to jump the gap; the colour you see is the complement of the colour absorbed. The size of the splitting — and hence the colour — depends on the ligand, the oxidation state, and the coordination number, which is why exchanging ligands changes the colour.

28.4Stereoisomerism in transition element complexes

Cis–trans isomerism occurs in square planar complexes such as Pt(NH₃)₂Cl₂ and in octahedral complexes of the type [M(A)₄(B)₂]. Optical isomerism occurs in octahedral complexes with three bidentate ligands, such as [Ni(NH₂CH₂CH₂NH₂)₃]²⁺ — the two forms are non-superimposable mirror images.

28.5Stability constants, K_stab

K_stab is the equilibrium constant for the formation of a complex from the aqueous metal ion. A larger K_stab means a more stable complex, so a ligand with a high K_stab will displace one with a lower value.

29.1Formulas, functional groups and naming (A2)

Extends AS naming to arenes, phenols, acyl chlorides, amides, amino acids and aromatic amines. Learn benzene ring numbering and the ortho/meta/para positions in the substituted products of electrophilic substitution.

29.2Characteristic organic reactions (A2)

Adds electrophilic substitution (arenes), nucleophilic addition–elimination (acyl chlorides), condensation (polymers, amides, esters) and diazotisation and coupling to the AS list of reaction types.

29.3Shapes of aromatic organic molecules; σ and π bonds

Benzene is planar, with all C–C bonds equal in length (between a single and a double bond) and all angles 120°. Each carbon contributes one electron to a delocalised π system above and below the ring. Evidence for delocalisation: the enthalpy of hydrogenation of benzene is about 152 kJ mol⁻¹ less exothermic than three times that of cyclohexene, and benzene resists addition reactions that would destroy the delocalised system.

29.4Optical isomerism

A chiral centre is a carbon with four different groups. The two enantiomers are non-superimposable mirror images that rotate plane-polarised light in opposite directions and are otherwise physically identical. A 50:50 mixture — a racemic mixture — shows no net rotation.

30.1Arenes

Benzene undergoes electrophilic substitution, not addition, because substitution preserves the stable delocalised π system.

Directing effects. Electron-donating groups (–OH, –NH₂, –CH₃) activate the ring and direct to the 2- and 4- (ortho/para) positions. Electron-withdrawing groups (–NO₂, –COOH) deactivate and direct to the 3- (meta) position.

Side-chain chemistry: methylbenzene's side chain is oxidised by hot alkaline KMnO₄ (then acidified) to benzoic acid, and undergoes free-radical substitution with Cl₂ in UV light — a useful contrast with ring substitution using a halogen carrier in the dark.

31.1Halogen compounds (A2)

Explain the relative inertness of aryl and vinyl halides: the halogen's lone pair overlaps with the ring's (or the C=C's) π system, giving the C–X bond partial double-bond character. It is shorter and stronger, and the carbon is less δ+, so nucleophilic substitution is far harder than for a halogenoalkane. Chlorobenzene does not give a precipitate with warm aqueous AgNO₃; chloroalkanes do.

CFCs and the ozone layer: UV light homolytically breaks the C–Cl bond, releasing Cl• radicals that catalyse ozone destruction — Cl• + O₃ → ClO• + O₂, then ClO• + O → Cl• + O₂. The chlorine radical is regenerated, so one radical destroys many ozone molecules.

32.1Alcohols (A2)

Extends the AS chemistry with the tri-iodomethane test as a structural probe and with acidity comparisons across alcohols, water, phenol and carboxylic acids.

32.2Phenol

Phenol is more acidic than an alcohol because a lone pair on the oxygen is delocalised into the ring, so the phenoxide ion formed on losing H⁺ is stabilised by delocalisation. It is less acidic than a carboxylic acid, which is why phenol reacts with NaOH but not with Na₂CO₃ — no effervescence. That single test distinguishes phenol from a carboxylic acid.

The same delocalisation makes the ring much more reactive towards electrophiles: phenol decolourises bromine water at room temperature with no catalyst, giving a white precipitate of 2,4,6-tribromophenol. Phenol also gives a violet colour with neutral FeCl₃(aq).

33.1Carboxylic acids (A2)

Explain relative acid strengths using inductive effects: electron-withdrawing groups near the COOH stabilise the carboxylate anion and increase acidity. So CCl₃COOH > CHCl₂COOH > CH₂ClCOOH > CH₃COOH, and chlorine closer to the COOH has a bigger effect than chlorine further away. Alkyl groups are electron-releasing, so they reduce acidity.

33.2Esters (A2)

Acid and alkaline hydrolysis as at AS, plus the use of esters in polyester formation (topic 35) and in fats and oils. Triglycerides hydrolysed with NaOH give soap (the sodium carboxylate) and glycerol.

33.3Acyl chlorides

Acyl chlorides are the most reactive carboxylic acid derivative: the carbonyl carbon is strongly δ+ (two electronegative atoms), and Cl⁻ is a good leaving group. They react vigorously at room temperature with:

34.1Primary and secondary amines

Alkyl groups are electron-releasing, so they increase the electron density on the nitrogen and make the lone pair more available to accept a proton — a stronger base. In phenylamine the lone pair is delocalised into the benzene ring, so it is far less available: phenylamine is a much weaker base than ammonia.

34.2Phenylamine and azo compounds

Preparation: nitrobenzene + Sn/concentrated HCl, reflux, then NaOH(aq) → phenylamine.

Diazotisation: phenylamine + HNO₂ (from NaNO₂ + HCl) at below 10 °C gives the benzenediazonium salt — unstable above that temperature, decomposing to phenol and nitrogen. Coupling the diazonium salt with phenol in alkaline solution gives a brightly coloured azo dye, whose colour comes from the extensive delocalisation across the –N=N– bridge joining the two rings.

34.3Amides

Formed from an acyl chloride with ammonia or an amine. Hydrolysed by acid (giving the carboxylic acid and an ammonium salt) or by alkali (giving the carboxylate and ammonia/amine). Amides are not basic in the way amines are — the nitrogen lone pair is delocalised onto the carbonyl oxygen.

34.4Amino acids

Amino acids are amphoteric — the COOH is acidic and the NH₂ basic. In the solid and at intermediate pH they exist as zwitterions (⁺H₃N–CHR–COO⁻), which is why they have unexpectedly high melting points and dissolve in water rather than in organic solvents.

The isoelectric point is the pH at which the zwitterion form predominates and there is no net charge. In acid the amino acid gains a proton and becomes a cation; in alkali it loses one and becomes an anion. Amino acids link through peptide (amide) bonds to form proteins.

35.1Condensation polymerisation

Monomers join with the loss of a small molecule, usually water or HCl. Polyesters (e.g. Terylene) form from a diol + a dicarboxylic acid; polyamides (nylon, Kevlar) from a diamine + a dicarboxylic acid, or from an amino acid. Be able to draw the repeat unit from the monomers, and — the more common exam task — to deduce the monomers by cutting the polymer at the ester or amide link.

35.2Predicting the type of polymerisation

Look at the monomer. A C=C and nothing else → addition. Two functional groups capable of condensing (diol/diacid, diamine/diacid, amino acid) → condensation. In a polymer chain, an ester or amide link in the backbone means condensation; a plain carbon backbone means addition.

35.3Degradable polymers

Condensation polymers can be hydrolysed at their ester or amide links, so they are biodegradable and can be chemically recycled to the monomers. Addition polymers, with inert C–C backbones, cannot. Photodegradable polymers include carbonyl groups that absorb UV and break the chain.

36.1Organic synthesis (A2)

Multi-step routes across the whole A Level, aliphatic and aromatic. The method: identify the change in functional group and the change in carbon number, work backwards from the target, and give reagents and conditions at every arrow.

You also need practical technique: reflux for reactions needing prolonged heating without loss of volatiles, distillation to remove a product as it forms, solvent extraction with a separating funnel, recrystallisation to purify a solid, and melting point determination to check purity — a sharp melting point at the literature value means pure; a low, broad range means impure.

37.1Thin-layer chromatography

Separation depends on the balance between adsorption on the stationary phase and solubility in the mobile phase — a partition/adsorption equilibrium. R_f values are constant for a given compound, stationary phase and solvent, so they identify components by comparison with standards.

37.2Gas / liquid chromatography

A volatile sample is carried by an inert gas through a column coated with a liquid stationary phase. Components emerge at characteristic retention times, and the peak area is proportional to the amount present. Retention time identifies; peak area quantifies. A limitation: two different compounds can share a retention time, so GC is often coupled to mass spectrometry for confirmation.

37.3Carbon-13 NMR spectroscopy

The number of peaks equals the number of chemically different carbon environments. Chemical shifts (relative to TMS at δ = 0) are given in a data table; you interpret rather than recall. There is no splitting to worry about, which makes ¹³C spectra the quicker of the two to read.

37.4Proton (¹H) NMR spectroscopy

Three pieces of information, and every question uses all three:

1. Number of peaks = number of different hydrogen environments.
2. Integration (relative peak area) = the ratio of hydrogens in each environment.
3. Splitting pattern — the n + 1 rule: a signal is split into (n + 1) lines by n hydrogens on the adjacent carbon.

Deuterium exchange: adding D₂O makes the O–H or N–H signal disappear, identifying it. CDCl₃ is used as a solvent because it contains no ordinary protons; TMS is the reference standard — inert, volatile, and with 12 equivalent protons giving a single sharp peak at δ = 0.

QA 1Reactions of cations

QA 2Reactions of anions, tests for gases and elements

Organic tests you may also be asked to carry out or interpret in Paper 3: Fehling's reagent (orange/red ppt → aldehyde); Tollens' reagent (silver mirror / black ppt → aldehyde); alkaline aqueous iodine (yellow ppt → CH₃CO or CH₃CH(OH) group); acidified KMnO₄ turning from purple to colourless (→ a compound that can be oxidised).

P3Paper 3: Advanced Practical Skills

Two or three questions totalling 40 marks. One is qualitative — investigating unknown substances, with the QA notes supplied. The others are quantitative — a titration, or measuring time, temperature, mass or gas volume, with a table or graph to draw, calculations to perform and conclusions to reach.

The precision rules Cambridge states explicitly: record burette readings to the nearest 0.05 cm³; with a thermometer calibrated at 1 °C intervals, record to the nearest 0.5 °C; a measuring cylinder calibrated at 1.0 cm³ is read to the nearest 0.5 cm³. Concordant titres means two titres within 0.10 cm³ of each other — and only concordant titres go into the mean.

Table conventions: a single table of results with headings and units in an accepted form — volume / cm³, volume (cm³) or volume in cm³. All raw readings of a quantity to the same number of decimal places, matching the instrument.

Recording observations: use simple colour words ("blue", "yellow"), and where fine discrimination is needed use "pale", "dark" and comparisons — "darker brown than at three minutes", "paler green than with 0.2 mol dm⁻³".

P5Paper 5: Planning, Analysis and Evaluation

Two or more questions totalling 30 marks, written, no laboratory. You may be asked to design an investigation, state a hypothesis linking independent and dependent variables, sketch the expected graph, or analyse and evaluate data you are given.

Planning checklist — work through it and you collect most of the marks:

1. State the independent variable, the dependent variable and the controlled variables explicitly.
2. Say how you will vary the independent variable, over what range, and with how many values.
3. Name the apparatus for measuring each quantity, chosen for appropriate precision — a burette rather than a measuring cylinder, a balance reading to 0.01 g.
4. Say how each controlled variable is kept constant (thermostatic water bath, same concentration of catalyst, same total volume).
5. Give a labelled diagram or a clear numbered procedure.
6. State the analysis: what you plot against what, and how the required quantity comes from the gradient or intercept. Show the rearrangement into y = mx + c.
7. Add a specific safety precaution with a reason — "wear gloves because concentrated H₂SO₄ is corrosive", not a generic list.

Analysis and evaluation: process the data into a table with headings and units, calculate the uncertainty in each processed value, plot with error bars, draw the best and worst acceptable lines, take gradients from both, and quote the answer as a value with an absolute uncertainty and a unit.

MapAliphatic conversions

MapAromatic conversions

LearnThe definitions examiners want verbatim

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Official (free)

• Cambridge International — 9701 subject page: syllabus, specimen papers, past papers, mark schemes and examiner reports.
• Examiner reports name the exact questions candidates got wrong each series — read them for every paper you attempt.

Free archives

• GCE Guide · PastPapers.co — full CAIE past-paper archives.
• Physics & Maths Tutor — topic-sorted questions.

How to revise this subject

1. Learn the qualitative analysis tables anyway. They are printed in Paper 3, but Papers 1, 2 and 4 can ask you to predict or explain an observation with no table in sight.
2. Moles first, always. Nearly every calculation in 9701 starts by converting a mass, volume or concentration into moles. Write "n = ..." as your first line and the rest usually follows.
3. Draw mechanisms properly. Curly arrows start at a lone pair or a bond, not at an atom, and point to where the electron pair goes. Show the dipoles, the intermediates and the lone pairs. Sloppy arrows lose marks even when the products are right.
4. Build the organic reaction map early. Every organic question is a path through it. Learn reagent + conditions + product as a single unit — "reagent" alone is half an answer.
5. State conditions. "Heat under reflux with acidified potassium dichromate(VI)" scores; "oxidise it" does not.
6. Explain with the right vocabulary. Ionic radius, shielding, nuclear charge, lattice energy, electronegativity, activation energy — examiners mark for the named concept, not for the story around it.