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A-level Chemistry revision checklist

All 478 objectives in the curriculum, taken straight from the lessons that teach them. Open a unit and tap the dot beside an objective to rate how confident you are.

Jump to: Working in the laboratory · Atomic structure · Amount of substance · Bonding and structure · Energetics · Kinetics · Chemical equilibria · Redox · Thermodynamics · Rate equations · Electrochemistry · Acids and bases · Periodicity and Period 3 · Group 2 and Group 7 · Transition metals · Ions in aqueous solution · Introduction to organic chemistry · Alkanes and halogenoalkanes · Alkenes and alcohols · Organic analysis · Carbonyls and carboxylic acids · Aromatic chemistry and amines · Organic synthesis

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Working in the laboratory

Measurement, uncertainty and significant figures in chemistry AQA MS 1.1, MS 1.3; 8 (practical skills) · OCR A 1.1.3, 1.1.4 · Edexcel 5.12, 5.13 · CAIE Papers 3 and 5

State the uncertainty of a single reading on a burette, a pipette, a balance and a thermometer, and explain why a difference of two readings carries twice that.
Calculate a percentage uncertainty, and combine percentage uncertainties for a result calculated from several measurements.
Suggest how to reduce the percentage uncertainty of a measurement by changing the quantity measured or the apparatus.
Give a calculated answer to the number of significant figures the data justify.
Distinguish accuracy from precision, and random error from systematic error, with an example of each from a titration.
Rearrange an equation, convert units, use logarithms, and find a gradient and intercept from a straight-line graph.

Apparatus, techniques and safety: the twelve required practicals in outline AQA Practical skills 8 (AT a, b, d, e, g, h, k) · OCR A 1.1 · Edexcel 5.16, 18.21 · CAIE Paper 3

Choose between a pipette, a burette, a volumetric flask and a measuring cylinder for a given volume, and describe how each is used.
Describe weighing by difference and explain why it records the mass that actually reaches the reaction.
Describe heating under reflux and simple distillation, and explain when each is used.
Outline filtration under reduced pressure and recrystallisation for purifying a solid, and the melting point test of its purity.
Identify the hazards of concentrated acids, bromine, sodium and organic solvents, and give a control measure for each.
Name the lesson in which each of the twelve required practicals is taught.

Formulae, equations and state symbols AQA 3.1.2.5 · OCR A 2.1.2, 2.1.4 · Edexcel 3.10, 5.6, 5.15 · CAIE 2.3

Work out the formula of an ionic compound from the charges on its ions, using brackets for polyatomic ions.
Give the formulae of the elements that exist as diatomic molecules.
Balance an equation by inspection, changing only the numbers in front of formulae.
Add the correct state symbol to every species, and explain what (aq) means.
Use the solubility rules to identify a precipitate, and write an ionic equation by cancelling spectator ions.
Explain how electrons balance the charge in a half-equation.

Atomic structure

Protons, neutrons, electrons and isotopes AQA 3.1.1.1, 3.1.1.2 · OCR A 2.1.1 · Edexcel 1.1, 1.2, 1.3, 1.4, 1.5, 1.6, 1.15 · CAIE 1.1, 1.2

State the relative mass, relative charge and position in the atom of the proton, the neutron and the electron.
Use atomic number and mass number to find the numbers of protons, neutrons and electrons in an atom or an ion.
Define isotopes, and explain why the isotopes of an element have the same chemical properties but can differ in physical properties.
Deduce the charge on an ion from its numbers of protons and electrons, and write its symbol.
Predict the direction and relative size of the deflection of beams of protons, neutrons, electrons and ions in an electric field.
Describe how the model of the atom changed as new evidence appeared, from the plum pudding to the nuclear atom, energy levels and orbitals.

Time-of-flight mass spectrometry and relative atomic mass AQA 3.1.1.2 · OCR A 2.1.1 · Edexcel 1.8, 1.9, 1.10 · CAIE 2.1, 22.2

Describe the four stages of a time-of-flight mass spectrometer, including both electron impact and electrospray ionisation.
Explain why every ion leaves the acceleration stage with the same kinetic energy, and why lighter ions reach the detector first.
Calculate the flight time of an ion from KE = ½mv² and t = d/v, and the ratio of the flight times of two ions.
Calculate a relative atomic mass from the isotopic masses and abundances in a mass spectrum, and an abundance from a relative atomic mass.
Interpret the molecular ion peaks of chlorine and bromine as combinations of isotopes.
Use the M+1 peak to calculate the number of carbon atoms in a molecule, and the M+2 peak to identify chlorine or bromine in an organic compound.

Electron configuration: shells, subshells and orbitals AQA 3.1.1.3 · OCR A 2.2.1 · Edexcel 1.16, 1.17, 1.18, 1.19, 1.20, 1.21, 1.22, 1.23, 15.1 · CAIE 1.3

State the number of orbitals in s, p, d and f subshells and the number of electrons each subshell can hold.
Describe the shapes of s and p orbitals.
Write the full and shortened electron configurations of the atoms from hydrogen to krypton, filling 4s before 3d.
Apply the Pauli principle and Hund's rule to draw electrons in boxes and count the unpaired electrons in an atom or ion.
Write the configurations of ions, removing 4s electrons before 3d electrons, and give the configurations of chromium and copper.
Identify the block of an element from its electron configuration.

Ionisation energies and the evidence for subshells AQA 3.1.1.3 · OCR A 3.1.1 · Edexcel 1.11, 1.12, 1.13, 1.14, 1.15, 1.25 · CAIE 1.4

Define first ionisation energy, and write equations with state symbols for the first, second and third ionisations of an element.
Explain how nuclear charge, distance from the nucleus and shielding affect an ionisation energy.
Explain the decrease in first ionisation energy down a group.
Explain the general increase across Period 3, and the dips at aluminium and sulfur, in terms of subshell energy and spin-pair repulsion.
Use successive ionisation energies to deduce the group of an element and the number of electrons in each shell.
Calculate the energy needed to form a gaseous ion from successive ionisation energies, and use them to explain the charge an element's ion carries.

Amount of substance

The mole, the Avogadro constant and molar mass AQA 3.1.2.1, 3.1.2.2 · OCR A 2.1.3; PAG 1 · Edexcel 1.6, 1.7, 5.1, 5.2, 5.3, 5.7 · CAIE 2.1, 2.2, 2.3

Define the mole and the Avogadro constant, and use L to convert between an amount in moles and a number of particles.
Calculate relative molecular mass and relative formula mass from Ar values, and say which term applies to which kind of substance.
Use n = m ÷ M in either direction, with mass in grams and molar mass in g mol⁻¹.
Use the mole ratio in a balanced equation to calculate the mass of one substance from the mass of another.
Describe heating to constant mass, and calculate the number of molecules of water of crystallisation in a hydrated salt.

Gases and the ideal gas equation AQA 3.1.2.3 · OCR A 2.1.3 (f); PAG 1 · Edexcel 5.5, 5.7, 5.9, 5.10; core practical 1 · CAIE 2.4, 4.1

State the ideal gas equation and the SI unit of each quantity in it.
Convert volumes from cm³ and dm³ to m³, pressures from kPa to Pa, and temperatures from °C to K.
Calculate any one of p, V, n and T from the other three.
Use the molar gas volume of 24.0 dm³ mol⁻¹ at room temperature and pressure, and use gas volumes as reacting ratios.
Describe the gas syringe method, and calculate the Mr of a volatile liquid from its results.
Explain the conditions under which a real gas departs from ideal behaviour.

Empirical and molecular formulae AQA 3.1.2.4 · OCR A 2.1.3 · Edexcel 5.4, 5.5, 18.18 · CAIE 2.3

Define empirical formula and molecular formula, and distinguish both from a structural formula.
Calculate an empirical formula from the masses of the elements in a compound, or from its percentage composition by mass.
Recognise ratios such as 1 : 1.5 and 1 : 1.33 that must be multiplied up, not rounded.
Find a molecular formula from an empirical formula and a relative molecular mass.
Use the masses of carbon dioxide and water from combustion analysis to find the masses of carbon, hydrogen and oxygen in a sample.

Concentration, volumetric solutions and titration calculations AQA 3.1.2.5; RP1 · OCR A 2.1.3; PAG 2 · Edexcel 5.7, 5.11, 5.12; core practicals 2 and 3 · CAIE 2.4

Calculate concentrations in mol dm⁻³ and g dm⁻³, and use n = cV with the volume in dm³.
Describe how to make up a standard solution, and give the reason for each step.
Carry out an acid–base titration to concordant titres, and calculate the mean titre.
Calculate an unknown concentration from a titre, using the mole ratio in the balanced equation.
Calculate the result of a back titration, and the percentage uncertainty in a titre.
Predict the effect of a procedural error on the calculated concentration.

Reacting masses, percentage yield and atom economy AQA 3.1.2.5, 3.3.14 · OCR A 2.1.3 (h), 2.1.3 (j) · Edexcel 5.8, 5.14 · CAIE 2.4

Calculate the mass of a product from the mass of a reactant using a balanced equation.
Identify the limiting reagent from the amounts of two reactants.
Calculate percentage yield, and explain why it is below 100% in a real preparation.
Calculate atom economy from a balanced equation, and explain why addition reactions have an atom economy of 100%.
Distinguish atom economy from percentage yield, and evaluate a process using both.

Bonding and structure

Ionic bonding and ionic lattices AQA 3.1.3.1, 3.1.3.4 · OCR A 2.2.2 · Edexcel 2.1, 2.2, 2.3, 2.4, 2.5, 2.6, 2.27 · CAIE 3.2, 4.2

Describe ionic bonding as the electrostatic attraction between oppositely charged ions, and predict the charge on a simple ion from its group.
Draw dot-and-cross diagrams for ionic compounds such as NaCl, MgO and CaCl₂, showing outer electrons and the charge on each ion.
Describe the sodium chloride lattice, and explain why its formula is a ratio of ions and not a molecule.
Explain how the charge and the radius of the ions affect the strength of an ionic lattice, and use that to compare melting points.
Explain the trends in ionic radius down a group and along an isoelectronic series.
Explain why ionic compounds are brittle, have high melting points, and conduct electricity when molten or dissolved but not as solids.

Covalent, dative covalent and metallic bonding AQA 3.1.3.2, 3.1.3.3 · OCR A 2.2.2, 3.1.1 · Edexcel 2.7, 2.8, 2.9, 2.22 · CAIE 3.3, 3.4, 3.7, 12.1

Describe a covalent bond as a shared pair of electrons attracted to the nuclei of both bonded atoms.
Draw dot-and-cross diagrams for simple molecules, including double and triple bonds and lone pairs.
Explain how a dative covalent bond forms, show it with an arrow, and identify it in NH₄⁺, H₃O⁺ and Al₂Cl₆.
Describe metallic bonding as a lattice of positive ions in a sea of delocalised electrons.
Explain how ionic charge, ionic radius and the number of delocalised electrons affect the strength of metallic bonding.
Explain why metals conduct electricity, are malleable and have high melting points.

The four crystal structures and their properties AQA 3.1.3.4 · OCR A 2.2.2, 3.1.1 · Edexcel 2.18, 2.23, 2.24, 2.25, 2.26, 2.27 · CAIE 4.2

Name the four types of crystal structure and give an example of each.
Describe the structures of diamond, graphite and silicon dioxide, and explain their melting points, hardness and conductivity.
Explain why molecular substances such as iodine and ice have low melting points, naming what is overcome on melting.
Explain why ice is less dense than liquid water.
Deduce the structure of an unknown solid from its melting point, conductivity and solubility.

Shapes of molecules and ions AQA 3.1.3.5 · OCR A 2.2.2 · Edexcel 2.10, 2.11, 2.12 · CAIE 3.5, 13.3

State the principle of electron-pair repulsion, and the order lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.
Deduce the number of bonding pairs and lone pairs round the central atom of a molecule or ion.
Name the shape and give the bond angles for molecules and ions with two to six electron pairs, with and without lone pairs.
Explain why the bond angle falls from 109.5° to 107° to 104.5° along CH₄, NH₃ and H₂O.
Draw molecules in three dimensions with wedge and dash bonds, showing lone pairs and a labelled bond angle.

Electronegativity, polarity and intermolecular forces AQA 3.1.3.6, 3.1.3.7 · OCR A 2.2.2 · Edexcel 2.13, 2.14, 2.15, 2.16, 2.17, 2.18, 2.19, 2.20, 2.21 · CAIE 3.1, 3.6

Define electronegativity, describe its trends in the periodic table and use it to explain why a bond is polar.
Explain why some molecules with polar bonds, such as CO₂ and CCl₄, are not polar.
Describe the origin of induced dipole–dipole forces, permanent dipole–dipole forces and hydrogen bonds, and put them in order of strength.
Explain the boiling points of the halogens, the alkanes and the hydrides of Groups 4 to 7, including the anomalous values for NH₃, H₂O and HF.
Draw hydrogen bonding between molecules, showing partial charges, the lone pair and the 180° angle at hydrogen.
Explain the solubility of a substance in water in terms of the intermolecular forces it can form.

Energetics

Enthalpy change, standard conditions and the definitions AQA 3.1.4.1 · OCR A 3.2.1 · Edexcel 8.1, 8.2, 8.3, 8.4 · CAIE 5.1

Define enthalpy change, and give the sign of ΔH for an exothermic and an endothermic reaction with the reason for each.
Draw and label an enthalpy profile diagram for each kind of reaction, marking ΔH and the activation energy.
State the standard conditions and identify the standard state of an element or compound.
Define the standard enthalpies of formation, combustion, neutralisation and reaction in the words a mark scheme uses.
Write the equation that represents each standard enthalpy change, with one mole of the right substance and every state symbol.
Calculate the heat released when a given mass of fuel burns, from its standard enthalpy of combustion.

Measuring an enthalpy change: calorimetry AQA 3.1.4.2 · OCR A 3.2.1, PAG 3 · Edexcel 6.38, 8.5, 8.6, 12.23 · CAIE 5.1

Use q = mcΔT, stating what each symbol is and the unit it takes.
Calculate ΔH in kJ mol⁻¹ from a temperature change, giving the sign and an appropriate number of significant figures.
Describe how to measure the enthalpy change of a reaction in solution, including the cooling-curve extrapolation.
Describe how to measure an enthalpy of combustion with a spirit burner.
Explain why calorimetry values are smaller than data-book values, and suggest improvements that bring them closer.
Explain why the enthalpy of neutralisation of a weak acid is less exothermic than that of a strong acid.

Hess's law and enthalpy cycles AQA 3.1.4.3 · OCR A 3.2.1 · Edexcel 8.5, 8.7, 8.8 · CAIE 5.2

State Hess's law and explain why it must hold.
Draw an enthalpy cycle using enthalpies of formation, with the arrows pointing up from the elements, and calculate ΔrH° from it.
Draw an enthalpy cycle using enthalpies of combustion, with the arrows pointing down to the combustion products, and calculate ΔrH° or ΔfH° from it.
Use a cycle to find an enthalpy change that cannot be measured directly, including the hydration of anhydrous copper(II) sulfate.
Evaluate a value calculated from a cycle against the experimental errors in the measured steps.
Calculate the enthalpy change of a thermal decomposition from two measured reactions with acid, scaling each step by its amount.

Mean bond enthalpies AQA 3.1.4.4 · OCR A 3.2.1 · Edexcel 8.9, 8.10, 8.11 · CAIE 5.1, 5.2

Define bond dissociation enthalpy and mean bond enthalpy, and explain why most tabulated values are means.
Explain why bond breaking is endothermic and bond making is exothermic.
Calculate ΔH for a reaction from mean bond enthalpies as bonds broken minus bonds made.
Calculate an unknown bond enthalpy from a measured ΔH.
Explain why ΔH from bond enthalpies differs from a value found by calorimetry or from enthalpies of formation.
Use enthalpies of formation and atomisation in a cycle to calculate a mean bond enthalpy, and relate bond length to bond strength.

Kinetics

Collision theory, activation energy and the rate of reaction AQA 3.1.5.1, 3.1.5.3, 3.1.5.4 · OCR A 3.2.2 · Edexcel 9.1, 9.2, 9.3, 9.6 · CAIE 8.1, 8.2

Define rate of reaction and give its units.
State the two conditions a collision must meet to lead to a reaction, and explain why most collisions do not.
Define activation energy and mark it, with ΔH, on the enthalpy profile of an exothermic and an endothermic reaction.
Explain the effect of concentration, gas pressure and surface area on the rate of a reaction in terms of the frequency of collisions.
Explain why the rate falls as a reaction proceeds, and calculate a rate from the gradient of a tangent to a graph.

The Maxwell–Boltzmann distribution, temperature and catalysts AQA 3.1.5.2, 3.1.5.3, 3.1.5.5 · OCR A 3.2.2 · Edexcel 9.1, 9.4, 9.5, 9.6, 9.7, 9.8, 9.9, 16.1 · CAIE 8.2, 8.3

Sketch the Maxwell–Boltzmann distribution correctly, marking the most probable energy, the mean energy and the activation energy.
State what the area under the curve, and the area beyond Ea, represent.
Sketch the distribution at a higher temperature on the same axes, and explain why a small rise in temperature gives a large rise in rate.
Explain how a catalyst increases a rate, using an enthalpy profile and the distribution with two activation energies marked.
Compare the effects of a rise in temperature and of a catalyst, and explain why industry uses catalysts.

Measuring the rate of a reaction AQA 3.1.5.3 (RP3) · OCR A 3.2.2; PAG 9, PAG 10 · Edexcel 9.3, 16.3 · CAIE 8.1

Choose a method to follow a given reaction, by gas volume, mass loss, colour or a precipitate, and justify the choice with a calculation where one helps.
Explain why 1/t is proportional to the rate when t is the time taken to reach a fixed end point.
Describe the practical that measures the effect of temperature on the rate of the reaction between sodium thiosulfate and hydrochloric acid, including its hazards.
Calculate 1/t from a results table, plot it against temperature and describe the trend.
Evaluate the errors in a rate experiment and suggest improvements that reduce them.
Calculate a rate in mol dm⁻³ s⁻¹ from a clock reaction's results, using the fixed amount of product formed at the end point.

Chemical equilibria

Dynamic equilibrium, Le Chatelier's principle and industrial compromise AQA 3.1.6.1 · OCR A 3.2.3 · Edexcel 10.1, 10.2, 10.3 · CAIE 7.1

State the features of a dynamic equilibrium, and explain why one can only be set up in a closed system.
Use Le Chatelier's principle to predict the effect of a change in concentration, pressure or temperature on the position of equilibrium.
Explain why a catalyst does not change the position of equilibrium but does change the time taken to reach it.
Explain the temperature and pressure chosen for the Haber process as a compromise between yield, rate and cost.
Apply the same reasoning to the Contact process, the hydration of ethene and the synthesis of methanol.

The equilibrium constant Kc AQA 3.1.6.2 · OCR A 3.2.3 (f) and (g), 5.1.2 · Edexcel 10.4, 11.2, 11.3, 11.4, 11.5 · CAIE 7.1, 25.2

Write an expression for Kc from a balanced equation for a homogeneous system, and deduce its units.
Calculate Kc from equilibrium amounts in a known volume, using an ICE table when only the initial amounts and one equilibrium amount are given.
Use a value of Kc to calculate an equilibrium amount.
Explain why changes in concentration and pressure and the addition of a catalyst leave Kc unchanged, and predict from the sign of ΔH how Kc changes with temperature.
Compare a reaction quotient with Kc to predict which way a mixture will move.
Write Kc for a heterogeneous equilibrium, leaving out solids and a solvent present in large excess.Edexcel · OCR A

Kp: equilibria in the gas phase AQA 3.1.10 · OCR A 5.1.2 · Edexcel 11.1, 11.2, 11.3, 11.4, 11.5 · CAIE 7.1

Calculate the mole fraction and the partial pressure of each gas in a mixture.
Write an expression for Kp for a gas-phase equilibrium and deduce its units.
Calculate Kp from partial pressures, and from initial amounts, one equilibrium amount and the total pressure.
Use a degree of dissociation, or a value of Kp, to find the composition of an equilibrium mixture.
Explain why Kp changes with temperature but not with the total pressure or a catalyst.
Write Kp for an equilibrium with a solid present, and convert a value of Kp between atmospheres and kilopascals.CAIE · Edexcel · OCR A

Redox

Oxidation states and the rules for assigning them AQA 3.1.7 · OCR A 2.1.5 · Edexcel 3.1, 3.2, 3.3, 3.4, 3.5, 3.6, 3.8, 3.9, 3.10, 3.11, 3.12, 4.3, 4.12, 14.1, 14.2 · CAIE 6.1

Define oxidation and reduction in terms of electrons and in terms of oxidation state.
Apply the rules for oxidation states in their order of priority, including the exceptions for peroxides, OF₂ and metal hydrides.
Calculate the oxidation state of an element in a compound or a polyatomic ion from the sum rule.
Deduce the oxidation state of a metal in a complex ion, counting each ligand at its own charge.
Identify the oxidising agent and the reducing agent in a redox reaction, and state what happens to each.
Use Roman numerals to name compounds and ions such as iron(III) chloride, manganate(VII) and chlorate(I).

Half-equations and building redox equations AQA 3.1.7, 3.2.3.2 · OCR A 2.1.5, 3.1.3, 5.2.3 · Edexcel 3.7, 3.8, 3.13, 14.1 · CAIE 6.1, 11.4, 24.2

Write a half-equation for a simple oxidation or reduction, with the electrons on the correct side.
Construct a half-equation in acidic solution by balancing atoms, then oxygen with H₂O, hydrogen with H⁺ and charge with electrons.
Combine two half-equations into an overall redox equation by matching the numbers of electrons, and balance an equation directly from changes in oxidation number.
Write a half-equation in alkaline solution by balancing as in acid and then adding OH⁻ to both sides.
Identify the species oxidised and the species reduced in an overall equation, and recognise spectator ions.
Explain disproportionation using the reactions of chlorine with water and with cold dilute sodium hydroxide.

Redox titrations: manganate(VII) and thiosulfate AQA 3.2.5.5 · OCR A 5.2.3 · Edexcel 14.18, 14.19 · CAIE 28.2, 33.1

Describe how to carry out a manganate(VII) titration, including why it needs no indicator and why it is acidified with dilute sulfuric acid.
Use the 1 : 5 and 2 : 5 ratios to calculate the amount of iron(II) or ethanedioate from a manganate(VII) titre.
Describe an iodine–thiosulfate titration, including the colour changes and the point at which starch is added.
Calculate the mass of iron in a tablet and the percentage of copper in an alloy from titration results.
Calculate a percentage purity, and apply the same chain to a redox system given in a question.
Explain how a thiosulfate solution is standardised against potassium iodate(V), and how dichromate(VI) differs from manganate(VII) as a titrant.

Thermodynamics

Lattice enthalpy and Born–Haber cycles AQA 3.1.8.1 · OCR A 5.2.1 · Edexcel 2.14, 4.6, 13.1, 13.2, 13.3, 13.4, 13.5, 13.6, 13.7, 13.8, 13.11 · CAIE 23.1, 27.1

Define lattice enthalpy of formation and lattice enthalpy of dissociation, and give the sign of each.
Define the enthalpy of atomisation, the first and second ionisation energies and the first and second electron affinities, and write an equation with state symbols for each.
Construct a Born–Haber cycle for a Group 1 or Group 2 halide or oxide as an energy-level diagram, labelling every level with its species.
Calculate a lattice enthalpy, or any other single missing term, from a Born–Haber cycle.
Explain why an experimental lattice enthalpy can be more exothermic than the value from the perfect ionic model, using polarisation and covalent character.

Enthalpies of solution and hydration AQA 3.1.8.1 · OCR A 5.2.1 · Edexcel 2.21, 13.9, 13.10, 13.11, 13.14 · CAIE 23.2

Define enthalpy of solution and enthalpy of hydration, and write an equation with state symbols for each.
Describe how water molecules are arranged around a cation and around an anion, and explain why hydration is always exothermic.
Explain how the charge and the radius of an ion affect its enthalpy of hydration.
Calculate an enthalpy of solution, a hydration enthalpy or a lattice enthalpy from the other terms in the solution cycle.
Explain how the balance between lattice enthalpy and hydration enthalpy decides a trend in solubility, and why an endothermic enthalpy of solution does not stop a salt dissolving.
Explain why the entropy change on dissolving is positive for most salts but negative for some salts of small, highly charged ions.

Entropy and Gibbs free-energy change AQA 3.1.8.2 · OCR A 5.2.2 · Edexcel 13.12, 13.13, 13.14, 13.15, 13.16, 13.17, 13.18, 13.19, 13.20, 13.21, 13.22 · CAIE 23.3, 23.4

Describe entropy as a measure of disorder, and predict the sign of an entropy change from the states and the moles of gas in an equation.
Calculate the entropy change of a reaction from standard entropies, in J K⁻¹ mol⁻¹.
Use ΔG = ΔH − TΔS to calculate a Gibbs free-energy change, converting the units correctly, and state the condition for a reaction to be feasible.
Calculate the temperature at which a reaction becomes feasible, and read ΔH, ΔS and that temperature from a graph of ΔG against T.
Explain why a reaction that is feasible may not take place at a measurable rate.
Calculate the entropy change of the surroundings and the total entropy change, and use ΔG = −RT ln K to relate ΔG to the equilibrium constant.CAIE · Edexcel

Rate equations

Rate equations, orders of reaction and the rate constant AQA 3.1.9.1 · OCR A 5.1.1 · Edexcel 16.1, 16.2, 16.6 · CAIE 26.1

Write a rate equation from given orders, and identify the order with respect to each reactant and the overall order.
Explain why the orders in a rate equation cannot be read from the balanced equation.
Calculate a rate constant from a rate and the concentrations, and deduce its units for any overall order from zero to three.
Sketch the rate–concentration and concentration–time graphs for a reactant that is zero, first or second order, and identify the order from the shape of each.
Calculate the new rate when the concentrations of the reactants are changed by given factors.

Finding orders from initial rates, half-lives and graphs AQA 3.1.9.2, RP7 · OCR A 5.1.1 (b) to (e); PAG 9, PAG 10 · Edexcel 16.3, 16.4, 16.5, 16.6, 16.7, 16.8; core practicals 13a and 13b · CAIE 26.1

Deduce the order with respect to each reactant from a table of initial rates, including runs in which two concentrations change.
Explain why 1/t from a clock reaction is proportional to the initial rate.
Find a rate from the tangent to a concentration–time graph, and use successive half-lives to identify the order.
Explain the isolation method, and why a run with the other reactants in large excess gives the order with respect to one reactant only.
Describe how to follow the iodine and propanone reaction by colorimetry and by titration, and explain the purpose of each step.
Evaluate the errors in that practical and suggest improvements.

The Arrhenius equation and the rate-determining step AQA 3.1.9.1, 3.1.9.2 · OCR A 5.1.1 (i) to (k) · Edexcel 6.34, 16.1, 16.8, 16.9, 16.10, 16.11, 16.12; core practical 14 · CAIE 26.1

Explain how the Arrhenius equation accounts for the effect of temperature and of activation energy on the rate constant.
Calculate the activation energy from the gradient of a graph of ln k against 1/T, and the pre-exponential factor from a point on the line.
Calculate the activation energy from rate constants at two temperatures.
Deduce the rate equation from a mechanism with a given slow step, and propose a mechanism consistent with a given rate equation.
Explain the different rate equations for the hydrolysis of primary and tertiary halogenoalkanes in terms of the SN2 and SN1 mechanisms.

Electrochemistry

Electrode potentials and the standard hydrogen electrode AQA 3.1.11.1 · OCR A 5.2.3 · Edexcel 14.3, 14.4, 14.5, 14.6, 14.7, 14.8, 14.9 · CAIE 24.2

Describe a half-cell, and explain how a potential difference arises between a metal and a solution of its ions.
Describe the standard hydrogen electrode and state the standard conditions under which E° is measured.
Explain the function of the salt bridge and of a high-resistance voltmeter in measuring the EMF of a cell.
Write the conventional representation of a cell, including half-cells that need a platinum electrode.
Calculate the EMF of a cell from two standard electrode potentials using E°cell = E°right − E°left.
Describe how to measure the EMF of a cell in the laboratory, and identify the errors that affect the reading.

The electrochemical series and predicting feasibility AQA 3.1.11.1 · OCR A 5.2.3 · Edexcel 13.22, 14.9, 14.10, 14.11, 14.12, 14.13, 14.14 · CAIE 24.2

Use a table of E° values to rank species as oxidising and reducing agents.
Predict whether a redox reaction is feasible from the sign of E°cell, and write the overall equation from two half-equations.
Use the anticlockwise rule to read the direction of a reaction from a pair of half-equations.
Explain why a reaction predicted to be feasible may not be observed, in terms of activation energy and non-standard conditions.
Predict how a change in concentration alters an electrode potential and the EMF of a cell, using Le Chatelier's principle.
Calculate an electrode potential with the Nernst equation and ΔG° from E°cell, and relate E°cell to ln K.CAIE · Edexcel

Commercial cells, rechargeable cells and fuel cells AQA 3.1.11.2 · OCR A 5.2.3 · Edexcel 14.15, 14.16, 14.17 · CAIE 24.1

Distinguish non-rechargeable, rechargeable and fuel cells by what happens to their reactants.
Write the electrode reactions of the lithium-ion cell during discharge and during charging, and calculate its EMF from given data.
Write the half-equations and overall equation for the hydrogen–oxygen fuel cell in alkaline and in acidic conditions, and explain why the EMF is the same in both.
Deduce the electrode reactions and EMF of an unfamiliar cell from electrode potential data.
Evaluate the benefits and risks to society of lithium-ion cells and of hydrogen fuel cells.
Calculate the mass liberated in electrolysis from Q = It and the Faraday constant.CAIE

Acids and bases

Brønsted–Lowry acids and bases, pH and strong acids AQA 3.1.12.1, 3.1.12.2 · OCR A 5.1.3 · Edexcel 12.1, 12.2, 12.3, 12.4, 12.5, 12.6, 12.7, 12.8, 12.14 · CAIE 7.2, 25.1

Define a Brønsted–Lowry acid and base, and identify the conjugate acid–base pairs in an equation.
Distinguish a strong acid from a weak acid by the extent of dissociation, and strength from concentration.
Write ionic equations for the reactions of acids with metals, carbonates and oxides, and state the basicity of an acid.
Calculate pH from [H⁺] and [H⁺] from pH, giving pH to two decimal places.
Calculate the pH of a strong monoprotic acid, of sulfuric acid, and of a strong acid after dilution.
Calculate the pH of the solution left when a strong acid is mixed with a strong base, with either in excess.

Weak acids, Ka and pKa; Kw and strong bases AQA 3.1.12.3, 3.1.12.4 · OCR A 5.1.3 · Edexcel 12.9, 12.10, 12.11, 12.12, 12.13, 12.15 · CAIE 25.1

State the expression for Kw and its value at 298 K, and explain why it increases with temperature.
Calculate the pH of a strong base, and of pure water at a temperature other than 298 K.
Write the expression for Ka of a weak acid with its units, and convert between Ka and pKa.
Calculate the pH of a weak acid from Ka, stating the two approximations used and when they fail.
Calculate Ka and pKa from the concentration and measured pH of a weak acid.

Titration curves and choosing an indicator AQA 3.1.12.5 · OCR A 5.1.3 · Edexcel 12.14, 12.16, 12.17, 12.22 · CAIE 25.1

Sketch the pH curves for the four combinations of strong and weak acid and base, with the start pH, vertical section and equivalence pH of each.
Explain why the pH at equivalence is above 7 for a weak acid with a strong base and below 7 for a strong acid with a weak base.
Show that pH = pKa at the half-neutralisation point, and use it to find Ka from a pH curve.
Choose an indicator whose range lies on the vertical section, and explain why none suits a weak acid–weak base titration.
Carry out and evaluate the pH-curve practical, including the calibration of a pH meter.

Buffer solutions and buffer calculations AQA 3.1.12.6 · OCR A 5.1.3 · Edexcel 12.18, 12.19, 12.20, 12.21, 12.24 · CAIE 25.1

Define a buffer solution, and describe how acidic and basic buffers are made.
Explain, with equations, how an acidic buffer and a basic buffer resist a change in pH when a small amount of acid or base is added.
Calculate the pH of an acidic buffer from Ka and the concentrations or amounts of the weak acid and its salt.
Calculate the pH of an acidic buffer after a small amount of strong acid or strong base has been added.
Outline how the carbonic acid–hydrogencarbonate system controls the pH of blood.
Write and use the expression for Ksp, and explain the common ion effect on the solubility of a salt.CAIE

Periodicity and Period 3

Periodicity: trends across Period 3 and the blocks of the table AQA 3.2.1.1, 3.2.1.2 · OCR A 3.1.1 · Edexcel 1.13, 1.22, 1.24, 1.25, 1.26 · CAIE 9.1

Classify an element as s, p, d or f block from its electron configuration or its position in the table.
Describe and explain the trend in atomic radius across Period 3.
Explain the general rise in first ionisation energy across Period 3 and the falls at aluminium and sulfur.
Explain the melting points of the Period 3 elements in terms of structure and bonding, including the sharp fall from silicon to phosphorus.
Explain why Period 3 positive ions are smaller, and negative ions larger, than their atoms.
Apply the same explanations to the trends across Period 2.

The Period 3 elements: reactions with oxygen and water AQA 3.2.4 · CAIE 9.2

Describe the reactions of sodium and magnesium with water, with equations and the pH of the resulting solution.
Explain why magnesium reacts so much faster with steam than with cold water, and why the product is different.
Write equations for the reactions of sodium, magnesium, aluminium, silicon, phosphorus and sulfur with oxygen, and describe what is seen.
Deduce the oxidation state of the Period 3 element in each oxide, and relate the highest oxidation state to the number of outer electrons.
Calculate gas volumes and reacting masses for these reactions, and deduce an oxide's formula from combustion data.
Describe the reactions of sodium to phosphorus with chlorine, and relate the oxidation state in each chloride to the number of outer electrons.CAIE

The Period 3 oxides: structure, bonding and acid–base character AQA 3.2.4 · CAIE 9.2

Explain the melting points of the Period 3 oxides in terms of their structure and bonding.
Write equations for the reactions of Na₂O, MgO, P₄O₁₀, SO₂ and SO₃ with water, and give the pH of each resulting solution.
Explain why the ionic oxides are basic and the covalent oxides acidic, and why Al₂O₃ and SiO₂ do not react with water.
Describe the structures of the acids and anions formed when P₄O₁₀, SO₂ and SO₃ react with water.
Write equations for the reactions of the oxides with acids and bases, including both reactions of amphoteric aluminium oxide.
Describe and explain the reactions of the Period 3 chlorides with water, with the pH of each solution, from their bonding.CAIE

Group 2 and Group 7

Group 2: trends down the group and reactions with water AQA 3.2.2 · OCR A 3.1.2 (a), 3.1.2 (b), 3.1.2 (d) · Edexcel 4.1, 4.2, 4.3, 4.4 · CAIE 10.1

State the electron configuration of each Group 2 element and explain why the oxidation state is +2 in every compound.
Explain the trends in atomic radius and first ionisation energy from magnesium to barium in terms of shells, shielding and nuclear charge.
Describe the trend in melting point, identify magnesium as the exception, and explain the general fall using metallic bonding.
Write equations with state symbols for the reactions of the Group 2 metals with water and of magnesium with steam, and describe what is seen.
Explain why magnesium is used to extract titanium, writing the equation and identifying what is oxidised and what is reduced.
Write equations for the metals with oxygen and dilute acids and for the oxides with water, and explain why the pH rises down the group.CAIE · Edexcel · OCR A

Group 2 compounds: solubility, uses and the sulfate test AQA 3.2.2 · OCR A 3.1.2 (e); 3.1.4 · Edexcel 4.4, 4.5, 4.6, 4.7, 4.8, 4.15 · CAIE 10.1

State the trends in solubility of the Group 2 hydroxides and sulfates, naming the least soluble compound in each series.
Give a use for magnesium hydroxide, calcium hydroxide, calcium oxide, calcium carbonate and barium sulfate, and write the equation that explains each.
Describe the test for sulfate ions, and explain why the barium chloride solution is acidified and why sulfuric acid cannot be the acid.
Identify Mg²⁺, Ca²⁺, Sr²⁺ and Ba²⁺ in solution from their reactions with sodium hydroxide and with dilute sulfuric acid.
Calculate reacting masses for the neutralisation and desulfurisation reactions.
Describe and explain the trend in thermal stability of the Group 2 carbonates and nitrates, and compare it with Group 1.CAIE · Edexcel

Group 7: trends down the group and displacement reactions AQA 3.2.3.1 · OCR A 3.1.3 · Edexcel 4.9, 4.10, 4.11, 4.12, 4.14 · CAIE 11.1, 11.2

Describe the state and colour of each halogen at room temperature, and explain the trend in boiling point using van der Waals forces.
Explain the trends in atomic radius and electronegativity down Group 7.
Explain why oxidising power decreases from fluorine to iodine.
Predict the outcome of mixing a halogen with a halide solution, and write the ionic equation and half-equations with state symbols.
Describe the colours of chlorine, bromine and iodine in water and in cyclohexane, and explain why the organic solvent is added.
Explain the trends in bond enthalpy of the halogens and in the thermal stability of the hydrogen halides.CAIE

Halide ions, the silver nitrate test, concentrated sulfuric acid, and chlorine in water AQA 3.2.3.1, 3.2.3.2 · OCR A 3.1.3, 3.1.4 · Edexcel 4.12, 4.13, 4.15 · CAIE 11.3, 11.4

Explain why the reducing ability of the halide ions increases from fluoride to iodide.
Describe and explain the reactions of solid sodium chloride, bromide and iodide with concentrated sulfuric acid, with equations, observations and the oxidation states of sulfur.
Describe the silver nitrate test, including the ammonia test, and explain why the silver nitrate is acidified with nitric acid.
Write equations for chlorine with water, in the dark and in sunlight, and with cold dilute sodium hydroxide, identifying disproportionation.
Evaluate the use of chlorine in treating drinking water.
Carry out test-tube tests to identify Group 2 ions, ammonium, hydroxide, carbonate, sulfate and halide ions in solution.

Transition metals

Transition metal characteristics and complex formation AQA 3.2.5.1, 3.2.5.2 · OCR A 5.3.1 · Edexcel 15.1, 15.2, 15.3, 15.4, 15.5, 15.6, 15.11, 15.12, 15.17, 15.18, 15.19 · CAIE 28.1, 28.2

State the definition of a transition metal and explain why scandium and zinc are excluded.
Write the electron configurations of the atoms titanium to copper and of their common ions, removing the 4s electrons first.
Name the four characteristic properties of transition metals.
Define ligand, coordinate bond, complex and coordination number, and deduce the charge or the oxidation state in a complex.
Distinguish monodentate, bidentate and multidentate ligands, with a named example of each.
Explain how haem carries oxygen and why carbon monoxide is toxic.

Ligands, the shapes of complexes, isomerism and ligand substitution AQA 3.2.5.2, 3.2.5.3 · OCR A 5.3.1 · Edexcel 15.12, 15.13, 15.14, 15.15, 15.16, 15.24, 15.26, 15.27, 15.28 · CAIE 28.2, 28.4, 28.5

Describe the octahedral, tetrahedral, square planar and linear shapes of complexes, with a named example and the bond angle of each.
Write equations for ligand substitution by ammonia and by chloride ions, with the colours and any change in coordination number.
Explain the chelate effect in terms of the number of particles and the entropy change.
Write the expression for a stability constant and use log₁₀ Kstab values to compare the stabilities of complexes.CAIE
Draw cis and trans isomers of octahedral and square planar complexes, and the optical isomers of a complex with three bidentate ligands.
Describe how cisplatin acts against cancer, and weigh its benefit against its side effects.

Colour in transition metal complexes and colorimetry AQA 3.2.5.4 · OCR A 5.3.1 (colour as a property only) · Edexcel 15.7, 15.8, 15.9, 15.10 · CAIE 28.3

Explain why transition metal complexes are coloured, using d-orbital splitting, the ground and excited states, and complementary colours.
Use ΔE = hν = hc/λ to calculate the energy, frequency or wavelength of the light absorbed.
Explain how a change of oxidation state, ligand or coordination number changes ΔE and the colour, with an example of each.
Explain why the compounds of Sc³⁺, Zn²⁺ and Cu⁺ are colourless.
Describe how to find the concentration of a coloured ion with a colorimeter and a calibration curve.

Variable oxidation states: vanadium, chromium and manganate(VII) titrations AQA 3.2.5.5 · OCR A 5.3.1 · Edexcel 15.3, 15.20, 15.21, 15.22, 15.23, 15.24 · CAIE 28.1, 28.2

Describe the reduction of vanadium(V) to vanadium(II) by zinc in acid, with the species, oxidation states and colours.
Use standard electrode potentials to explain how far a reducing agent can reduce a transition metal ion.
Explain how pH and the ligand change the stability of an oxidation state, with iron(II) hydroxide in air and Tollens' reagent as examples.
Describe a manganate(VII) titration of Fe²⁺ or ethanedioate, including the choice of acid and the end point.
Calculate an amount, a concentration, a formula or a percentage from the results of a manganate(VII) titration.

Catalysis: heterogeneous, homogeneous and autocatalytic AQA 3.2.5.6 · OCR A 5.3.1 · Edexcel 6.13, 9.8, 15.29, 15.30, 15.31, 15.32, 15.33, 15.34, 15.35 · CAIE 26.2

Describe heterogeneous catalysis as adsorption, reaction and desorption, and explain why the strength of adsorption matters.
Write the two steps by which vanadium(V) oxide catalyses the Contact process, and explain the part variable oxidation states play.
Explain how a support medium reduces cost, and how lead and sulfur poison catalysts.
Explain why the reaction between peroxodisulfate and iodide ions is slow, and how Fe²⁺ catalyses it, with equations.
Describe and explain autocatalysis by Mn²⁺, including the shape of its concentration–time curve.AQA · Edexcel

Ions in aqueous solution

Metal-aqua ions and the acidity of their solutions AQA 3.2.6 · Edexcel 12.14 · CAIE 9.2

Write the formulae of the hexaaqua ions of iron(II), copper(II), iron(III) and aluminium, and give the colour of each solution.
Write the hydrolysis equilibrium of a metal-aqua ion with state symbols, identifying the aqua ion as a Brønsted–Lowry acid.
Explain, in terms of charge/size ratio, why [M(H₂O)₆]³⁺ is more acidic than [M(H₂O)₆]²⁺.
Calculate the pH of a metal salt solution from Ka, Ka from a measured pH, and the change in pH when the solution is diluted.
Describe and explain the reactions of 2+ and 3+ aqua ions with sodium carbonate, and use them to tell the two charges apart.

Reactions of aqua ions with hydroxide, ammonia and carbonate AQA 3.2.6 · OCR A 5.3.1 (j), 5.3.2 · Edexcel 15.24, 15.25, 15.26 · CAIE 9.2, 28.2

Describe the reactions of the iron(II), copper(II), iron(III) and aluminium aqua ions with sodium hydroxide, ammonia and sodium carbonate, giving each observation in the words the mark scheme uses.
Write each reaction as an acid–base reaction of the aqua ion, balanced and with state symbols.
Explain why aluminium hydroxide dissolves in excess sodium hydroxide, and state what amphoteric means.
Explain why copper(II) hydroxide dissolves in excess ammonia, and write the equation for the ligand substitution.
Describe the reactions of the chromium(III), manganese(II) and cobalt(II) aqua ions with sodium hydroxide and ammonia.Edexcel · OCR A
Carry out test-tube reactions to identify a metal ion in solution, and identify an unknown from its observations.

Introduction to organic chemistry

Naming and representing organic molecules AQA 3.3.1.1 · OCR A 4.1.1 · Edexcel 6.1, 6.2, 6.3, 6.4, 6.9 · CAIE 13.1, 13.3

Write the empirical, molecular, general, structural, displayed and skeletal formula of the same compound, and state what each one shows.
Define a homologous series and a functional group, and identify the functional group in a compound from its formula.
Apply the IUPAC rules to name alkanes, alkenes, halogenoalkanes and alcohols with up to six carbons in the main chain.
Name aldehydes, ketones, carboxylic acids, esters, nitriles, amines, acyl chlorides and amides, including a molecule with two functional groups.
Draw a structure from a name, and use it to correct a name that does not follow the rules.

Structural isomerism and E/Z isomerism AQA 3.3.1.3 · OCR A 4.1.1, 4.1.3 · Edexcel 6.6, 6.7 · CAIE 13.4

Define structural isomers and stereoisomers, and tell the two apart from a pair of formulae.
Identify chain, position and functional group isomers, and draw all the structural isomers of a small molecular formula.
Explain how restricted rotation about a C=C bond gives rise to E/Z isomerism, and state the condition each carbon must meet.
Apply the Cahn–Ingold–Prelog priority rules to assign E or Z to an alkene, including one where the first atoms tie.
Explain why cis and trans agree with Z and E only in simple cases, and describe how the shape of an isomer changes its physical properties.

Reaction types, curly arrows and how a mechanism is written AQA 3.3.1.2 · OCR A 4.1.1 · Edexcel 6.5, 6.15, 6.21, 6.23, 6.31 · CAIE 13.2

Classify an organic reaction as addition, substitution, elimination, oxidation, reduction, hydrolysis, condensation or polymerisation.
Describe homolytic and heterolytic fission, and state what each produces.
Define a radical, a nucleophile and an electrophile, and identify each from its formula.
Explain what a curly arrow and a fishhook arrow represent, and where each may start and end.
Draw a two-arrow nucleophilic substitution mechanism with its dipole, lone pair and charges, and check that the charges balance.

Alkanes and halogenoalkanes

Alkanes: fractional distillation, cracking and combustion AQA 3.3.2.1, 3.3.2.2, 3.3.2.3 · OCR A 4.1.2 · Edexcel 2.20, 6.8, 6.9, 6.10, 6.11, 6.12, 6.13, 6.16 · CAIE 12.1, 14.1

Explain why the boiling points of alkanes rise with chain length and fall with branching, in terms of van der Waals forces.
Describe how crude oil is separated by fractional distillation, and say where in the column each fraction condenses.
Compare thermal and catalytic cracking: conditions, intermediates, products and the economic reason for each.
Write equations for complete and incomplete combustion, and give the origin of each pollutant from an engine or a power station.
Write equations for the reactions in a catalytic converter and in flue-gas desulfurisation, and outline the greenhouse effect.
Outline reforming, and explain the part nitrogen oxides play in photochemical smog and acid rain.CAIE · Edexcel

Free-radical substitution AQA 3.3.2.4 · OCR A 4.1.2 (f), 4.1.2 (g) · Edexcel 6.15, 6.16, 6.17 · CAIE 14.1

State the conditions for the reaction of methane with chlorine, and explain why ultraviolet light is needed.
Write equations for the initiation, propagation and termination steps, with the radical dot on the correct atom.
Explain why the reaction is a chain reaction, using the enthalpy changes of the two propagation steps.
Explain why a mixture of products forms, and why free-radical substitution is a poor way to make one halogenoalkane.
Apply the mechanism to another alkane or another halogen, including the termination products it predicts.
Use bond enthalpies to explain why fluorine reacts explosively with methane, bromine slowly and iodine not at all.

Halogenoalkanes: nucleophilic substitution and rates of hydrolysis AQA 3.3.3.1 · OCR A 4.2.1 (e); 4.2.2 (a), 4.2.2 (b); 6.2.1; 6.2.4 · Edexcel 6.30, 6.31, 6.32, 6.33, 6.34, 6.35, 6.36, 17.5 · CAIE 15.1, 19.2

Explain why the carbon of a C–X bond is attacked by nucleophiles, and define a nucleophile.
Outline the mechanisms for the reactions of a primary halogenoalkane with hydroxide ions, cyanide ions and ammonia, with the conditions for each.
Name the alcohol, nitrile or amine formed, and explain why excess ammonia is used.
Explain why the rate of hydrolysis depends on the C–X bond enthalpy rather than the bond polarity.
Describe how silver nitrate in ethanol is used to compare rates of hydrolysis, with the colour of each precipitate.
Explain why tertiary halogenoalkanes are hydrolysed fastest, and how optical activity distinguishes SN1 from SN2.CAIE · Edexcel

Elimination from halogenoalkanes, and CFCs and the ozone layer AQA 3.3.3.2, 3.3.3.3 · OCR A 4.2.2 (e) · Edexcel 6.32 · CAIE 15.1

State the conditions for elimination of HBr from a halogenoalkane, and outline the mechanism with hydroxide acting as a base.
Explain how the conditions and the type of halogenoalkane decide whether substitution or elimination is the main reaction.
Deduce the alkenes formed by elimination from an unsymmetrical halogenoalkane such as 2-bromobutane, including E/Z isomers.
Explain how ozone in the stratosphere is formed and why it matters.
Write the equations by which chlorine radicals from CFCs catalyse the decomposition of ozone, and evaluate the replacements for CFCs.

Alkenes and alcohols

Alkenes: structure, electrophilic addition and Markovnikov's rule AQA 3.3.4.1, 3.3.4.2 · OCR A 4.1.3 · Edexcel 6.18, 6.19, 6.20, 6.21, 6.22, 6.24, 6.25 · CAIE 13.3, 14.2

Describe the C=C double bond as a σ bond and a π bond, and explain why it makes an alkene attract electrophiles.
Outline the electrophilic addition mechanism for the reactions of alkenes with hydrogen bromide, bromine and concentrated sulfuric acid, with every curly arrow, dipole, lone pair and charge.
Explain how a bromine molecule, which is non-polar, becomes an electrophile as it approaches a double bond.
Explain the order of stability of primary, secondary and tertiary carbocations, and use it to predict the major product from an unsymmetrical alkene.
Describe the bromine water test for unsaturation and state its result.
Describe the oxidation of alkenes by cold dilute and by hot concentrated acidified potassium manganate(VII).CAIE · Edexcel

Addition polymers AQA 3.3.4.3 · OCR A 4.1.3 · Edexcel 6.26, 6.27, 6.28, 6.29 · CAIE 20.1, 35.3

Draw the repeating unit of an addition polymer from its monomer, with brackets, extended bonds and n.
Identify the monomer from a section of an addition polymer chain.
Explain the properties and uses of poly(ethene), poly(propene), PVC and PTFE in terms of their chains and the forces between them.
Explain how a plasticiser makes PVC flexible.
Explain why polyalkenes are unreactive and non-biodegradable, and evaluate landfill, incineration and recycling as ways of dealing with the waste.
Explain how a photodegradable polymer breaks down, and compare it with a biodegradable one.CAIE · OCR A

Alcohols: hydration, fermentation and classification AQA 3.3.5.1, 3.3.5.2 · OCR A 4.2.1 · Edexcel 6.14, 6.37, 6.38 · CAIE 16.1

Classify an alcohol as primary, secondary or tertiary from its structure or its name.
Explain the boiling points and solubilities of alcohols in terms of hydrogen bonding.
State and justify the conditions for making ethanol by fermentation and by the hydration of ethene, and outline the mechanism of acid-catalysed hydration.
Compare the two routes to ethanol on rate, purity, cost, energy and renewability.
Write the equations that support the claim that bioethanol is carbon neutral, and give the reasons the claim is not fully valid.
Describe the reaction of an alcohol with sodium, and explain why ethanol is a weaker acid than water.CAIE

Oxidation and elimination of alcohols AQA 3.3.5.2, 3.3.5.3 · OCR A 4.2.1 · Edexcel 6.38, 6.39, 17.11 · CAIE 15.1, 16.1

Write equations, using [O], for the oxidation of primary and secondary alcohols, and explain why tertiary alcohols are not oxidised.
Explain how distillation gives an aldehyde and heating under reflux gives a carboxylic acid from the same primary alcohol, and describe both sets of apparatus.
Describe the distillation of an aldehyde from a reaction mixture, and the errors that affect its yield and purity.
Outline the mechanism of the acid-catalysed elimination of water from an alcohol.
Deduce the alkenes formed by dehydrating an unsymmetrical alcohol, including E and Z isomers.
Write equations for the conversion of alcohols into halogenoalkanes by hydrogen halides and phosphorus halides.CAIE · Edexcel · OCR A

Organic analysis

Test-tube reactions for functional groups, and high-resolution mass spectrometry AQA 3.3.6.1, 3.3.6.2 · OCR A 4.2.4, 6.1.2, 6.3.1 · Edexcel 1.10, 6.25, 7.1, 19.1 · CAIE 17.1, 22.2

State the reagent, the conditions and the observation for the tests for an alkene, an aldehyde, a primary or secondary alcohol and a carboxylic acid.
Explain why an aldehyde reduces Tollens' reagent and Fehling's solution and a ketone does not.
Identify an unknown compound from the results of a sequence of tests, taking account of reagents that respond to more than one group.
Carry out the functional group tests practical safely, recording observations in the words a mark scheme uses.
Calculate a precise relative molecular mass and use it to decide between molecular formulae with the same nominal mass.
Describe the 2,4-dinitrophenylhydrazine test for a carbonyl group, and the use of the derivative's melting point.CAIE · Edexcel · OCR A

Infrared spectroscopy AQA 3.3.6.3 · OCR A 4.2.4 · Edexcel 7.2 · CAIE 22.1

Explain why a covalent bond absorbs infrared radiation of a particular wavenumber.
Identify the functional groups in a compound from its infrared spectrum and the data-sheet table.
Distinguish an alcohol, an aldehyde or ketone, a carboxylic acid and an ester from their absorptions above 1500 cm⁻¹.
Describe how the fingerprint region is used to identify a compound by comparison with a database.
Explain how absorption of infrared by carbon dioxide, methane and water vapour contributes to global warming.

Proton and carbon-13 NMR spectroscopy AQA 3.3.15 · OCR A 6.3.2 · Edexcel 19.2, 19.3, 19.4, 19.5 · CAIE 37.3, 37.4

Explain why TMS is used as the standard and why the solvent contains no hydrogen-1 atoms.
Deduce the number of carbon environments in a molecule and assign carbon-13 shifts with the data sheet.
Use the chemical shift, the integration ratio and the splitting of each peak in a proton NMR spectrum to describe the hydrogen environments.
Apply the n + 1 rule to predict the splitting of a peak from the number of hydrogens on the neighbouring carbons.
Deduce the structure of a compound from its molecular formula and its NMR spectra.

Chromatography: TLC, column and gas AQA 3.3.16 · OCR A 6.3.1 · Edexcel 19.6, 19.7, 19.8 · CAIE 37.1, 37.2

Explain how a mixture is separated by the balance between adsorption to the stationary phase and solubility in the mobile phase.
Calculate Rf values from a thin-layer chromatogram and use them to identify components against standards.
Carry out the thin-layer chromatography practical, separating a mixture by TLC and locating colourless spots.
Describe column chromatography and gas chromatography, and use retention times and peak areas to identify and compare components.
Explain the advantage of GC–MS over GC alone, and state the limits of chromatography for identifying a compound.
Describe high-performance liquid chromatography, and use a calibration curve of peak area against concentration.Edexcel

Carbonyls and carboxylic acids

Optical isomerism AQA 3.3.7 · OCR A 6.2.2 · Edexcel 17.1, 17.2, 17.3, 17.4, 17.5 · CAIE 13.4, 29.4

Identify a chiral centre in a structure as a carbon atom bonded to four different atoms or groups.
Draw the two enantiomers of a chiral molecule in three dimensions, with wedge and dashed bonds, as mirror images.
Describe how a pair of enantiomers affects plane-polarised light.
Explain why a racemic mixture is optically inactive.
Explain why nucleophilic addition to a planar carbonyl group gives a racemic mixture.
Suggest why a drug is made as a single enantiomer, using thalidomide and ibuprofen as examples.

Aldehydes and ketones: nucleophilic addition and the tests that tell them apart AQA 3.3.8 · OCR A 6.1.2 · Edexcel 17.6, 17.7, 17.8, 17.12 · CAIE 16.1, 17.1, 19.2, 29.4

Explain why the carbonyl group is polar and why its carbon is attacked by nucleophiles.
Describe the Tollens' and Fehling's tests, the observations for an aldehyde and a ketone, and write the half-equations.
Outline the nucleophilic addition mechanism for the reduction of an aldehyde or ketone by NaBH₄.
Outline the nucleophilic addition mechanism for the reaction of KCN followed by dilute acid, and name the hydroxynitrile formed.
Explain why the product of either addition can be a racemic mixture, and state the hazard of using cyanide.
Describe the 2,4-dinitrophenylhydrazine test for a carbonyl compound and, for Edexcel and Cambridge, the triiodomethane test for a CH₃CO group.CAIE · Edexcel · OCR A

Carboxylic acids, esters and biodiesel AQA 3.3.9.1 · OCR A 6.1.3, 6.2.2, 6.2.4 · Edexcel 6.14, 17.9, 17.10, 17.11, 17.12, 17.13, 17.15 · CAIE 16.1, 18.1, 18.2, 19.2, 33.1

Explain why carboxylic acids are weak acids, and write equations for their reactions with carbonates, bases and metals.
Describe the esterification of a carboxylic acid with an alcohol, and name the ester formed.
Give uses of esters linked to their properties.
Compare the hydrolysis of an ester in acid with its hydrolysis in alkali.
Describe how soap and biodiesel are made from a triglyceride, with equations.
Write equations for the hydrolysis of a nitrile to a carboxylic acid in acid and in alkali.CAIE · Edexcel · OCR A

Acylation: acyl chlorides and acid anhydrides AQA 3.3.9.2 · OCR A 6.1.3 · Edexcel 17.12, 17.13, 17.14, 18.8, 18.9, 18.13 · CAIE 29.2, 32.1, 33.1, 33.2, 33.3, 34.1, 34.3

Name acyl chlorides, acid anhydrides, amides and N-substituted amides from their formulae.
Write equations for the reactions of acyl chlorides and acid anhydrides with water, alcohols, ammonia and primary amines, and state the observations.
Outline the nucleophilic addition–elimination mechanism for an acyl chloride with water, an alcohol, ammonia or a primary amine.
Explain why acyl chlorides react faster than carboxylic acids.
Evaluate the use of ethanoic anhydride rather than ethanoyl chloride in the manufacture of aspirin.
Write equations for making an acyl chloride from a carboxylic acid with SOCl₂ or PCl₅.CAIE · Edexcel · OCR A

Condensation polymers and the disposal of polymers AQA 3.3.12.1, 3.3.12.2 · OCR A 6.2.3 · Edexcel 17.16, 18.14, 18.15 · CAIE 35.1, 35.2, 35.3

Describe condensation polymerisation and compare it with addition polymerisation.
Draw the repeating unit of a polyester or polyamide from its monomers, and identify the monomers from a repeating unit.
Explain the strength of nylon and Kevlar in terms of hydrogen bonding between chains.
Explain why polyesters and polyamides can be hydrolysed and are biodegradable, while polyalkenes are not.
Evaluate landfill, incineration and recycling as ways of disposing of polymers.

Aromatic chemistry and amines

Benzene: its structure and the evidence for delocalisation AQA 3.3.10.1 · OCR A 6.1.1 (a) to (c), 6.1.1 (h) · Edexcel 18.1, 18.2, 18.3, 18.4 · CAIE 29.1, 29.3, 30.1, 31.1, 32.2

Describe the Kekulé structure of benzene and state the three pieces of evidence that show it is wrong.
Describe the delocalised model of benzene: a planar ring of carbon atoms, p orbitals overlapping above and below it, and six delocalised electrons.
Calculate the delocalisation energy of benzene from enthalpies of hydrogenation, and explain what it shows about the stability of the ring.
Explain why benzene does not undergo the addition reactions that alkenes undergo readily.
Name substituted benzenes, using numbers to give the positions of two or more substituents.
Explain why phenol is a stronger acid than ethanol yet does not react with carbonates, and, for Cambridge, why chlorobenzene resists hydrolysis.CAIE · OCR A

Electrophilic substitution: nitration and Friedel–Crafts acylation AQA 3.3.10.2 · OCR A 6.1.1 (d), 6.1.1 (e), 6.1.1 (g), 6.1.1 (i) to (l), 6.2.4 (d) · Edexcel 18.4, 18.5, 18.6, 18.7 · CAIE 29.2, 30.1, 32.2, 34.2

Explain why benzene reacts with electrophiles by substitution and not by addition.
State the reagents and conditions for the nitration of benzene, and write equations for the formation of the nitronium ion.
Outline the mechanism of nitration with curly arrows, drawing the intermediate correctly, and show how the catalyst is regenerated.
State the reagents and conditions for Friedel–Crafts acylation, write the equation for the formation of the acylium ion, and outline the mechanism.
Describe why nitration and acylation are important in making explosives, dyes and new carbon–carbon bonds.
Explain why phenol substitutes more readily than benzene, and predict where a group already on the ring sends the next one.CAIE · OCR A

Amines: preparation, basicity and nucleophilic reactions AQA 3.3.11.1, 3.3.11.2, 3.3.11.3 · OCR A 6.2.1 (a), 6.2.1 (b) · Edexcel 18.8, 18.9, 18.10, 18.11, 18.12, 18.22 · CAIE 12.1, 19.1, 34.1, 34.2, 34.3

Identify primary, secondary and tertiary amines and quaternary ammonium salts, and describe the use of quaternary ammonium salts as cationic surfactants.
Explain why amines are bases, and explain the order of base strength of phenylamine, ammonia and primary aliphatic amines.
Outline the mechanism of the reaction between ammonia and a halogenoalkane, and explain why further substitution gives a mixture of products.
Describe the reactions of amines with acyl chlorides to give N-substituted amides.
Describe the preparation of aliphatic amines from halogenoalkanes and by reducing nitriles, and of phenylamine by reducing nitrobenzene.
Describe the hydrolysis and reduction of amides and the making of diazonium salts and azo dyes.CAIE

Amino acids, proteins, enzymes and DNA AQA 3.3.13.1 to 3.3.13.5 · OCR A 6.2.2 · Edexcel 15.16, 18.8, 18.15, 18.16, 18.17 · CAIE 34.4, 35.1

Describe the structure of an amino acid, identify its chiral centre, and draw the zwitterion and the ions it forms in acid and in alkali.
Explain how a peptide link forms by condensation, and give the two dipeptides made from two different amino acids.
Describe the primary, secondary and tertiary structure of a protein and the bonds that hold each in place.
Describe how a protein is hydrolysed and its amino acids identified by thin-layer chromatography, calculate Rf values, and predict how amino acids move in electrophoresis.
Explain how enzymes act as stereospecific catalysts and how a drug can act as an enzyme inhibitor.
Describe the structure of DNA, and explain how cisplatin prevents DNA replication and why its use carries risks.

Organic synthesis

Planning a synthetic route AQA 3.3.14 · OCR A 4.2.3, 6.2.4, 6.2.5 · Edexcel 17.11, 18.19, 18.20, 18.21 · CAIE 18.1, 19.2, 21.1, 36.1

State the reagents and conditions for each conversion between functional groups in the aliphatic and aromatic reactions of the course.
Devise a synthesis of two to four steps from a starting material to a target, naming every intermediate.
Identify when a route must add a carbon to the chain, and use a nitrile to do it.
Calculate the overall yield of a route with several steps and the atom economy of a step.
Evaluate two routes to the same product on yield, atom economy, hazard and the purity of the product.
Use a Grignard reagent to add a carbon: a carboxylic acid from carbon dioxide, or an alcohol from an aldehyde or ketone.Edexcel

Preparing, purifying and testing an organic solid and liquid AQA 3.3.9.2; RP10 · OCR A 4.2.3, 6.2.5; PAG 5, PAG 6 · Edexcel 6.39, 18.21, 18.22; CP16 · CAIE Papers 3 and 5 (practical skills)

Explain why organic reactions are heated under reflux, and when distillation is used instead.
Describe how an impure organic solid is purified by recrystallisation, giving the reason for each step.
Explain how a melting point shows whether a solid is pure, and describe how it is measured.
Describe how an organic liquid is purified by washing, drying and redistillation, giving the reason for each step.
Calculate the percentage yield of a preparation from masses or volumes and densities.

Putting the analysis together: structure from spectra AQA 3.3.15 · OCR A 4.2.4, 6.3.2 · Edexcel 18.18, 19.5 · CAIE 37.4

Deduce a molecular formula from a precise mass, or from combustion data and the molecular ion.
Calculate the degree of unsaturation from a molecular formula, and use it to limit the possible structures.
Identify the functional groups in an unknown from its infrared spectrum.
Deduce a structure from carbon-13 and proton NMR data, using shift, integration and splitting.
Explain how symmetry, O–H peaks and integration ratios can mislead an analysis, and check a structure against every piece of evidence.

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