IGCSE Chemistry syllabus checklist

CoreEvery candidate

• 1.1State the distinguishing properties of solids, liquids and gases.
• 1.1Describe the structures of solids, liquids and gases in terms of particle separation, arrangement and motion.
• 1.1Describe changes of state in terms of melting, boiling, evaporating, freezing and condensing.
• 1.1Describe the effects of temperature and pressure on the volume of a gas.
• 1.2Describe and explain diffusion in terms of the kinetic particle theory.

SupplementExtended candidates only

• 1.1Explain changes of state in terms of the kinetic particle theory, including the interpretation of heating and cooling curves.
• 1.1Explain, using the kinetic particle theory, the effects of temperature and pressure on the volume of a gas.
• 1.2Describe and explain the effect of relative molecular mass on the rate of diffusion of gases.

CoreEvery candidate

• 2.1Describe the differences between elements, compounds and mixtures.
• 2.2Describe the atom as a central nucleus of protons and neutrons surrounded by electrons in shells.
• 2.2State the relative charges and the relative masses of a proton, a neutron and an electron.
• 2.2Define proton number (atomic number) as the number of protons in the nucleus.
• 2.2Define mass number (nucleon number) as the total number of protons and neutrons in the nucleus.
• 2.2Determine the electronic configuration of elements and their ions for proton numbers 1–20.
• 2.2State that the Group VIII noble gases have a full outer shell, that the number of outer-shell electrons equals the group number for Groups I–VII, and that the number of occupied shells equals the period number.
• 2.3Define isotopes as atoms of the same element that have the same number of protons but different numbers of neutrons.
• 2.3Interpret and use symbols for atoms and ions, such as ¹²₆C and ³⁵₁₇Cl⁻.
• 2.4Describe the formation of positive ions (cations) and negative ions (anions).
• 2.4State that an ionic bond is a strong electrostatic attraction between oppositely charged ions.
• 2.4Describe ionic bond formation between Group I and Group VII elements, using dot-and-cross diagrams.
• 2.4Describe the properties of ionic compounds: high melting and boiling points, and conduction when aqueous or molten but not when solid.
• 2.5State that a covalent bond is a shared pair of electrons between two atoms, giving both atoms the noble-gas configuration.
• 2.5Describe covalent bond formation in simple molecules including H₂, Cl₂, H₂O, CH₄, NH₃ and HCl, using dot-and-cross diagrams.
• 2.5Describe the properties of simple molecular compounds: low melting and boiling points, and poor electrical conductivity.
• 2.6Describe the giant covalent structures of graphite and diamond.
• 2.6Relate the structures and bonding of graphite and diamond to their uses: graphite as a lubricant and an electrode, and diamond in cutting tools.

SupplementExtended candidates only

• 2.3State that isotopes of an element have the same chemical properties because they have the same number of electrons and the same electronic configuration.
• 2.3Calculate relative atomic mass from the relative masses and abundances of an element’s isotopes.
• 2.4Describe the giant lattice structure of ionic compounds as regular alternating positive and negative ions.
• 2.4Describe ionic bond formation between ions of metallic and non-metallic elements, using dot-and-cross diagrams.
• 2.4Explain the melting and boiling points and the conductivity of ionic compounds in terms of their structure and bonding.
• 2.5Describe covalent bond formation in simple molecules including CH₃OH, C₂H₄, O₂, CO₂ and N₂, using dot-and-cross diagrams.
• 2.5Explain the low melting and boiling points of simple molecular compounds in terms of weak intermolecular forces, and explain their poor electrical conductivity.
• 2.6Describe the giant covalent structure of silicon(IV) oxide, SiO₂.
• 2.6Describe the similarity in properties between diamond and silicon(IV) oxide, linked to their structures.
• 2.7Describe metallic bonding as the electrostatic attraction between positive ions in a giant metallic lattice and a sea of delocalised electrons.
• 2.7Explain the electrical conductivity, malleability and ductility of metals in terms of their structure and bonding.

CoreEvery candidate

• 3.1State the formulae of the elements and compounds named in the subject content.
• 3.1Define the molecular formula as the number and type of atoms in one molecule.
• 3.1Deduce a simple compound’s formula from the relative numbers of atoms in a model or diagram.
• 3.1Construct word and symbol equations showing how reactants form products, including state symbols.
• 3.2Describe relative atomic mass, Aᵣ, as the average mass of an element’s isotopes compared with 1/12 the mass of a carbon-12 atom.
• 3.2Define relative molecular mass, Mᵣ, as the sum of the relative atomic masses of the atoms in a formula (relative formula mass for ionic compounds).
• 3.2Calculate reacting masses in simple proportions; these calculations do not involve the mole concept.
• 3.3State that concentration can be measured in g/dm³.

SupplementExtended candidates only

• 3.1Define the empirical formula as the simplest whole-number ratio of atoms or ions in a compound.
• 3.1Deduce the formula of an ionic compound from relative ion numbers, a model or diagram, or ion charges.
• 3.1Construct symbol equations with state symbols, including ionic equations.
• 3.1Deduce the symbol equation with state symbols for a reaction from the relevant information.
• 3.3State that the mole (mol) is the unit of amount of substance and contains 6.02 × 10²³ particles (the Avogadro constant).
• 3.3Use amount (mol) = mass (g) ÷ molar mass (g/mol) to find amount, mass, molar mass or relative atomic mass.
• 3.3Use the molar gas volume of 24 dm³ at room temperature and pressure.
• 3.3Calculate stoichiometric reacting masses, limiting-reagent amounts, gas volumes at r.t.p., and solution concentrations in mol/dm³ and g/dm³.
• 3.3Use titration data to calculate concentrations and reacting volumes.
• 3.3Calculate empirical formulae and molecular formulae.
• 3.3Calculate percentage yield, percentage purity and percentage composition by mass.

CoreEvery candidate

• 4.1Define electrolysis as the decomposition of an ionic compound, molten or in aqueous solution, by the passage of electricity.
• 4.1Identify, in simple cells, the anode, the cathode and the electrolyte.
• 4.1Identify the products at the electrodes in specified examples.
• 4.1State that metals or hydrogen form at the cathode and that non-metals form at the anode.
• 4.1Predict the products at each electrode in specified molten and aqueous cases.
• 4.1State that metal objects are electroplated to improve their appearance and to resist corrosion.
• 4.1Describe how metals are electroplated.
• 4.2State that a hydrogen–oxygen fuel cell uses hydrogen and oxygen to produce electricity, with water as the only product.

SupplementExtended candidates only

• 4.1Describe the transfer of charge during electrolysis: electrons in the metal and ions in the electrolyte.
• 4.1Construct ionic half-equations for the reactions at the electrodes.
• 4.1Identify the products at the electrodes in a wider range of specified cases.
• 4.1Predict the products at each electrode in a wider range of specified molten and aqueous cases.
• 4.2Describe the advantages and disadvantages of hydrogen–oxygen fuel cells.

CoreEvery candidate

• 5.1State that an exothermic reaction transfers thermal energy to the surroundings, so the temperature of the surroundings increases.
• 5.1State that an endothermic reaction takes in thermal energy from the surroundings, so the temperature of the surroundings decreases.
• 5.1Interpret reaction pathway diagrams that show the reactants, the products and the overall energy change.

SupplementExtended candidates only

• 5.1Explain that the thermal energy transferred during a reaction comes from the breaking and forming of bonds.
• 5.1Define activation energy, Eₐ, as the minimum energy that colliding particles need in order to react.
• 5.1Interpret reaction pathway diagrams that also show the activation energy.
• 5.1State that breaking a bond is endothermic and forming a bond is exothermic.
• 5.1Calculate the enthalpy change of a reaction from the bond energies of the bonds broken and the bonds formed.

CoreEvery candidate

• 6.1Identify physical and chemical changes and describe the differences between them.
• 6.2Describe how the surface area of a solid, the concentration of a solution, the temperature and a catalyst affect the rate of a reaction.
• 6.2State that a catalyst increases the rate of a reaction and is not used up.
• 6.2Describe practical methods for investigating the rate of a reaction, including the effects of concentration and temperature.
• 6.2Interpret data, including graphs, from rate experiments.
• 6.3State that some reactions are reversible and that the products can react to re-form the reactants.
• 6.3Describe how changing the conditions can change the direction of a reversible reaction.
• 6.4Use a Roman numeral to indicate the oxidation number of an element in a compound, such as Fe in Fe₂O₃ is III.
• 6.4Define redox reactions as reactions involving oxidation and reduction.
• 6.4Define oxidation as the gain of oxygen and reduction as the loss of oxygen.
• 6.4Identify redox reactions as reactions involving the gain or loss of oxygen.
• 6.4Identify oxidation and reduction in redox reactions.

SupplementExtended candidates only

• 6.2Describe collision theory in terms of particle collisions, activation energy and the effect of changing conditions.
• 6.2Describe and explain the effect on the rate of reaction of surface area, concentration, temperature and catalysts using collision theory.
• 6.2State that a catalyst decreases the activation energy of a reaction.
• 6.2Evaluate practical methods for investigating the rate of a reaction.
• 6.3State that a reversible reaction in a closed system reaches dynamic equilibrium, and predict and explain the effect of changing conditions.
• 6.3State the symbol equations for the Haber process (N₂ + 3H₂ ⇌ 2NH₃) and the Contact process (2SO₂ + O₂ ⇌ 2SO₃), the sources of the reactants and the typical industrial conditions.
• 6.3Explain the choice of conditions in the Haber and Contact processes in terms of rate and position of equilibrium.
• 6.4Define oxidation and reduction in terms of electron loss or gain, changes in oxidation number, and hydrogen loss or gain.
• 6.4Identify redox reactions by changes in oxidation number.
• 6.4Identify redox reactions by the colour changes seen with acidified aqueous potassium manganate(VII) (KMnO₄) or aqueous potassium iodide (KI).
• 6.4Define an oxidising agent and a reducing agent.
• 6.4Identify oxidising agents and reducing agents in reactions.

CoreEvery candidate

• 7.1Describe how acids react with metals, with bases and with carbonates.
• 7.1Describe acids by their effect on litmus, thymolphthalein and methyl orange.
• 7.1State that bases are the oxides and hydroxides of metals, and that alkalis are soluble bases.
• 7.1Describe how bases react with acids and with ammonium salts.
• 7.1Describe alkalis by their effect on litmus, thymolphthalein and methyl orange.
• 7.1State that aqueous acids contain H⁺(aq) ions and aqueous alkalis contain OH⁻(aq) ions.
• 7.1Compare hydrogen ion concentration, neutrality, relative acidity and relative alkalinity using colour and pH with universal indicator paper.
• 7.1Describe the neutralisation reaction between an acid and an alkali to produce water, H⁺ + OH⁻ → H₂O.
• 7.2Classify oxides as acidic, for example SO₂ and CO₂, or basic, for example CuO and CaO, related to metallic and non-metallic character.
• 7.3Describe the preparation, separation and purification of soluble salts by reacting an acid with an alkali by titration, or with excess metal, excess insoluble base or excess insoluble carbonate.
• 7.3Describe the general solubility rules for salts in water.
• 7.3Define a hydrated substance as one that is chemically combined with water, and an anhydrous substance as one containing no water.

SupplementExtended candidates only

• 7.1Define acids as proton donors and bases as proton acceptors.
• 7.1Define a strong acid as completely dissociated in aqueous solution and a weak acid as partially dissociated.
• 7.1State that hydrochloric acid is a strong acid, HCl(aq) → H⁺(aq) + Cl⁻(aq), and that ethanoic acid is a weak acid, CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq).
• 7.2Describe amphoteric oxides as oxides that react with both acids and bases to produce a salt and water.
• 7.2Classify aluminium oxide, Al₂O₃, and zinc oxide, ZnO, as amphoteric.
• 7.3Describe the preparation of insoluble salts by precipitation.
• 7.3Define water of crystallisation, including CuSO₄•5H₂O and CoCl₂•6H₂O.

CoreEvery candidate

• 8.1Describe the Periodic Table as an arrangement of elements in periods and groups, in order of increasing proton number.
• 8.1Describe the change from metallic to non-metallic character across a period.
• 8.1Describe the relationship between group number and the charge of the ions formed by elements in that group.
• 8.1Explain similarities in the chemical properties of elements in the same group in terms of their electronic configurations.
• 8.1Explain how the position of an element in the Periodic Table can be used to predict its properties.
• 8.2Describe lithium, sodium and potassium as relatively soft metals, with trends down the group of decreasing melting point, increasing density and increasing reactivity.
• 8.2Predict the properties of other Group I elements, given information about the group.
• 8.3Describe chlorine, bromine and iodine as diatomic non-metals, with trends down the group of increasing density and decreasing reactivity.
• 8.3State the appearance of the halogens at room temperature and pressure: chlorine a pale yellow-green gas, bromine a red-brown liquid, iodine a grey-black solid.
• 8.3Describe and explain the displacement reactions of halogens with other halide ions.
• 8.3Predict the properties of other Group VII elements, given information about the group.
• 8.4Describe transition elements as metals with high densities, high melting points, coloured compounds and which often act as catalysts.
• 8.5Describe the Group VIII noble gases as unreactive, monatomic gases, and explain this unreactivity in terms of their electronic configurations.

SupplementExtended candidates only

• 8.1Identify trends in groups, given information about the elements.
• 8.4Describe transition elements as having ions with variable oxidation numbers, including iron(II) and iron(III).

CoreEvery candidate

• 9.1Compare the general physical properties of metals and non-metals: thermal and electrical conductivity, malleability and ductility, and melting and boiling points.
• 9.1Describe the general chemical properties of metals: reactions with dilute acids, with cold water and steam, and with oxygen.
• 9.2Describe uses of metals linked to their physical properties: aluminium in aircraft (low density), aluminium in overhead cables (low density and good electrical conductivity), aluminium in food containers (resistance to corrosion), and copper in electrical wiring (good electrical conductivity and ductility).
• 9.3Describe an alloy as a mixture of a metal with other elements, including brass (copper and zinc) and stainless steel (iron with chromium, nickel and carbon).
• 9.3State that alloys can be harder and stronger than pure metals and are therefore more useful.
• 9.3Describe uses of alloys linked to physical properties, including stainless steel in cutlery because of its hardness and resistance to rusting.
• 9.3Identify representations of alloys from diagrams of their structure.
• 9.4State the order of the reactivity series: potassium, sodium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, gold.
• 9.4Describe the reactions, if any, of potassium, sodium and calcium with cold water; of magnesium with steam; and of magnesium, zinc, iron, copper, silver and gold with dilute hydrochloric acid, explaining these using the reactivity series.
• 9.4Deduce an order of reactivity from given experimental results.
• 9.5State the conditions required for the rusting of iron and steel to form hydrated iron(III) oxide.
• 9.5State common barrier methods of rust prevention: painting, greasing and coating with plastic.
• 9.5Describe how barrier methods prevent rusting by excluding oxygen or water.
• 9.6Describe the ease of obtaining metals from their ores in relation to the position of the metal in the reactivity series.
• 9.6Describe the extraction of iron from hematite in the blast furnace: burning of carbon (coke) to provide heat and carbon dioxide; reduction of carbon dioxide to carbon monoxide; reduction of iron(III) oxide by carbon monoxide; thermal decomposition of calcium carbonate; and formation of slag.
• 9.6State that the main ore of aluminium is bauxite and that aluminium is extracted by electrolysis.

SupplementExtended candidates only

• 9.3Explain, in terms of structure, how alloys are harder and stronger than pure metals: the different-sized atoms disturb the regular layers and prevent them sliding.
• 9.4Describe relative reactivities in terms of the tendency to form positive ions, using displacement reactions with aqueous ions of magnesium, zinc, iron, copper and silver.
• 9.4Explain the apparent unreactivity of aluminium in terms of its protective oxide layer.
• 9.5Describe the use of zinc in galvanising as a barrier method and as sacrificial protection.
• 9.5Explain sacrificial protection in terms of the reactivity series and the loss of electrons.
• 9.6State the symbol equations for the reactions in the blast furnace: C + O₂ → CO₂; C + CO₂ → 2CO; Fe₂O₃ + 3CO → 2Fe + 3CO₂; CaCO₃ → CaO + CO₂; CaO + SiO₂ → CaSiO₃.
• 9.6Describe the extraction of aluminium from purified bauxite (aluminium oxide), including the role of cryolite, why the carbon anodes need regular replacement, and the reactions at the electrodes including ionic half-equations.

CoreEvery candidate

• 10.1Describe chemical tests for water using anhydrous cobalt(II) chloride and anhydrous copper(II) sulfate.
• 10.1Describe how to test the purity of water using melting point and boiling point.
• 10.1Explain that distilled water is used in practical chemistry because it has fewer chemical impurities than tap water.
• 10.1State that natural water may contain dissolved oxygen, metal compounds, plastics, sewage, harmful microbes, nitrates from fertilisers, and phosphates from fertilisers and detergents.
• 10.1State that some of these substances are beneficial: dissolved oxygen for aquatic life, and some metal compounds as essential minerals.
• 10.1State that some are potentially harmful: toxic metal compounds, plastics harming aquatic life, sewage carrying disease-causing microbes, and nitrates and phosphates causing deoxygenation that damages aquatic life.
• 10.1Describe domestic water treatment: sedimentation and filtration to remove solids, use of carbon to remove tastes and odours, and chlorination to kill microbes.
• 10.2State that ammonium salts and nitrates are used as fertilisers.
• 10.2Describe the use of NPK fertilisers to supply nitrogen, phosphorus and potassium for improved plant growth.
• 10.3State the composition of clean, dry air: approximately 78% nitrogen, 21% oxygen, and the remainder a mixture of noble gases and carbon dioxide.
• 10.3State the sources of air pollutants: carbon dioxide from the complete combustion of carbon-containing fuels; carbon monoxide and particulates from incomplete combustion; methane from decomposing vegetation and animal digestion; oxides of nitrogen from car engines; sulfur dioxide from burning fossil fuels that contain sulfur compounds.
• 10.3State the adverse effects of these pollutants: carbon dioxide and methane raise global warming and drive climate change; carbon monoxide is toxic; particulates increase the risk of respiratory problems and cancer; oxides of nitrogen cause acid rain, photochemical smog and respiratory problems; sulfur dioxide causes acid rain.
• 10.3State and explain strategies to reduce the effects of these environmental issues: planting trees, reducing livestock farming, using less fossil fuel and using more hydrogen and renewable energy for climate change; catalytic converters, low-sulfur fuels and flue-gas desulfurisation with calcium oxide for acid rain.
• 10.3Describe photosynthesis as the reaction of carbon dioxide and water to produce glucose and oxygen, using light energy absorbed by chlorophyll, and state the word equation.

SupplementExtended candidates only

• 10.3Describe how carbon dioxide and methane cause global warming in terms of the absorption, reflection and emission of thermal energy, and the reduced loss of thermal energy to space.
• 10.3Explain how oxides of nitrogen form in car engines and how catalytic converters remove them, for example 2CO + 2NO → 2CO₂ + N₂.
• 10.3State the symbol equation for photosynthesis: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂.

CoreEvery candidate

• 11.1Draw and interpret displayed formulae, in which every atom and bond is shown.
• 11.1Write and interpret the general formulae of the homologous series: alkanes CₙH₂ₙ₊₂, alkenes CₙH₂ₙ, alcohols CₙH₂ₙ₊₁OH and carboxylic acids CₙH₂ₙ₊₁COOH.
• 11.1Identify a functional group as the atom or group of atoms that determines the chemical properties of a homologous series.
• 11.1State that a homologous series is a family of similar compounds with similar chemical properties, due to the same functional group.
• 11.1State that a saturated compound has molecules in which all carbon–carbon bonds are single bonds.
• 11.1State that an unsaturated compound has molecules containing one or more carbon–carbon bonds that are not single bonds.
• 11.2Name and draw the displayed formulae of methane, ethane, ethene, ethanol and ethanoic acid, and of the products of the reactions in topics 11.4–11.7.
• 11.2State the type of compound from an ending -ane, -ene, -ol or -oic acid, or from a molecular or displayed formula.
• 11.3Name the fossil fuels: coal, natural gas and petroleum.
• 11.3Name methane as the main constituent of natural gas.
• 11.3State that hydrocarbons are compounds that contain hydrogen and carbon only.
• 11.3State that petroleum is a mixture of hydrocarbons and describe its separation into useful fractions by fractional distillation.
• 11.3Describe the properties of fractions changing from the bottom to the top of the fractionating column: shorter chain length, higher volatility, lower boiling points and lower viscosity.
• 11.3Name the uses of the fractions: refinery gas for heating and cooking; gasoline fraction for cars; naphtha fraction as a chemical feedstock; kerosene fraction for jet fuel; diesel fraction for diesel engines; fuel oil for ships and home heating; lubricating oil for lubricants, waxes and polishes; bitumen for road surfaces.
• 11.4State that alkanes are saturated hydrocarbons in which all the bonds are single covalent bonds.
• 11.4Describe alkanes as generally unreactive, except in terms of combustion and substitution by chlorine.
• 11.5State that alkenes are unsaturated hydrocarbons containing a carbon–carbon double bond.
• 11.5Describe the manufacture of alkenes and of hydrogen by the cracking of larger alkanes, using high temperature and a catalyst.
• 11.5Describe the reasons for cracking, including the demand for shorter-chain alkenes and hydrogen.
• 11.5Describe the test to distinguish saturated from unsaturated hydrocarbons using aqueous bromine.
• 11.6Describe the manufacture of ethanol by fermentation of aqueous glucose at 25–35 °C using yeast and the absence of oxygen, and by the catalytic addition of steam to ethene at 300 °C and 6000 kPa (60 atmospheres) using an acid catalyst.
• 11.6Describe the combustion of ethanol.
• 11.6State the uses of ethanol as a solvent and as a fuel.
• 11.7Describe the reaction of ethanoic acid with metals, bases and carbonates, including the names and formulae of the salts produced.
• 11.8Define polymers as large molecules built up from many smaller molecules called monomers.
• 11.8Describe the formation of poly(ethene) as an example of addition polymerisation of ethene monomers.
• 11.8State that plastics are made from polymers.
• 11.8Describe how the properties of plastics influence the way they are used and disposed of.
• 11.8Describe the environmental challenges of plastics: disposal in landfill, accumulation in oceans, and the production of toxic gases when they are burned.

SupplementExtended candidates only

• 11.1State that a structural formula unambiguously describes the arrangement of atoms, for example CH₂=CH₂, CH₃CH₂OH and CH₃COOCH₃.
• 11.1Define structural isomers as compounds with the same molecular formula but different structural formulae, including examples with the formula C₄H₁₀ and C₄H₈.
• 11.1Describe the general characteristics of a homologous series: the same functional group, the same general formula, successive members differing by a –CH₂– unit, a trend in physical properties, and similar chemical properties.
• 11.2Name and draw the structural and displayed formulae of unbranched alkanes, alkenes including but-1-ene and but-2-ene, alcohols including propan-1-ol, propan-2-ol, butan-1-ol and butan-2-ol, and carboxylic acids containing up to four carbon atoms.
• 11.2Name and draw the displayed formulae of unbranched esters made from unbranched alcohols and carboxylic acids, each containing up to four carbon atoms.
• 11.4State that in a substitution reaction one atom or group is replaced by another atom or group.
• 11.4Describe the photochemical reaction of alkanes with chlorine, in which ultraviolet light provides the activation energy, and draw the products limited to monosubstitution.
• 11.5State that in an addition reaction only one product is formed.
• 11.5Describe the addition reactions of alkenes with bromine and aqueous bromine, with hydrogen using a nickel catalyst, and with steam using an acid catalyst, and draw the products.
• 11.6Describe the advantages and disadvantages of manufacturing ethanol by fermentation and by the catalytic addition of steam to ethene.
• 11.7Describe the formation of ethanoic acid by the oxidation of ethanol with acidified aqueous potassium manganate(VII), and by bacterial oxidation during vinegar production.
• 11.7Describe the reaction of a carboxylic acid with an alcohol, using an acid catalyst, to form an ester.
• 11.8Identify the repeat units and/or linkages in addition and condensation polymers.
• 11.8Deduce the repeat unit of an addition polymer from a given alkene, and vice versa.
• 11.8Deduce the repeat unit of a condensation polymer from given monomers, and vice versa, limited to polyamides made from a dicarboxylic acid and a diamine, and polyesters made from a dicarboxylic acid and a diol.
• 11.8Describe the differences between addition and condensation polymerisation.
• 11.8Describe and draw the structures of nylon, a polyamide, and PET, a polyester.
• 11.8State that PET can be converted back into monomers and re-polymerised.
• 11.8Describe proteins as natural polyamides formed from amino acid monomers with the general structure given in the syllabus, and describe and draw the protein structure.

CoreEvery candidate

• 12.1Name appropriate apparatus for the measurement of time, temperature, mass and volume, including stopwatches and the balance; thermometers; and burettes, volumetric pipettes, measuring cylinders and gas syringes.
• 12.1Suggest advantages and disadvantages of experimental methods and apparatus for a given task.
• 12.1Describe the terms solvent, solute, solution, saturated solution, residue and filtrate.
• 12.2Describe an acid–base titration, including the use of a burette, a volumetric pipette and a suitable indicator.
• 12.2Describe how to identify the end-point of a titration using an indicator.
• 12.3Describe how paper chromatography is used to separate mixtures of soluble coloured substances, using a suitable solvent.
• 12.3Interpret simple chromatograms, including the use of known substances to identify unknowns, and the identification of pure and impure substances.
• 12.4Describe and explain methods of separation and purification using: a suitable solvent; filtration; crystallisation; simple distillation; and fractional distillation.
• 12.4Suggest suitable separation and purification techniques, given information about the substances involved.
• 12.4Identify substances and assess their purity using melting point and boiling point information.
• 12.5Describe tests for the anions: carbonate (add dilute acid, then test for carbon dioxide); chloride, bromide and iodide (acidify with dilute nitric acid, then add aqueous silver nitrate); nitrate (add aluminium foil and aqueous sodium hydroxide, then test for ammonia); sulfate (acidify with dilute nitric acid, then add aqueous barium nitrate); sulfite (acidified aqueous potassium manganate(VII)).
• 12.5Describe tests using aqueous sodium hydroxide and aqueous ammonia to identify the aqueous cations Al³⁺, NH₄⁺, Ca²⁺, Cr³⁺, Cu²⁺, Fe²⁺, Fe³⁺ and Zn²⁺.
• 12.5Describe tests for the gases: ammonia (damp red litmus paper); carbon dioxide (limewater); chlorine (damp litmus paper); hydrogen (lighted splint); oxygen (glowing splint); and sulfur dioxide (acidified aqueous potassium manganate(VII)).
• 12.5Describe the flame test for the cations Li⁺, Na⁺, K⁺, Ca²⁺, Ba²⁺ and Cu²⁺.

SupplementExtended candidates only

• 12.3Describe how paper chromatography separates mixtures of soluble colourless substances, using a solvent and a locating agent.
• 12.3State and use the equation Rf = distance moved by substance ÷ distance moved by solvent.