G3 Pure Chemistry / O-Level Chemistry
Study the shared G3 Pure Chemistry / O-Level Chemistry course topic by topic, with qualification-specific exam guidance where needed.
Learning goals
- name appropriate apparatus for the measurement of time, temperature, mass and volume; including burettes, pipettes, measuring cylinders and gas syringes
- suggest suitable apparatus, given relevant information, for a variety of simple experiments, including drying and collection of gases and measurement of rates of reaction (drying agents will be limited to calcium oxide, concentrated sulfuric acid and fused calcium chloride).
- describe methods of separation and purification for the components of mixtures, to include: — use of a suitable solvent, filtration and crystallisation or evaporation
- describe methods of separation and purification for the components of mixtures, to include: — sublimation
- describe methods of separation and purification for the components of mixtures, to include: — distillation and fractional distillation (see also 11.1(b))
- describe methods of separation and purification for the components of mixtures, to include: — use of a separating funnel
- describe methods of separation and purification for the components of mixtures, to include: — paper chromatography
- suggest suitable separation and purification methods, given information about the substances involved in the following types of mixtures: — solid-solid
- suggest suitable separation and purification methods, given information about the substances involved in the following types of mixtures: — solid-liquid
- suggest suitable separation and purification methods, given information about the substances involved in the following types of mixtures: — liquid-liquid (miscible and immiscible)
- interpret paper chromatograms including comparison with ‘known’ samples and the use of Rf values
- explain the need to use locating agents in the chromatography of colourless compounds (knowledge of specific locating agents is not required)
- deduce from given melting point and boiling point data the identities of substances and their purity.
- explain the importance of measuring the purity in substances used in everyday life, e.g. foodstuffs and drugs.
- follow a sequence of instructions;
- use techniques, apparatus and materials;
- make and record observations, measurements and estimates;
- interpret and evaluate observations and experimental results;
- plan investigations, select techniques, apparatus and materials;
- evaluate methods and suggest possible improvements.
- describe the solid, liquid and gaseous states of matter and explain their interconversion in terms of the kinetic particle theory and of the energy changes involved.
- describe and explain evidence for the movement of particles in liquids and gases (the treatment of Brownian motion is not required)
- explain everyday effects of diffusion in terms of particles, e.g. the spread of perfumes and cooking aromas; tea and coffee grains in water
- state qualitatively the effect of molecular mass on the rate of diffusion and explain the dependence of rate of diffusion on temperature.
- state the relative charges and approximate relative masses of a proton, a neutron and an electron
- describe, with the aid of diagrams, the structure of an atom as consisting of protons and neutrons (nucleons) in the nucleus and electrons arranged in shells (energy levels) (knowledge of s, p, d and f classification is not required; a copy of the Periodic Table will be available in Papers 1 and 2)
- define proton (atomic) number and nucleon (mass) number
- interpret and use nuclide notations such as 126 C
- define the term isotopes
- deduce the numbers of protons, neutrons and electrons in atoms and ions given proton and nucleon numbers.
- describe the formation of ions by electron loss/gain and that these ions usually have the electronic configuration of a noble gas
- describe, including the use of ‘dot-and-cross’ diagrams, the formation of ionic bonds between metals and non-metals, e.g. NaCl; MgCl2
- state that ionic materials contain a giant lattice in which the ions are held by electrostatic attraction, e.g. NaCl (candidates will not be required to draw diagrams of ionic lattices)
- relate the physical properties (including electrical property) of ionic compounds to their lattice structure (see also 3.4(g)).
- describe the formation of a covalent bond by the sharing of a pair of electrons and that the atoms in the molecules usually have the electronic configuration of a noble gas
- describe, using ‘dot-and-cross’ diagrams, the formation of covalent bonds between non-metallic elements, e.g. H2; O2; H2O; CH4; CO2
- deduce the arrangement of electrons in other covalent molecules
- relate the physical properties (including electrical property) of covalent substances to their structure and bonding (see also 3.4(g)).
- describe metals as a lattice of positive ions in a ‘sea of electrons’
- describe the general physical properties of metals as solids having high melting and boiling points, malleable, good conductors of heat and electricity in terms of their structure (see also 3.4(g)).
- describe the differences between elements, compounds and mixtures
- describe an alloy as a mixture of a metal with another element, e.g. brass; stainless steel
- identify representations of metals and alloys from diagrams of structures
- explain why alloys have different physical properties to their constituent elements
- compare the structures of the following substances in order to deduce their properties: — simple molecular substances, e.g. methane, iodine
- compare the structures of the following substances in order to deduce their properties: — macromolecules, e.g. poly(ethene)
- compare the structures of the following substances in order to deduce their properties: — giant covalent substances, e.g. sand (silicon dioxide), diamond, graphite (see also 3.4(g))
- compare the bonding and structures of diamond and graphite in order to deduce their properties such as electrical conductivity, lubricating or cutting action (candidates will not be required to draw the structures)
- deduce the physical and chemical properties of substances from their structures and bonding and vice versa (see also 3.1(d), 3.2(d), 3.3(b) and 3.4(e)).
- state the symbols of the elements and formulae of the compounds mentioned in the syllabus
- deduce the formulae of simple compounds from the relative numbers of atoms present and vice versa
- deduce the formulae of ionic compounds from the charges on the ions present and vice versa
- interpret chemical equations with state symbols
- construct chemical equations, with state symbols, including ionic equations.
- define relative atomic mass, Ar
- define relative molecular mass, Mr, and calculate relative molecular mass (and relative formula mass) as the sum of relative atomic masses
- define the term mole in terms of the Avogadro constant
- calculate the percentage mass of an element in a compound when given appropriate information
- calculate empirical and molecular formulae from relevant data
- calculate stoichiometric reacting masses and volumes of gases (one mole of gas occupies 24 dm3 at room temperature and pressure); calculations involving the idea of limiting reactants may be set (knowledge of the gas laws and the calculations of gaseous volumes at different temperatures and pressures are not required)
- apply the concept of solution concentration (in mol/dm3 or g/dm3) to process the results of volumetric experiments (e.g. titration) and to solve simple problems (appropriate guidance will be provided where unfamiliar reactions are involved)
- calculate % yield and % purity.
- describe the meanings of the terms acid and alkali in terms of the ions they produce in aqueous solution and their effects on Universal Indicator
- describe neutrality and relative acidity and alkalinity, in terms of — relative H+ and OH– ion concentrations,
- describe neutrality and relative acidity and alkalinity, in terms of — colour in Universal Indicator, and
- describe neutrality and relative acidity and alkalinity, in terms of — the pH scale (calculation of pH from hydrogen ion concentration is not required)
- describe qualitatively the difference between strong and weak acids in terms of the extent of ionisation
- describe the characteristic properties of acids as in reactions with metals, bases and carbonates to form salts
- describe the reaction between hydrogen ions and hydroxide ions to produce water, H+ + OH– → H2O, as neutralisation
- describe the importance of controlling the pH in soils and how excess acidity can be treated using calcium hydroxide
- describe the characteristic properties of bases in reactions with acids and with ammonium salts
- classify oxides as acidic, basic, amphoteric or neutral based on metallic/non-metallic character.
- describe the techniques used in the preparation, separation and purification of salts as examples of some of the techniques specified in Section 1.2(a) (methods for preparation should include precipitation and titration together with reactions of acids with metals, insoluble bases and insoluble carbonates)
- describe the general rules of solubility for common salts to include nitrates, chlorides (including those of silver and lead), sulfates (including those of barium, calcium and lead), carbonates, hydroxides, salts of Group 1 cations and ammonium salts
- suggest a method of preparing a given salt from suitable starting materials, given appropriate information.
- describe the use of nitrogen, from air, and hydrogen, from the cracking of crude oil, in the manufacture of ammonia
- state that some chemical reactions are reversible, e.g. manufacture of ammonia
- interpret data relating to the conditions used in industry for processes involving reversible reactions, e.g. manufacture of ammonia by the Haber Process (knowledge of Le Chatelier’s Principle is not required).
- describe the use of aqueous sodium hydroxide and/or aqueous ammonia to identify the following aqueous cations through the formation of precipitates (if any) and their subsequent solubility: aluminium, ammonium (together with evolution of ammonia gas upon warming), calcium, copper(II), iron(II), iron(III) and zinc (formulae of complex ions are not required)
- describe tests to identify the following anions: carbonate (by the addition of dilute acid and subsequent use of limewater); chloride (by reaction of an aqueous solution with nitric acid and aqueous silver nitrate); iodide (by reaction of an aqueous solution with nitric acid and aqueous silver nitrate); nitrate (by reduction with aluminium in aqueous sodium hydroxide to ammonia and subsequent use of damp red litmus paper) and sulfate (by reaction of an aqueous solution with nitric acid and aqueous barium nitrate)
- describe tests to identify the following gases: ammonia (using damp red litmus paper); carbon dioxide (using limewater); chlorine (using damp litmus paper); hydrogen (using a burning splint); oxygen (using a glowing splint) and sulfur dioxide (using acidified potassium manganate(VII)).
- define oxidation and reduction (redox) in terms of oxygen/hydrogen gain/loss
- define redox in terms of electron transfer and changes in oxidation state
- identify redox reactions in terms of oxygen/hydrogen gain/loss, electron gain/loss and changes in oxidation state
- describe the use of aqueous potassium iodide and acidified potassium manganate(VII) in testing for oxidising and reducing agents from the resulting colour changes.
- describe electrolysis as the conduction of electricity through an ionic compound (an electrolyte), when molten or dissolved in water, leading to chemical changes (including decomposition) at the electrodes
- describe electrolysis as evidence for the existence of ions which are held in a lattice when solid but which are free to move when molten or in solution
- describe, in terms of the mobility of ions present and the electrode products, the electrolysis of molten sodium chloride, using inert electrodes
- predict the likely products of the electrolysis of a molten binary ionic compound using inert electrodes
- apply the idea of selective discharge based on — cations: linked to the reactivity series (see also 8.4)
- apply the idea of selective discharge based on — anions: halides, hydroxides and sulfates (e.g. aqueous copper(II) sulfate and dilute sodium chloride solution (as essentially the electrolysis of water))
- apply the idea of selective discharge based on — concentration effects (as in the electrolysis of concentrated and dilute aqueous sodium chloride) (in all cases above, inert electrodes are used)
- predict the likely products of the electrolysis of an aqueous electrolyte, given relevant information
- construct ionic equations for the reactions occurring at the electrodes during the electrolysis, given relevant information
- describe the electrolysis of aqueous copper(II) sulfate with copper electrodes as a means of purifying copper (no technical details are required)
- describe the electroplating of metals, e.g. copper plating, and state one use of electroplating
- describe the production of electrical energy from simple cells (i.e. two electrodes in an electrolyte) linked to the reactivity series (see also 8.4) and redox reactions (in terms of electron transfer)
- describe hydrogen, derived from water or hydrocarbons, as a potential fuel, reacting with oxygen to generate electricity directly in a hydrogen fuel cell (details of the construction and operation of a fuel cell are not required).
- describe the Periodic Table as an arrangement of the elements in the order of increasing proton (atomic) number
- describe how the position of an element in the Periodic Table is related to proton number and electronic configuration
- describe the relationship between number of outer (valence) electrons and the ionic charge of an ion for the first twenty elements
- explain the similarities between the elements in the same group of the Periodic Table in terms of their electronic configuration
- describe the change from metallic to non-metallic character from left to right across a period of the Periodic Table
- describe the relationship between number of outer (valence) electrons and metallic/non-metallic character
- predict the properties of elements in Group 1 and Group 17 using the Periodic Table.
- describe lithium, sodium and potassium in Group 1 (the alkali metals) as a collection of relatively soft, low density metals showing a trend in melting point and in their reaction with water
- describe chlorine, bromine and iodine in Group 17 (the halogens) as a collection of diatomic non- metals showing a trend in colour, state and their displacement reactions with solutions of other halide ions
- describe the elements in Group 18 (the noble gases) as a collection of monoatomic elements that are chemically unreactive and hence important in providing an inert environment, e.g. argon and neon in light bulbs; helium in balloons; argon in the manufacture of steel
- describe the lack of reactivity of the noble gases in terms of their electronic configurations.
- describe typical transition elements as metals having high melting point, high density, variable oxidation state and forming coloured compounds
- state that the elements and/or their compounds are often able to act as catalysts (see also 10(d)).
- place in order of reactivity calcium, copper, (hydrogen), iron, lead, magnesium, potassium, silver, sodium and zinc by reference to the reactions, if any, of the metals with water, steam and dilute hydrochloric acid
- place in order of reactivity calcium, copper, (hydrogen), iron, lead, magnesium, potassium, silver, sodium and zinc by reference to — the reduction, if any, of their oxides by carbon and/or by hydrogen
- describe the reactivity series as related to the tendency of a metal to form its positive ion, illustrated by its reaction with — the aqueous ions of the other listed metals
- describe the reactivity series as related to the tendency of a metal to form its positive ion, illustrated by its reaction with — the oxides of the other listed metals
- deduce the order of reactivity from a given set of experimental results
- describe the action of heat on the carbonates of the listed metals and relate thermal stability to the reactivity series
- describe the ease of obtaining metals from their ores by relating the elements to their positions in the reactivity series
- describe the essential conditions for the corrosion (rusting) of iron as the presence of oxygen and water; prevention of rusting can be achieved by placing a barrier around the metal, e.g. painting; greasing; plastic coating; galvanising
- describe the sacrificial protection of iron by a more reactive metal in terms of the reactivity series where the more reactive metal corrodes preferentially, e.g. underwater pipes have a piece of magnesium attached to them.
- describe the meaning of enthalpy change in terms of exothermic (∆H negative) and endothermic (∆H positive) reactions
- represent energy changes by energy profile diagrams, including reaction enthalpy changes and activation energies (see also 10(c), 10(d))
- describe bond breaking as an endothermic process and bond making as an exothermic process
- explain qualitatively overall enthalpy changes in terms of the energy changes associated with the breaking and making of covalent bonds.
- describe the effect of concentration, pressure, particle size and temperature on the rates of reactions and explain these effects in terms of collisions between reacting particles
- define the term catalyst and describe the effect of catalysts (including enzymes) on the rates of reactions
- explain how pathways with lower activation energies account for the increase in rates of reactions (see also 9(b))
- state that some compounds act as catalysts in a range of industrial processes and that enzymes are biological catalysts (see also 8.3(b), 9(b), 10(c) and 12(d))
- suggest a suitable method for investigating the effect of a given variable on the rate of a reaction
- interpret data obtained from experiments concerned with rate of reaction.
- name natural gas, mainly methane, and crude oil as non-renewable sources of energy
- describe crude oil as a mixture of hydrocarbons and its separation by fractional distillation to yield fractions which have competing uses as fuels and as a source of chemicals (see also 1.2(a))
- describe biofuel (exemplified by bioethanol from sugarcane) as a renewable alternative to natural gas and crude oil
- describe how biofuel, when compared to fossil fuels, can be more environmentally sustainable in terms of carbon dioxide emission (see also 12(g)).
- describe a homologous series as a group of compounds with a general formula, similar chemical properties and showing a gradation in physical properties as a result of increase in the size and mass of the molecules, e.g. melting and boiling points; viscosity
- describe the alkanes as a homologous series of saturated hydrocarbons with the general formula CnH2n+2
- draw the structures of branched and unbranched alkanes, C1 to C4, and name the unbranched alkanes methane to butane
- define isomerism and identify isomers
- describe alkanes (exemplified by methane) as being generally unreactive except in terms of combustion and substitution by chlorine
- describe the alkenes as a homologous series of unsaturated hydrocarbons with the general formula CnH2n
- draw the structures of branched and unbranched alkenes, C2 to C4, and name the unbranched alkenes ethene to butene
- describe the manufacture of alkenes and hydrogen by cracking hydrocarbons and recognise that cracking is essential to match the demand for fractions containing smaller molecules from the refinery process
- describe the difference between saturated and unsaturated hydrocarbons from their molecular structures and by using aqueous bromine
- describe the reactions of alkenes (exemplified by ethene) in terms of combustion, polymerisation (see also 11.4(b)), and the addition with bromine, steam and hydrogen
- state the meaning of polyunsaturated when applied to food products
- describe the manufacture of margarine by the addition of hydrogen to unsaturated vegetable oils to form a solid product.
- describe the alcohols as a homologous series containing the –OH group
- draw the structures of branched and unbranched alcohols, C1 to C4, and name the unbranched alcohols methanol to butanol
- describe the reactions of alcohols in terms of combustion and oxidation to carboxylic acids
- describe the formation of ethanol by the catalysed addition of steam to ethene and by fermentation of glucose
- describe the carboxylic acids as a homologous series containing the –CO2H group
- draw the structures of carboxylic acids, C1 to C4, and name the unbranched acids methanoic acid to butanoic acid
- describe the carboxylic acids as weak acids, reacting with carbonates, bases and some metals
- describe the formation of ethanoic acid by the oxidation of ethanol by atmospheric oxygen or acidified potassium manganate(VII).
- describe the reaction of a carboxylic acid with an alcohol to form an ester, e.g. ethyl ethanoate
- deduce the name and formula of an ester from the unbranched carboxylic acid, C1 to C4, and alcohol, C1 to C4, and vice versa.
- describe polymers as large molecules built up from small units (monomers), different polymers having different units and/or different linkages
- describe the formation of poly(ethene) as an example of addition polymerisation of ethene as the monomer (see also 11.2(j))
- state some uses of poly(ethene) as a typical plastic, e.g. plastic bags; clingfilm
- deduce the structure of the polymer product from a given monomer and vice versa
- describe nylon, a polyamide, and Terylene, a polyester, as condensation polymers, the partial structure of nylon being represented as and the partial structure of Terylene as (details of manufacture and mechanisms of these polymerisations are not required)
- state some typical uses of man-made fibres such as nylon and Terylene, e.g. clothing; curtain materials; fishing line; parachutes; sleeping bags
- describe the pollution problems caused by the disposal of non-biodegradable plastics
- describe two methods of recycling plastics as
- physical method (exemplified by melting small pieces of poly(ethene) waste into pellets)
- chemical method (exemplified by depolymerisation and cracking of plastic waste into chemical feedstock and fuel respectively) — describe depolymerisation as a process in which polymers are broken down into their monomers, exemplified by hydrolysis of polyesters using acid as a catalyst (details of mechanisms are not required)
- discuss the social, economic and environmental issues of recycling plastics.
- describe the volume composition of gases present in dry air as being approximately 78% nitrogen, 21% oxygen and the remainder being noble gases (with argon as the main constituent) and carbon dioxide
- name some common atmospheric pollutants, e.g. carbon monoxide; methane; nitrogen oxides (NO and NO2); ozone; sulfur dioxide; unburned hydrocarbons
- state the sources of these pollutants as — carbon monoxide from incomplete combustion of carbon-containing substances
- state the sources of these pollutants as — nitrogen oxides from lightning activity and internal combustion engines
- state the sources of these pollutants as — sulfur dioxide from volcanoes and combustion of fossil fuels
- describe the reactions used in possible solutions to the problems arising from some of the pollutants named in (b) — the redox reactions in catalytic converters to remove combustion pollutants (see also 10(d))
- describe the reactions used in possible solutions to the problems arising from some of the pollutants named in (b) — the use of calcium carbonate to reduce the effect of ‘acid rain’ and in flue gas desulfurisation
- discuss some of the effects of these pollutants on health and on the environment — the toxic nature of carbon monoxide
- discuss some of the effects of these pollutants on health and on the environment — the role of nitrogen dioxide and sulfur dioxide in the formation of ‘acid rain’ and its effects on respiration and buildings
- discuss the importance of the ozone layer and the problems involved with the depletion of ozone by reaction with chlorine-containing compounds, chlorofluorocarbons (CFCs)
- describe the carbon cycle in simple terms, to include — the processes of combustion, respiration and photosynthesis
- describe the carbon cycle in simple terms, to include — how the carbon cycle regulates the amount of carbon dioxide in the atmosphere (see also 11.1(d))
- state that carbon dioxide and methane are greenhouse gases and may contribute to global warming; give the sources of these gases and describe the potential effects of increased levels of these greenhouse gases, including more extreme weather events and melting of polar ice.
G3 Pure Chemistry and O-Level Chemistry share one set of topics and lessons. Work through topics 1–12 in order, or jump to the topic you are revising.
G3 Pure or O-Level?
The lessons are the same. Where syllabus years, papers or exam formats differ, follow the guidance labelled for your qualification. Compare the G1, G2 and G3 courses.
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Learning the course? Start with topic 1, Experimental Chemistry, and work through in order. Revising? Open your weakest topic, then try the mixed quiz, the topic quizzes or structured questions.
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O Level structured practice currently covers 12/12 hubs.
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K324 paper-skills orientation
| Paper | What it tests (practically) |
|---|---|
| Paper 1 | Fast recall + quick application (MCQ). |
| Paper 2 | Structured questions: explanations, calculations, and data/QA. |
| Paper 3 | Practical skills: planning; manipulation, measurement and observation (MMO); presenting data and observations (PDO); analysis, conclusions and evaluation (ACE). |
Shared course topic map
1. Experimental Chemistry
Measurement, apparatus, planning, data, separation, chromatography and purity.
2–3. Matter, Bonding and Structure
Particle theory, atomic structure, ions and ionic, covalent and metallic bonding.
4. Chemical Calculations
Formulae, equations, moles, reacting quantities, concentration, yield and purity.
5. Acid-Base Chemistry
Acids, bases, salts, ammonia and reversible industrial manufacture.
6. Qualitative Analysis
Tests for cations, anions and gases using observation → inference reasoning.
7. Redox Chemistry
Oxidation and reduction, oxidation states, redox tests, electrolysis, electroplating, simple cells and fuel cells.
8. Patterns in the Periodic Table
Periodic patterns, group chemistry, transition elements, extraction and corrosion.
9. Chemical Energetics
Exothermic and endothermic changes, energy profiles and qualitative bond-energy reasoning.
10. Rate of Reactions
Rate measurement, collision theory, catalysts, investigations and data interpretation.
11. Organic Chemistry
Fuels, hydrocarbons, alcohols, acids, esters and polymers.
12. Maintaining Air Quality
Air composition, pollutants and treatment, ozone depletion, the carbon cycle and climate.
Practical skills support
K324 / 6092 practical skills hub
Planning, measurements, tables and graphs, titration, evaluation and qualitative analysis.
Mock Exams
K324 / 6092 Mock Paper 1 (MCQ)
Timed Paper 1 practice across the shared course. Select Version A, B or C.
Exam Skills
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