1.1

understand the three states of matter in terms of the arrangement, movement and

energy of the particles

1.2

understand the interconversions between the three states of matter in terms of:

• the names of the interconversions

• how they are achieved

• the changes in arrangement, movement and energy of the particles.

1.3

understand how the results of experiments involving the dilution of coloured solutions

and diffusion of gases can be explained

1.4

know what is meant by the terms:

• solvent

• solute

• solution

• saturated solution.

1.8

understand how to classify a substance as an element, compound or mixture

1.9

understand that a pure substance has a fixed melting and boiling point, but that a

mixture may melt or boil over a range of temperatures

1.10

describe these experimental techniques for the separation of mixtures:

• simple distillation

• fractional distillation

• filtration

• crystallisation

• paper chromatography.

1.11

understand how a chromatogram provides information about the composition of a

mixture

1.12

understand how to use the calculation of R_f values to identify the components of a

mixture

1.13

practical: investigate paper chromatography using inks/food colourings

1.14

know what is meant by the terms atom and molecule

1.15

know the structure of an atom in terms of the positions, relative masses and relative

charges of sub-atomic particles

1.16

know what is meant by the terms atomic number, mass number, isotopes and relative

atomic mass (A_r)

1.17

be able to calculate the relative atomic mass of an element (A_r) from isotopic

abundances

1.18

understand how elements are arranged in the Periodic Table:

• in order of atomic number

• in groups and periods.

1.19

understand how to deduce the electronic configurations of the first 20 elements from

their positions in the Periodic Table

1.20

understand how to use electrical conductivity and the acid-base character of oxides to

classify elements as metals or non-metals

1.21

identify an element as a metal or a non-metal according to its position in the Periodic

Table

1.22

understand how the electronic configuration of a main group element is related to its

position in the Periodic Table

1.23

understand why elements in the same group of the Periodic Table have similar

chemical properties

1.24

understand why the noble gases (Group 0) do not readily react

1.25

write word equations and balanced chemical equations (including state symbols):

• for reactions studied in this specification

• for unfamiliar reactions where suitable information is provided

1.26

calculate relative formula masses (including relative molecular masses) (M_r) from

relative atomic masses (A_r)

1.27

know that the mole (mol) is the unit for the amount of a substance

1.28

understand how to carry out calculations involving amount of substance, relative

atomic mass (A_r) and relative formula mass (M_r)

1.29

calculate reacting masses using experimental data and chemical equations

1.30

calculate percentage yield

1.31

understand how the formulae of simple compounds can be obtained experimentally,

including metal oxides, water and salts containing water of crystallisation

1.32

know what is meant by the terms empirical formula and molecular formula

1.33

calculate empirical and molecular formulae from experimental data

1.36

practical: know how to determine the formula of a metal oxide by combustion

(e.g. magnesium oxide) or by reduction (e.g. copper(II) oxide)

1.37

understand how ions are formed by electron loss or gain

1.38 know the charges of these ions:

• metals in Groups 1, 2 and 3

• non-metals in Groups 5, 6 and 7

• Ag⁺, Cu²⁺, Fe²⁺, Fe³⁺, Pb²⁺, Zn²⁺

• hydrogen (H⁺), hydroxide (OH^–), ammonium (NH₄⁺), carbonate (CO₃^2–), nitrate

(NO₃^-), sulfate (SO₄^2–).

1.39

write formulae for compounds formed between the ions listed above

1.40

draw dot-and-cross diagrams to show the formation of ionic compounds by electron

transfer, limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7

only outer electrons need be shown

1.41

understand ionic bonding in terms of electrostatic attractions

1.42

understand why compounds with giant ionic lattices have high melting and boiling points

1.43

know that ionic compounds do not conduct electricity when solid, but do conduct

electricity when molten and in aqueous solution

1.44

know that a covalent bond is formed between atoms by the sharing of a pair of

electrons

1.45

understand covalent bonds in terms of electrostatic attractions

1.46

understand how to use dot-and-cross diagrams to represent covalent bonds in:

• diatomic molecules, including hydrogen, oxygen, nitrogen, halogens and hydrogen

halides

• inorganic molecules including water, ammonia and carbon dioxide

• organic molecules containing up to two carbon atoms, including methane, ethane,

ethene and those containing halogen atoms.

1.47

explain why substances with a simple molecular structures are gases or liquids, or

solids with low melting and boiling points

the term intermolecular forces of attraction can be used to represent all forces

between molecules

1.48

explain why the melting and boiling points of substances with simple molecular

structures increase, in general, with increasing relative molecular mass

1.49

explain why substances with giant covalent structures are solids with high melting and

boiling points

1.50

explain how the structures of diamond, graphite and C₆₀ fullerene influence their

physical properties, including electrical conductivity and hardness

1.51

know that covalent compounds do not usually conduct electricity

2.1

understand how the similarities in the reactions of these elements with water provide

evidence for their recognition as a family of elements

2.2

understand how the differences between the reactions of these elements with air and

water provide evidence for the trend in reactivity in Group 1

2.3

use knowledge of trends in Group 1 to predict the properties of other alkali metals

2.5

know the colours, physical states (at room temperature) and trends in physical

properties of these elements

2.6

use knowledge of trends in Group 7 to predict the properties of other halogens

2.7

understand how displacement reactions involving halogens and halides provide

evidence for the trend in reactivity in Group 7

2.9

know the approximate percentages by volume of the four most abundant gases in dry

air

2.10

understand how to determine the percentage by volume of oxygen in air using

experiments involving the reactions of metals (e.g. iron) and non-metals

(e.g. phosphorus) with air

2.11

describe the combustion of elements in oxygen, including magnesium, hydrogen and

sulfur

2.12

describe the formation of carbon dioxide from the thermal decomposition of metal

carbonates, including copper(II) carbonate

2.13

know that carbon dioxide is a greenhouse gas and that increasing amounts in the

atmosphere may contribute to climate change

2.14

practical: determine the approximate percentage by volume of oxygen in air using a

metal or a non-metal

2.15

understand how metals can be arranged in a reactivity series based on their reactions

with:

• water

• dilute hydrochloric or sulfuric acid.

2.16

understand how metals can be arranged in a reactivity series based on their

displacement reactions between:

• metals and metal oxides

• metals and aqueous solutions of metal salts.

2.17

know the order of reactivity of these metals: potassium, sodium, lithium, calcium,

magnesium, aluminium, zinc, iron, copper, silver, gold

2.18

know the conditions under which iron rusts

2.19

understand how the rusting of iron may be prevented by:

• barrier methods

• galvanising

• sacrificial protection.

2.20

the terms:

• oxidation

• reduction

• redox

• oxidising agent

• reducing agent

in terms of gain or loss of oxygen and loss or gain of electrons.

2.21

practical: investigate reactions between dilute hydrochloric and sulfuric acids and

metals (e.g. magnesium, zinc and iron)

2.28

describe the use of litmus, phenolphthalein and methyl orange to distinguish between

acidic and alkaline solutions

2.29

understand how to use the pH scale, from 0–14, can be used to classify solutions as

strongly acidic (0–3), weakly acidic (4–6), neutral (7), weakly alkaline (8–10) and

strongly alkaline (11–14)

2.30

describe the use of universal indicator to measure the approximate pH value of an

aqueous solution

2.31

know that acids in aqueous solution are a source of hydrogen ions and alkalis in a

aqueous solution are a source of hydroxide ions

2.32

know that alkalis can neutralise acids

2.34

know the general rules for predicting the solubility of ionic compounds in water:

• common sodium, potassium and ammonium compounds are soluble

• all nitrates are soluble

• common chlorides are soluble, except those of silver and lead(II)

• common sulfates are soluble, except for those of barium, calcium and lead(II)

• common carbonates are insoluble, except for those of sodium, potassium and

ammonium

• common hydroxides are insoluble except for those of sodium, potassium and

calcium (calcium hydroxide is slightly soluble).

2.35

understand acids and bases in terms of proton transfer

2.36

understand that an acid is a proton donor and a base is a proton acceptor

2.37

describe the reactions of hydrochloric acid, sulfuric acid and nitric acid with metals,

bases and metal carbonates (excluding the reactions between nitric acid and metals)

to form salts

2.38

know that metal oxides, metal hydroxides and ammonia can act as bases, and that

alkalis are bases that are soluble in water

2.39

describe an experiment to prepare a pure, dry sample of a soluble salt, starting from

an insoluble reactant

2.42

practical: prepare a sample of pure, dry hydrated copper(II) sulfate crystals starting

from copper(II) oxide

2.44

describe tests for these gases:

• hydrogen

• oxygen

• carbon dioxide

• ammonia

• chlorine.

2.45

describe how to carry out a flame test

2.46

know the colours formed in flame tests for these cations:

• Li⁺ is red

• Na⁺ is yellow

• K⁺ is lilac

• Ca²⁺ is orange-red

• Cu²⁺ is blue-green.

2.47

describe tests for these cations:

• NH⁴⁺ using sodium hydroxide solution and identifying the gas evolved

• Cu²⁺, Fe²⁺ and Fe³⁺ using sodium hydroxide solution.

2.48

describe tests for these anions:

• Cl^–, Br^– and I^– using acidified silver nitrate solution

• SO₄^2– using acidified barium chloride solution

• CO₃^2– using hydrochloric acid and identifying the gas evolved.

2.49

describe a test for the presence of water using anhydrous copper(II) sulfate

2.50

describe a physical test to show whether a sample of water is pure

3.1

know that chemical reactions in which heat energy is given out are described as

exothermic, and those in which heat energy is taken in are described as endothermic

3.2

describe simple calorimetry experiments for reactions such as combustion,

displacement, dissolving and neutralisation

3.3

calculate the heat energy change from a measured temperature change using the

expression Q = mcΔT

3.4

calculate the molar enthalpy change (ΔH) from the heat energy change, Q

3.8

practical: investigate temperature changes accompanying some of the following types

of change:

• salts dissolving in water

• neutralisation reactions

• displacement reactions

• combustion reactions.

3.9

describe experiments to investigate the effects of changes in surface area of a solid,

concentration of a solution, temperature and the use of a catalyst on the rate of a

reaction

3.10

describe the effects of changes in surface area of a solid, concentration of a solution,

pressure of a gas, temperature and the use of a catalyst on the rate of a reaction

3.11

explain the effects of changes in surface area of a solid, concentration of a solution,

pressure of a gas and temperature on the rate of a reaction in terms of particle

collision theory

3.12

know that a catalyst is a substance that increases the rate of a reaction, but is

chemically unchanged at the end of the reaction

3.13

know that a catalyst works by providing an alternative pathway with lower activation

energy

3.15

practical: investigate the effect of changing the surface area of marble chips and of

changing the concentration of hydrochloric acid on the rate of reaction between

marble chips and dilute hydrochloric acid

3.16

practical: investigate the effect of different solids on the catalytic decomposition of

hydrogen peroxide solution

3.17

know that some reactions are reversible and this is indicated by the symbol ⇌ in

equations

3.18

describe reversible reactions such as the dehydration of hydrated copper(II) sulfate

and the effect of heat on ammonium chloride

4.1

know that a hydrocarbon is a compound of hydrogen and carbon only

4.2

understand how to represent organic molecules using empirical formulae, molecular

formulae, general formulae, structural formulae and displayed formulae

4.3

know what is meant by the terms homologous series, functional group and isomerism

4.4

understand how to name compounds relevant to this specification using the rules of

International Union of Pure and Applied Chemistry (IUPAC) nomenclature

students will be expected to name compounds containing up to six carbon atoms

4.5

understand how to write the possible structural and displayed formulae of an organic

molecule given its molecular formula

4.6

understand how to classify reactions of organic compounds as substitution, addition

and combustion

knowledge of reaction mechanisms is not required

4.7

know that crude oil is a mixture of hydrocarbons

4.8

describe how the industrial process of fractional distillation separates crude oil into

fractions

4.9

know the names and uses of the main fractions obtained from crude oil:

refinery gases, gasoline, kerosene, diesel, fuel oil and bitumen

4.10

know the trend in colour, boiling point and viscosity of the main fractions

4.11

know that a fuel is a substance that, when burned, releases heat energy

4.12

know the possible products of complete and incomplete combustion of hydrocarbons

with oxygen in the air

4.13

understand why carbon monoxide is poisonous, in terms of its effect on the capacity

of blood to transport oxygen

references to haemoglobin are not required

4.14

know that, in car engines, the temperature reached is high enough to allow nitrogen

and oxygen from air to react, forming oxides of nitrogen

4.15

explain how the combustion of some impurities in hydrocarbon fuels results in the

formation of sulfur dioxide

4.16

understand how sulfur dioxide and oxides of nitrogen contribute to acid rain

4.17

describe how long-chain alkanes are converted to alkenes and shorter-chain alkanes

by catalytic cracking (using silica or alumina as the catalyst and a temperature in the

range of 600–700 ºC)

4.18

explain why cracking is necessary, in terms of the balance between supply and

demand for different fractions

4.19

know the general formula for alkanes

4.20

explain why alkanes are classified as saturated hydrocarbons

4.21

understand how to draw the structural and displayed formulae for alkanes with up to

five carbon atoms in the molecule, and to name the unbranched-chain isomers

4.22

describe the reactions of alkanes with halogens in the presence of ultraviolet

radiation, limited to mono-substitution

knowledge of reaction mechanisms is not required

4.23

know that alkenes contain the functional group >C=C<

4.24

know the general formula for alkenes

4.25

explain why alkenes are classified as unsaturated hydrocarbons

4.26

understand how to draw the structural and displayed formulae for alkenes with up to

four carbon atoms in the molecule, and name the unbranched-chain isomers

knowledge of cis/trans or E/Z notation is not required

4.27

describe the reactions of alkenes with bromine to produce dibromoalkanes

4.28

describe how bromine water can be used to distinguish between an alkane and an

alkene

4.44

know that an addition polymer is formed by joining up many small molecules called

monomers

4.45

understand how to draw the repeat unit of an addition polymer, including

poly(ethene), poly(propene), poly(chloroethene) and (poly)tetrafluoroethene

4.46

understand how to deduce the structure of a monomer from the repeat unit of an

addition polymer and vice versa

4.47

explain problems in the disposal of addition polymers, including:

• their inertness and inability to biodegrade

• the production of toxic gases when they are burned.