This topic illustrates how quantitative relationships can be established when different substances react. (The term relative formula mass or Mr will be used for all compounds including ionic compounds.)

1.1 Relative masses of atoms and molecules

a) define and use the terms relative atomic, isotopic, molecular and formula masses, based on the 12C scale

1.2 The mole and the Avogadro constant

a) define and use the term mole in terms of the Avogadro constant

1.3 The determination of relative atomic masses, Ar

a) analyse mass spectra in terms of isotopic abundances (knowledge of the working of the mass spectrometer is not required)

b) calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum

1.4 The calculation of empirical and molecular formulae

a) define and use the terms empirical and molecular formula

b) calculate empirical and molecular formulae, using combustion data or composition by mass

1.5 Reacting masses and volumes (of solutions and gases)

a) write and construct balanced equations

b) perform calculations, including use of the mole concept, involving:

(i) reacting masses (from formulae and equations)

(ii) volumes of gases (e.g. in the burning of hydrocarbons)

(iii) volumes and concentrations of solutions When performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question. When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified.

c) deduce stoichiometric relationships from calculations such as those in 1.5(b)

This topic describes the type, number and distribution of the fundamental particles which make up an atom and the impact of this on some atomic properties.

2.1 Particles in the atom

a) identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses

b) deduce the behaviour of beams of protons, neutrons and electrons in electric fields

c) describe the distribution of mass and charge within an atom

d) deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (atomic and mass numbers) and charge

2.2 The nucleus of the atom

a) describe the contribution of protons and neutrons to atomic nuclei in terms of proton (atomic) number and nucleon (mass) number

b) distinguish between isotopes on the basis of different numbers of neutrons present

c) recognise and use the symbolism x yA for isotopes, where x is the nucleon (mass) number and y is the proton (atomic) number

2.3 Electrons: energy levels, atomic orbitals, ionisation energy, electron affinity

a) describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals

b) describe and sketch the shapes of s and p orbitals

c) state the electronic configuration of atoms and ions given the proton (atomic) number and charge, using the convention 1s2 2s2 2p6 , etc.

d) (i) explain and use the term ionisation energy

(ii) explain the factors influencing the ionisation energies of elements

(iii) explain the trends in ionisation energies across a period and down a group of the Periodic Table.

e) deduce the electronic configurations of elements from successive ionisation energy data

f) interpret successive ionisation energy data of an element in terms of the position of that element within the Periodic Table

This topic introduces the different ways by which chemical bonding occurs and the effect this can have on physical properties.

3.1 Ionic bonding

a) describe ionic bonding, using the examples of sodium chloride, magnesium oxide and calcium fluoride, including the use of ‘dot-andcross’ diagrams

3.2 Covalent bonding and co-ordinate (dative covalent) bonding including shapes of simple molecules

a) describe, including the use of ‘dot-and-cross’ diagrams:

(i) covalent bonding, in molecules such as hydrogen, oxygen, chlorine, hydrogen chloride, carbon dioxide, methane, ethene

(ii) co-ordinate (dative covalent) bonding, such as in the formation of the ammonium ion and in the Al 2Cl 6 molecule

b) describe covalent bonding in terms of orbital overlap, giving σ and π bonds, including the concept of hybridisation to form sp, sp2 and sp3 orbitals

c) explain the shapes of, and bond angles in, molecules by using the qualitative model of electron-pair repulsion (including lone pairs), using as simple examples BF3 (trigonal planar), CO2 (linear), CH4 (tetrahedral), NH3 (pyramidal), H2O (non-linear), SF6 (octahedral), PF5 (trigonal bipyramidal)

d) predict the shapes of, and bond angles in, molecules and ions

3.3 Intermolecular forces, electronegativity and bond properties

a) describe hydrogen bonding, using ammonia and water as simple examples of molecules containing N–H and O–H groups

b) understand, in simple terms, the concept of electronegativity and apply it to explain the properties of molecules such as bond polarity, the dipole moments of moleculesand the behaviour of oxides with water

c) explain the terms bond energy, bond length and bond polarity and use them to compare the reactivities of covalent bonds

d) describe intermolecular forces (van der Waals’ forces), based on permanent and induced dipoles, as in, for example, CHCl 3(l); Br2(l) and the liquid Group 18 elements

3.4 Metallic bonding

a) describe metallic bonding in terms of a lattice of positive ions surrounded by delocalised electrons

3.5 Bonding and physical properties

a) describe, interpret and predict the effect of different types of bonding (ionic bonding, covalent bonding, hydrogen bonding, other intermolecular interactions, metallic bonding) on the physical properties of substances

b) deduce the type of bonding present from given information

c) show understanding of chemical reactions in terms of energy transfers associated with the breaking and making of chemical bonds

The study of the particles in solids, liquids and gases and the interactions between them is important in understanding the physical properties of substances.

4.1 The gaseous state: ideal and real gases and pV = nRT

a) state the basic assumptions of the kinetic theory as applied to an ideal gas

b) explain qualitatively in terms of intermolecular forces and molecular size:

(i) the conditions necessary for a gas to approach ideal behaviour

(ii) the limitations of ideality at very high pressures and very low temperatures c) state and use the general gas equation pV = nRT in calculations, including the determination of Mr

4.2 The liquid state

a) describe, using a kinetic-molecular model, the liquid state, melting, vaporisation and vapour pressure

4.3 The solid state: lattice structures

a) describe, in simple terms, the lattice structure of a crystalline solid which is:

(i) ionic, as in sodium chloride and magnesium oxide

(ii) simple molecular, as in iodine and the fullerene allotropes of carbon (C60 and nanotubes only)

(iii) giant molecular, as in silicon(IV) oxide and the graphite, diamond and graphene allotropes of carbon

(iv) hydrogen-bonded, as in ice

(v) metallic, as in copper b) discuss the finite nature of materials as a resource and the importance of recycling processes

c) outline the importance of hydrogen bonding to the physical properties of substances, including ice and water (for example, boiling and melting points, viscosity and surface tension)

d) suggest from quoted physical data the type of structure and bonding present in a substance

This topic illustrates the relationship between electricity and chemical changes. Chemical reactions can be investigated by looking at electrode potentials.

6.1 Redox processes: electron transfer and changes in oxidation number (oxidation state)

a) calculate oxidation numbers of elements in compounds and ions

b) describe and explain redox processes in terms of electron transfer and changes in oxidation number

c) use changes in oxidation numbers to help balance chemical equations

This topic illustrates that many chemical reactions are reversible and involve an equilibrium process. The consideration of the many factors that can affect an equilibrium is an important aspect of physical chemistry.

7.1 Chemical equilibria: reversible reactions, dynamic equilibrium

a) explain, in terms of rates of the forward and reverse reactions, what is meant by a reversible reaction and dynamic equilibrium

b) state Le Chatelier’s principle and apply it to deduce qualitatively (from appropriate information) the effects of changes in temperature, concentration or pressure on a system at equilibrium

c) state whether changes in temperature, concentration or pressure or the presence of a catalyst affect the value of the equilibrium constant for a reaction

d) deduce expressions for equilibrium constants in terms of concentrations, Kc, and partial pressures, Kp (treatment of the relationship between Kp and Kc is not required)

e) calculate the values of equilibrium constants in terms of concentrations or partial pressures from appropriate data

f) calculate the quantities present at equilibrium, given appropriate data (such calculations will not require the solving of quadratic equations)

g) describe and explain the conditions used in the Haber process and the Contact process, as examples of the importance of an understanding of chemical equilibrium in the chemical industry

7.2 Brønsted–Lowry theory of acids and bases

1. state the names and formulae of the common acids, limited to hydrochloric acid, HCl, sulfuric acid, H2SO4, nitric acid, HNO3, ethanoic acid, CH3COOH

2. state the names and formulae of the common alkalis, limited to sodium hydroxide, NaOH, potassium hydroxide, KOH, ammonia, NH3.

3. describe the Brønsted–Lowry theory of acids and bases

4. describe strong acids and strong bases as fully dissociated in aqueous solution and weak acids and weak bases as partially dissociated in aqueous solution

5. appreciate that water has pH of 7, acid solutions pH of below 7 and alkaline solutions pH of above 7

6. explain qualitatively the differences in behaviour between strong and weak acids including the reaction with a reactive metal and difference in pH values by use of a pH meter, universal indicator or conductivity

7. understand that neutralisation reactions occur when H+ (aq) and OH– (aq) form H2O(l)

8. understand that salts are formed in neutralisation reactions

9. sketch the pH titration curves of titrations using combinations of strong and weak acids with strong and weak alkalis

10. select suitable indicators for acid-alkali titrations, given appropriate data (pKa values will not be used)

The investigation of the factors that affect the rate of a chemical reaction is important in the study of physical chemistry. The temperature and the addition of a catalyst can both affect the progression of a chemical reaction.

8.1 Simple rate equations.

a) explain and use the term rate of reaction

b) explain qualitatively, in terms of collisions, the effect of concentration changes on the rate of a reaction

8.2 Effect of temperature: on reaction rates and rate constants and the concept of activation energy

a) explain and use the term activation energy, including reference to the Boltzmann distribution

b) explain qualitatively, in terms both of the Boltzmann distribution and of collision frequency, the effect of temperature change on the rate of a reaction

8.3 Homogeneous and heterogeneous catalysts including enzymes

a) explain and use the term catalysis

b) explain that catalysts can be homogeneous or heterogeneous

c) (i) explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower activation energy

(ii) interpret this catalytic effect in terms of the Boltzmann distribution

d) describe enzymes as biological catalysts (proteins) which may have specificity

This topic illustrates the regular patterns in chemical and physical properties of the elements in the Periodic Table.

9.1 Periodicity of physical properties of the elements in Period 3

a) describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet)

b) explain qualitatively the variation in atomic radius and ionic radius

c) interpret the variation in melting point and electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements d) explain the variation in first ionisation energy (see the Data Booklet) e) explain the strength, high melting point and electrical insulating properties of ceramics in terms of their giant structure; to include magnesium oxide, aluminium oxide and silicon dioxide

Organic chemistry involves the study of a large class of chemical compounds containing carbon. This topic introduces naming conventions, organic reaction terminology and structures of organic molecules.

14.1 Formulae, functional groups and the naming of organic compounds

a) interpret and use the general, structural, displayed and skeletal formulae of the following classes of compound:

(i) alkanes and alkenes

(ii) halogenoalkanes

(iii) alcohols (including primary, secondary and tertiary)

(iv) aldehydes and ketones

(v) carboxylic acids and esters

(vi) amines (primary only) and nitriles (Candidates are expected to recognise the shape of the benzene ring when it is present in organic compounds. Knowledge of benzene or its compounds is not required for AS Level.)

b) understand and use systematic nomenclature of simple aliphatic organic molecules with functional groups detailed in 14.1(a), up to six carbon atoms (six plus six for esters and amides, straight chains only)

c) deduce the possible isomers for an organic molecule of known molecular formula

d) deduce the molecular formula of a compound, given its structural, displayed or skeletal formula

14.2 Characteristic organic reactions

a) interpret and use the following terminology associated with types of organic reactions:

(i) functional group

(ii) homolytic and heterolytic fission

(iii) free radical, initiation, propagation, termination

(iv) nucleophile, electrophile (v) addition, substitution, elimination, hydrolysis, condensation

(vi) oxidation and reduction

(in equations for organic redox reactions, the symbols [O] and [H] are acceptable for oxidising and reducing agents)

14.3 Shapes of organic molecules; σ and π bonds

a) (i) describe and explain the shape of, and bond angles in, the ethane and ethene molecules in terms of σ and π bonds

(ii) predict the shapes of, and bond angles in, other related molecules .

14.4 Isomerism: structural and stereoisomerism

a) describe structural isomerism and its division into chain, positional and functional group isomerism

b) describe stereoisomerism and its division into geometrical (cis-trans) and optical isomerism (use of E, Z nomenclature is acceptable but is not required)

c) describe geometrical (cis-trans) isomerism in alkenes, and explain its origin in terms of restricted rotation due to the presence of π bonds

d) explain what is meant by a chiral centre and that such a centre normally gives rise to optical isomerism (Candidates should appreciate that compounds can contain more than one chiral centre, but knowledge of meso compounds, or nomenclature such as diastereoisomers is not required.)

e) identify chiral centres and geometrical (cis-trans) isomerism in a molecule of given structural formula

Compounds containing only carbon and hydrogen are called hydrocarbons. This class of compound can be subdivided into alkanes, alkenes and arenes.

15.1 Alkanes

a) understand the general unreactivity of alkanes, including towards polar reagents

b) describe the chemistry of alkanes as exemplified by the following reactions of ethane:

(i) combustion

(ii) substitution by chlorine and by bromine

c) describe the mechanism of free-radical substitution at methyl groups with particular reference to the initiation, propagation and termination reactions

d) explain the use of crude oil as a source of both aliphatic and aromatic hydrocarbons

e) suggest how cracking can be used to obtain more useful alkanes and alkenes of lower Mr from larger hydrocarbon molecules

15.2 Alkenes

a) describe the chemistry of alkenes as exemplified, where relevant, by the following reactions of ethene and propene (including the Markovnikov addition of asymmetric electrophiles to alkenes using propene as an example):

(i) addition of hydrogen, steam, hydrogen halides and halogens

(ii) oxidation by cold, dilute, acidified manganate(VII) ions to form the diol

(iii) oxidation by hot, concentrated, acidified manganate(VII) ions leading to the rupture of the carbon–carbon double bond in order to determine the position of alkene linkages in larger molecules

(iv) polymerisation

b) describe the mechanism of electrophilic addition in alkenes, including using bromine/ethene and hydrogen bromide/propene as examples

c) describe and explain the inductive effects of alkyl groups on the stability of cations formed during electrophilic addition

d) describe the characteristics of addition polymerisation as exemplified by poly(ethene) and PVC

e) deduce the repeat unit of an addition polymer obtained from a given monomer

f) identify the monomer(s) present in a given section of an addition polymer molecule

g) recognise the difficulty of the disposal of poly(alkene)s, i.e. nonbiodegradability and harmful combustion products.

15.3 Hydrocarbons as fuels

a) describe and explain how the combustion reactions of alkanes make them suitable to be used as fuels in industry, in the home and in transport

b) recognise the environmental consequences of:

(i) carbon monoxide, oxides of nitrogen and unburnt hydrocarbons arising from the internal combustion engine and of their catalytic removal

(ii) gases that contribute to the enhanced greenhouse effect

c) outline the use of infra-red spectroscopy in monitoring air pollution

The inclusion of a halogen atom within an organic molecule affects its reactivity. The reactions of halogenoalkanes are very important in organic chemistry.

16.1 Halogenoalkanes

a) recall the chemistry of halogenoalkanes as exemplified by:

(i) the following nucleophilic substitution reactions of bromoethane: hydrolysis, formation of nitriles, formation of primary amines by reaction with ammonia

(ii) the elimination of hydrogen bromide from 2-bromopropane

b) describe the SN1 and SN2 mechanisms of nucleophilic substitution in halogenoalkanes including the inductive effects of alkyl groups

c) recall that primary halogenoalkanes tend to react via the SN2 mechanism; tertiary halogenoalkanes via the SN1 mechanism; and secondary halogenoalkanes by a mixture of the two, depending on structure

16.2 Relative strength of the C–Hal bond

a) interpret the different reactivities of halogenoalkanes (with particular reference to hydrolysis and to the relative strengths of the C–Hal bonds)

b) explain the uses of fluoroalkanes and fluorohalogenoalkanes in terms of their relative chemical inertness

c) recognise the concern about the effect of chlorofluoroalkanes on the ozone layer

This topic introduces the chemistry of a versatile class of organic compounds, hydroxy compounds, which contain an R–OH group.

17.1 Alcohols

a) recall the chemistry of alcohols, exemplified by ethanol, in the following reactions:

(i) combustion

(ii) substitution to halogenoalkanes

(iii) reaction with sodium

(iv) oxidation to carbonyl compounds and carboxylic acids

(v) dehydration to alkenes

(vi) formation of esters by esterification with carboxylic acids

b) (i) classify hydroxy compounds into primary, secondary and tertiary alcohols

(ii) suggest characteristic distinguishing reactions, e.g. mild oxidation

c) deduce the presence of a CH3CH(OH)– group in an alcohol from its reaction with alkaline aqueous iodine to form tri-iodomethane

This topic introduces the chemistry of the carbonyl compounds, aldehydes and ketones.

18.1 Aldehydes and ketones

a) describe:

(i) the formation of aldehydes and ketones from primary and secondary alcohols respectively using Cr2O7 2–/H+

(ii) the reduction of aldehydes and ketones, e.g. using NaBH4 or LiAlH4

(iii) the reaction of aldehydes and ketones with HCN and NaCN or KCN

b) describe the mechanism of the nucleophilic addition reactions of hydrogen cyanide with aldehydes and ketones

c) describe the use of 2,4-dinitrophenylhydrazine (2,4-DNPH) to detect the presence of carbonyl compounds

d) deduce the nature (aldehyde or ketone) of an unknown carbonyl compound from the results of simple tests (Fehling’s and Tollens’ reagents; ease of oxidation)

e) describe the reaction of CH3CO– compounds with alkaline aqueous iodine to give tri-iodomethane

This topic introduces the chemistry of carboxylic acids and their derivatives.

19.1 Carboxylic acids

a) describe the formation of carboxylic acids from alcohols, aldehydes and nitriles

b) describe the reactions of carboxylic acids in the formation of:

(i) salts, by the use of reactive metals, alkalis or carbonates

(ii) alkyl esters

(iii) alcohols, by the use of LiAlH4

19.2 Esters

a) describe the acid and base hydrolysis of esters

b) state the major commercial uses of esters, e.g. solvents, perfumes, flavourings