A-Level CIE Chemistry 9701 Objectives

2.1.a identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses
2.1.b deduce the behaviour of beams of protons, neutrons and electrons in electric fields
2.1.c describe the distribution of mass and charge within an atom
2.1.d deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers and charge
2.2.a describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number
2.2.b distinguish between isotopes on the basis of different numbers of neutrons present
2.2.c recognise and use the symbolism .^x_yA for isotopes, where x is the nucleon number and y is the proton number
1.1.a define and use the terms relative atomic, isotopic, molecular and formula masses, based on the 12C scale

2.3.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
2.3.b describe and sketch the shapes of s and p orbitals
2.3.c state the electronic configuration of atoms and ions given the proton number and charge, using the convention 1s²2s²2p⁶, etc.
1.2.a define and use the term mole in terms of the Avogadro constant
1.3.a analyse mass spectra in terms of isotopic abundances (knowledge of the working of the mass spectrometer is not required)
1.3.b calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum
1.4.a define and use the terms empirical and molecular formula
1.4.b calculate empirical and molecular formulae, using combustion data or composition by mass

1.5 a write and construct balanced equations
1.5.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 (see also Practical Assessment, Paper 3, Display of calculation and reasoning)
1.5.c deduce stoichiometric relationships from calculations such as those in 1.5.b

3.1.a describe ionic bonding, as in sodium chloride, magnesium oxide and calcium fluoride, including the use of ‘dot-and-cross' diagrams
3.2.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 Al₂Cl₆ molecule
3.2.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), CO₂ (linear), CH₄ (tetrahedral), NH₃ (pyramidal), H₂O (non-linear), SF₆ (octahedral), PF₅ (trigonal bipyramidal)
3.2.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 (see also LO 14.3)
3.2.d predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.2.b (see also LO 14.3)
14.3.a predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.2(b) (see also LO 14.3)
3.4.a describe metallic bonding in terms of a lattice of positive ions surrounded by de ocalised electrons

4.3.a describe, in simple terms, the lattice structure of a crystalline solid which is: (i) ionic, as in sodium chloride, magnesium oxide (ii) simple molecular, as in iodine and the fullerene allotropes of carbon (C 60 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
4.2.a describe, using a kinetic-molecular model, the liquid state, melting, vaporisation, vapour pressure
4.1.a state the basic assumptions of the kinetic theory as applied to an ideal gas
4.1.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
4.1.c state and use the general gas equation pV = nRT in calculations, including the determination of M_r

3.5.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
3.3.b understand, in simple terms, the concept of electronegativity and apply it to explain the properties of molecules such as bond polarity (see also LO 3.3.c), the dipole moments of molecules (LO 3.3.d) and the behaviour of oxides with water (LO 9.2.c)
3.3.d describe intermolecular forces (van der Waals' forces), based on permanent and induced dipoles, as in, for example, CHCl₃(l); Br₂(l) and the liquid Group 18 elements
3.3.a describe hydrogen bonding, using ammonia and water as simple examples of molecules containing N-H and O-H groups
3.5.b deduce the type of bonding present from given information
4.3.d suggest from quoted physical data the type of structure and bonding present in a substance
4.3.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)
4.3.b discuss the finite nature of materials as a resource and the importance of recycling processes

5.1.a explain that chemical reactions are accompanied by energy changes, principally in the form of heat energy; the energy changes can be exothermic (Δ/H/ is negative) or endothermic (Δ/H/ is positive)
3.5.c show understanding of chemical reactions in terms of energy transfers associated with the breaking and making of chemical bonds
3.3.c explain the terms bond energy, bond length and bond polarity and use them to compare the reactivities of covalent bonds (see also LO 5.1.b(ii))
5.1.b (i) and (ii) only explain and use the terms: (i) enthalpy change of reaction and standard conditions, with particular reference to: formation, combustion, hydration, solution, neutralisation, atomisation. (ii) bond energy (ΔH positive, i.e. bond breaking). (iii) lattice energy (ΔH negative, i.e. gaseous ions to solid lattice)
5.1.c calculate enthalpy changes from appropriate experimental results, including the use of the relationship enthalpy change, ΔH = -mcΔT
5.2.a (i) & (ii) only apply Hess' Law to construct simple energy cycles, and carry out calculations involving such cycles and relevant energy terms, with particular reference to: (i) (determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion (ii) average bond energies (iii) the formation of a simple ionic solid and of its aqueous solution (iv) Born-Haber cycles (including ionisation energy and electron affinity)
15.3.a describe and explain how the combustion reactions of alkanes led to their use as fuels in industry, in the home and in transport
15.3.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
13.2.a describe the formation of atmospheric sulfur dioxide from the combustion of sulfur-contaminated fossil fuels
13.2.b state the role of sulfur dioxide in the formation of acid rain and describe the main environmental consequences of acid rain
15.3.c outline the use of infra-red spectroscopy in monitoring air pollution (see also LO 22.2)
6.1.a calculate oxidation numbers of elements in compounds and ions
6.1.b describe and explain redox processes in terms of electron transfer and changes in oxidation number
6.1.c use changes in oxidation numbers to help balance chemical equations

8.1.a explain and use the term rate of reaction
8.1.b explain qualitatively, in terms of collisions, the effect of concentration changes on the rate of a reaction
5.2.b construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy
8.3.a explain and use the term catalysis
8.3.b explain that catalysts can be homogenous or heterogeneous
8.3.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
8.3.d describe enzymes as biological catalysts (proteins) which may have specificity
8.2.a explain and use the term activation energy, including reference to the Boltzmann distribution
8.2.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

7.1.a explain, in terms of rates of the forward and reverse reactions, what is meant by a reversible reaction and dynamic equilibrium
7.1.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
7.1.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
7.1.d deduce expressions for equilibrium constants in terms of concentrations, K_c, and partial pressures, K_p (treatment of the relationship between K_p and K_c is not required)
7.1.e calculate the values of equilibrium constants in terms of concentrations or partial pressures from appropriate data
7.1.f calculate the quantities present at equilibrium, given appropriate data (such calculations will not require the solving of quadratic equations)
7.1.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.a show understanding of, and use, the Brønsted-Lowry theory of acids and bases, including the use of the acid-I base-I, acid-II base-II concept
7.2.b explain qualitatively the differences in behaviour between strong and weak acids and bases and the pH values of their aqueous solutions in terms of the extent of dissociation

14.1.a AS Level compounds only interpret and use the general, structural, displayed and skeletal formulae of the following classes of compound: (i) alkanes, alkenes and arenes (ii) halogenoalkanes and halogenoarenes (iii) alcohols (including primary, secondary and tertiary) and phenols (iv) aldehydes and ketones (v) carboxylic acids, esters and acyl chlorides (vi) amines (primary only), nitriles, amides and amino acids (Candidates will be 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.)
14.1.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)
14.1.d deduce the possible isomers for an organic molecule of known molecular formula
14.1.e deduce the molecular formula of a compound, given its structural, displayed or skeletal formula
14.2.a interpret and use the following terminology associated with 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)

15.1.a understand the general unreactivity of alkanes, including towards polar reagents
15.1.b describe the chemistry of alkanes as exemplified by the following reactions of ethane: (i) combustion. (ii) substitution by chlorine and by bromine
15.1.c describe the mechanism of free-radical substitution at methyl groups with particular reference to the initiation, propagation and termination reactions
15.1.d explain the use of crude oil as a source of both aliphatic and aromatic hydrocarbons
15.1.e suggest how cracking can be used to obtain more useful alkanes and alkenes of lower M_r from larger hydrocarbon molecules
15.2.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 (see also Section 21)
15.2.b describe the mechanism of electrophilic addition in alkenes, using bromine/ethene and hydrogen bromide/propene as examples
15.2.c describe and explain the inductive effects of alkyl groups on the stability of cations formed during electrophilic addition
15.2.d describe the characteristics of addition polymerisation as exemplified by poly(ethene) and PVC
15.2.e deduce the repeat unit of an addition polymer obtained from a given monomer
15.2.f identify the monomer(s) present in a given section of an addition polymer molecule
15.2.g recognise the difficulty of the disposal of poly(alkene)s, i.e. nonbiodegradability and harmful combustion products
22.2.a analyse an infra-red spectrum of a simple molecule to identify functional groups (see the Data Booklet for functional groups required in the syllabus)
14.4.a describe structural isomerism and its division into chain, positional and functional group isomerism
14.4.b describe stereoisomerism and its division into geometrical (cis-trans) and optical isomerism (use of E, Z nomenclature is acceptable but is not required)
14.4.c describe cis-trans isomerism in alkenes, and explain its origin in terms of restricted rotation due to the presence of π bonds
14.4.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.)
14.4.e identify chiral centres and cis-trans isomerism in a molecule of given structural formula

16.1.a recall the chemistry of halogenoalkanes as exemplified by: hydrolysis, formation of nitriles, formation of primary amines by reaction with ammonia (ii) the elimination of hydrogen bromide from 2- bromopropane
16.1.b describe the S_N1 and S_N2 mechanisms of nucleophilic substitution in halogenoalkanes including the inductive effects of alkyl groups (see LO 15.2.c)
16.1.c recall that primary halogenoalkanes tend to react via the S_N2 mechanism; tertiary halogenoalkanes via the S_N1 mechanism and secondary halogenoalkanes by a mixture of the two, depending on structure
16.2.a interpret the different reactivities of halogenoalkanes (with particular reference to hydrolysis and to the relative strengths of the C-Hal bonds)
16.2.b explain the uses of fluoroalkanes and fluorohalogenoalkanes in terms of their relative chemical inertness
16.2.c recognise the concern about the effect of chlorofluoroalkanes on the ozone layer

17.1.a (i) to (vi) only recall the chemistry of alcohols, exemplified by ethanol, in the following reactions: (i) combustion. (ii) substitution to give 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. (vii) formation of esters by acylation with acyl chlorides using ethyl ethanoate and phenyl benzoate as examples
17.1.b (i) classify hydroxy compounds into primary, secondary and tertiary alcohols. (ii) suggest characteristic distinguishing reactions, e.g. mild oxidation
17.1.c deduce the presence of a CH₃CH(OH)- group in an alcohol from its reaction with alkaline aqueous iodine to form tri-iodomethane

18.1.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
18.1.b describe the mechanism of the nucleophilic addition reactions of hydrogen cyanide with aldehydes and ketones
18.1.c describe the use of 2,4-dinitrophenylhydrazine (2,4-DNPH) reagent to detect the presence of carbonyl compounds
18.1.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)
18.1.e describe the reaction of CH₃CO- compounds with alkaline aqueous iodine to give tri-iodomethane

19.1.a describe the formation of carboxylic acids from alcohols, aldehydes and nitriles
19.1.b (i) to (iii) only 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 use of LiAlH 4 (iv) acyl chlorides
19.3.a describe the acid and base hydrolysis of esters
19.3.b state the major commercial uses of esters, e.g. solvents, perfumes, flavourings

9.1.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)
9.1.b explain qualitatively the variation in atomic radius and ionic radius
9.1.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
2.3.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 (see also LOs 9.1)
9.1.d explain the variation in first ionisation energy (see the Data Booklet)
2.3.e deduce the electronic configurations of elements from successive ionisation energy data
2.3.f interpret successive ionisation energy data of an element in terms of the position of that element within the Periodic Table
9.2.a describe the reactions, if any, of the elements with oxygen (to give Na₂O, MgO, Al₂O₃, P₄O₁₀, SO₂, SO₃), chlorine (to give NaCl, MgCl₂, Al₂Cl₆, SiCl₄, PCl₅) and water (Na and Mg only)
9.2.b state and explain the variation in oxidation number of the oxides (sodium to sulfur only) and chlorides (sodium to phosphorus only) in terms of their valence shell electrons
9.2.c describe the reactions of the oxides with water (treatment of peroxides and superoxides is not required)
9.2.d describe and explain the acid/base behaviour of oxides and hydroxides including, where relevant, amphoteric behaviour in reaction with acids and bases (sodium hydroxide only)
9.2.e describe and explain the reactions of the chlorides with water
9.2.f interpret the variations and trends in 9.2(b), (c), (d) and (e) in terms of bonding and electronegativity
9.2.g suggest the types of chemical bonding present in chlorides and oxides from observations of their chemical and physical properties
9.1.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
9.3.a predict the characteristic properties of an element in a given Group by using knowledge of chemical periodicity
9.3.b deduce the nature, possible position in the Periodic Table and identity of unknown elements from given information about physical and chemical properties

10.1.a describe the reactions of the elements with oxygen, water and dilute acids
10.1.b describe the behaviour of the oxides, hydroxides and carbonates with water and dilute acids
10.1.c describe the thermal decomposition of the nitrates and carbonates
10.1.d interpret, and make predictions from, the trends in physical and chemical properties of the elements and their compounds
10.1.e state the variation in the solubilities of the hydroxides and sulfates
10.2.a describe and explain the use of calcium hydroxide and calcium carbonaten (powdered limestone) in agriculture

11.1.a describe the colours and the trend in volatility of chlorine, bromine and iodine
11.1.b interpret the volatility of the elements in terms of van der Waals' forces
11.2.a describe the relative reactivity of the elements as oxidising agents (see also Section 6.3.f)
11.2.b describe and explain the reactions of the elements with hydrogen
11.2.c (i) describe and explain the relative thermal stabilities of the hydrides. (ii) interpret these relative stabilities in terms of bond energies
11.3.a describe and explain the reactions of halide ions with: (i) aqueous silver ions followed by aqueous ammonia (ii) concentrated sulfuric acid
11.4.a describe and interpret, in terms of changes of oxidation number, the reaction of chlorine with cold and with hot aqueous sodium hydroxide
11.5.a explain the use of chlorine in water purification
11.5.b state the industrial importance and environmental significance of the halogens and their compounds (e.g. for bleaches, PVC, halogenated hydrocarbons as solvents, refrigerants and in aerosols. See also Section 16.2)

13.1.a explain the lack of reactivity of nitrogen
13.1.b describe and explain: (i) the basicity of ammonia (see also LOs 7.2). (ii) the structure of the ammonium ion and its formation by an acid-base reaction. (iii) the displacement of ammonia from its salts
13.1.c state the industrial importance of ammonia and nitrogen compounds derived from ammonia
13.1.d state and explain the environmental consequences of the uncontrolled use of nitrate fertilisers
13.1.e state and explain the natural and man-made occurrences of oxides of nitrogen and their catalytic removal from the exhaust gases of internal combustion engines
13.1.f explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulfur dioxide (see also LO 8.3.e (iii))

7.2.c explain the terms pH, Ka, pKa and Kw and use them in calculations
7.2.d calculate [H+(aq)] and pH values for strong and weak acids and strong bases
7.2.e explain the choice of suitable indicators for acid-base titrations, given appropriate data
7.2.f describe the changes in pH during acid-base titrations and explain these changes in terms of the strengths of the acids and bases
7.2.g (i) explain how buffer solutions control pH (ii) describe and explain the uses of buffer solutions, including the role of HCO₃ in controlling pH in blood
7.2.h calculate the pH of buffer solutions, given appropriate data
7.2.i show understanding of, and use, the concept of solubility product, Ksp
7.2.j calculate Ksp from concentrations and vice versa
7.2.k show understanding of the common ion effect

6.2.a state and apply the relationship F=Le between the Faraday constant, the Avogadro constant and the charge on the electron
6.2.b predict the identity of the substance liberated during electrolysis from the state of electrolyte (molten or aqueous), position in the redox series (electrode potential) and concentration
6.2.c calculate: (i) the quantity of charge passed during electrolysis (ii) the mass and/or volume of substance liberated during electrolysis, including those in the electrolysis of H₂SO₄ (aq) and of Na₂SO₄ (aq)
6.2.d describe the determination of a value of the Avogadro constant by an electrolytic method
6.3.a define the terms: (i) standard electrode (redox) potential (ii) standard cell potential
6.3.b describe the standard hydrogen electrode
6.3.c describe methods used to measure the standard electrode potentials of: (i) metals or non-metals in contact with their ions in aqueous solution (ii) ions of the same element in different oxidation states
6.3.d calculate a standard cell potential by combining two standard electrode potentials
6.3.e use standard cell potentials to: (i) explain/deduce the direction of electron flow in a simple cell (ii) predict the feasibility of a reaction
6.3.f deduce from E^⦵ values the relative reactivity of elements of Group 17 (the halogens) as oxidising agents
6.3.g construct redox equations using the relevant half-equations (see also LO 12.2.d)
12.2.d describe and explain the use of Fe³⁺/Fe²⁺, MnO4^-/Mn²⁺ and Cr2O7^2-/Cr³⁺ as examples of redox systems (see also LO 6.3)
12.2.e predict, using E^⦵ values, the likelihood of redox reactions
6.3.h predict qualitatively how the value of an electrode potential varies with the concentrations of the aqueous ions
6.3.i use the Nernst equation, e.g. E = E⦵ (0.059/z) log [oxidised species] / [reduced species] to predict quantitatively how the value of an electrode potential varies with the concentrations of the aqueous ions; examples include Cu(s) + 2e^- ⇌ Cu²⁺(aq), Fe³⁺(aq) + e^- ⇌Fe²⁺(aq), Cl₂(g) + 2e^- ⇌ 2Cl^-(aq)
6.4.a state the possible advantages of developing other types of cell, e.g. the H₂/O₂ fuel cell and the nickel-metal hydride and lithium-ion rechargeable batteries

12.1.a explain what is meant by a transition element, in terms of d-block elements forming one or more stable ions with incomplete d orbitals
12.1.b sketch the shape of a d orbital
12.1.c state the electronic configuration of each of the first row transition elements and of their ions
12.1.d contrast, qualitatively, the melting points and densities of the transition elements with those of calcium as a typical s-block element
12.1.e describe the tendency of transition elements to have variable oxidation states
12.1.f predict from a given electronic configuration, the likely oxidation states of a transition element
12.2.a describe and explain the reactions of transition elements with ligands to form complexes, including the complexes of copper(II) and cobalt(II) ions with water and ammonia molecules and hydroxide and chloride ions
12.2.b (i) define the term ligand as a species that contains a lone pair of electrons that forms a dative bond to a central metal atom/ion including monodentate, bidentate and polydentate ligands. (ii) define the term complex as a molecule or ion formed by a central metal atom/ion surrounded by one or more ligands. (iii) describe transition metal complexes as linear, octahedral, tetrahedral or square planar. (iv) state what is meant by co-ordination number and predict the formula and charge of a complex ion, given the metal ion, its charge, the ligand and its co-ordination number.
12.2.c explain qualitatively that ligand exchange may occur, including the complexes of copper(II) ions with water and ammonia molecules and hydroxide and chloride ions
12.3.a describe the splitting of degenerate d orbitals into two energy levels in octahedral and tetrahedral complexes
12.3.b explain the origin of colour in transition element complexes resulting from the absorption of light energy as an electron moves between two non-degenerate d orbitals
12.3.c describe, in qualitative terms, the effects of different ligands on absorption, and hence colour, using the complexes of copper(II) ions with water and ammonia molecules and hydroxide and chloride ions as examples
12.3.d apply the above ideas of ligands and complexes to other metals, given information
12.5.a describe and explain ligand exchanges in terms of competing equilibria (also see syllabus topic 7)
12.5.b state that the stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions or molecules
12.5.c deduce expressions for the stability constant of a ligand substitution
12.5.d explain ligand exchange in terms of stability constants, Kstab, and understand that a large Kstab is due to the formation of a stable complex ion

5.1.b (iii) only explain and use the terms: (i) enthalpy change of reaction and standard conditions, with particular reference to: formation, combustion, hydration, solution, neutralisation, atomisation (ii) bond energy ( Δ H positive, i.e. bond breaking) (iii) lattice energy ( Δ H negative, i.e. gaseous ions to solid lattice)
2.3.g explain and use the term electron affinity
5.1.d explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy
5.2.a (iii) & (iv) only apply Hess' Law to construct simple energy cycles, and carry out calculations involving such cycles and relevant energy terms, with particular reference to: (i) (determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion (ii) average bond energies (iii) the formation of a simple ionic solid and of its aqueous solution (iv) Born-Haber cycles (including ionisation energy and electron affinity)
10.1.f interpret and explain qualitatively the trend in the thermal stability of the nitrates and carbonates in terms of the charge density of the cation and the polarisability of the large anion
10.1.g interpret and explain qualitatively the variation in solubility of the hydroxides and sulfates in terms of relative magnitudes of the enthalpy change of hydration and the corresponding lattice energy
5.3.a explain that entropy is a measure of the ‘disorder' of a system, and that a system becomes more stable when its energy is spread out in a more disordered state
5.3.b explain the entropy changes that occur: (i) during a change in state e.g. (s) → (I); (I) → (g); (s) → (aq) (ii) during a temperature change (iii) during a reaction in which there is a change in the number of gaseous molecules
5.3.c predict whether the entropy change for a given process is positive or negative
5.3.d calculate the entropy change for a reaction, ΔS^⦵, given the standard entropies, S^⦵, of the reactants and products
5.4.a define standard Gibbs free energy change of reaction by means of the equation }ΔG^⦵ = ΔH^⦵ - TΔS^⦵
5.4.b calculate G^⦵ for a reaction using the equation ΔG^⦵ = ΔH^⦵ - TΔS^⦵
5.4.c state whether a reaction or process will be spontaneous by using the sign of ΔG^⦵
5.4.d predict the effect of temperature change on the spontaneity of a reaction, given standard enthalpy and entropy changes }

8.1.c explain and use the terms rate equation, order of reaction, rate constant, half-life of a reaction, rate-determining step
8.1.d construct and use rate equations of the form rate = k[A]m[B]n (for which m and n are 0, 1 or 2), including: (i) deducing the order of a reaction, or the rate equation for a reaction, from concentration-time graphs or from experimental data relating to the initial rates method and half-life method (ii) interpreting experimental data in graphical form, including concentration-time and rate-concentration graphs (iii) calculating an initial rate using concentration data. integrated forms of rate equations are not required
8.1.e (i) show understanding that the half-life of a first-order reaction is independent of concentration (ii) use the half-life of a first-order reaction in calculations
8.1.f calculate the numerical value of a rate constant, for example by using the initial rates or half-life method
8.1.g for a multi-step reaction: (i) suggest a reaction mechanism that is consistent with the rate equation and the equation for the overall reaction (ii) predict the order that would result from a given reaction mechanism (and vice versa)
8.1.h devise a suitable experimental technique for studying the rate of a reaction, from given information
8.2.c explain qualitatively the effect of temperature change on a rate constant and hence the rate of a reaction

14.1.a (A Level compounds only) interpret and use the general, structural, displayed and skeletal formulae of the following classes of compound: (i) alkanes, alkenes and arenes. (ii) halogenoalkanes and halogenoarenes. (iii) alcohols (including primary, secondary and tertiary) and phenols. (iv) aldehydes and ketones. (v) carboxylic acids, esters and acyl chlorides. (vi) amines (primary only), nitriles, amides and amino acids. (Candidates will be 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.)
15.4.a describe the chemistry of arenes as exemplified by the following reactions of benzene and methylbenzene: (i) substitution reactions with chlorine and with bromine. (ii) nitration. (iii) Friedel-Crafts alkylation and acylation. (iv) complete oxidation of the side-chain to give a benzoic acid. (v) hydrogenation of the benzene ring to form a cyclohexane ring.
15.4.b (i) describe the mechanism of electrophilic substitution in arenes, as exemplified by the formation of nitrobenzene and bromobenzene (ii) suggest the mechanism of other electrophilic substitution reactions, given data (iii) describe the effect of the delocalisation of electrons in arenes in such reactions
15.4.c interpret the difference in reactivity between benzene and chlorobenzene
15.4.d predict whether halogenation will occur in the side-chain or in the aromatic ring in arenes depending on reaction conditions
15.4.e apply knowledge relating to position of substitution in the electrophilic substitution of arenes (see the syllabus Data Booklet Table 9)
14.1.c understand and use systematic nomenclature of simple aromatic molecules with one benzene ring and one or more simple substituents, for example 3-nitrobenzoic acid, 2,4,6-tribromophenol

17.2.a recall the chemistry of phenol, as exemplified by the following reactions: (i) with bases (ii) with sodium (iii) with diazonium salts (see also Section 20.1) (iv) nitration of, and bromination of, the aromatic ring
17.2.b describe and explain the relative acidities of water, phenol and ethanol

19.1.b (iv) only 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 use of LiAlH₄ (iv) acyl chlorides
19.1.c recognise that some carboxylic acids can be further oxidised: (i) the oxidation of methanoic acid, HCO₂H, with Fehling’s and Tollens’ reagents (ii) the oxidation of ethanedioic acid, HO₂CCO₂H, with warm acidified manganate(VII)
19.1.d explain the relative acidities of carboxylic acids, phenols and alcohols
19.1.e use the concept of electronegativity to explain the acidities of chlorine-substituted ethanoic acids
19.2.a describe the hydrolysis of acyl chlorides
19.2.b describe the reactions of acyl chlorides with alcohols, phenols, ammonia and primary amines
17.1.a (vii) only recall the chemistry of alcohols, exemplified by ethanol, in the following reactions: (i) combustion (ii) substitution to give 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 (vii) formation of esters by acylation with acyl chlorides using ethyl ethanoate and phenyl benzoate as examples
19.2.c explain the relative ease of hydrolysis of acyl chlorides, alkyl chlorides and aryl chlorides including the condensation (additionelimination) mechanism for the hydrolysis of acyl chlorides
20.1.a describe the formation of alkyl amines such as ethylamine (by the reaction of ammonia with halogenoalkanes; the reduction of amides with LiAlH4; the reduction of nitriles with LiAlH4 or H2/Ni) and of phenylamine (by the reduction of nitrobenzene with tin/concentrated HCl)
20.1.b describe and explain the basicity of amines
20.1.c explain the relative basicities of ammonia, ethylamine and phenylamine in terms of their structures
20.1.d describe the reaction of phenylamine with: (i) aqueous bromine (ii) nitrous acid to give the diazonium salt and phenol
20.1.e describe the coupling of benzenediazonium chloride and phenol and the use of similar reactions in the formation of dyestuff
20.2.a describe the formation of amides from the reaction between NH3 or RNH2 and R'COCl
20.2.b recognise that amides are neutral
20.2.c (i) describe amide hydrolysis on treatment with aqueous alkali or acid. (ii) describe the reduction of amides with LiAlH₄
20.3.a describe the acid/base properties of amino acids and the formation of zwitterions
20.3.b describe the formation of peptide bonds between amino acids to give di- and tri-peptides
20.3.c describe simply the process of electrophoresis and the effect of pH, using peptides and amino acids as examples

8.3.e outline the different characteristics and modes of action of homogeneous, heterogeneous and enzyme catalysts, including: (i) the Haber process (ii) the catalytic removal of oxides of nitrogen from the exhaust gases of car engines (see also LO 15.3.b(i)) (iii) the catalytic role of atmospheric oxides of nitrogen in the oxidation of atmospheric sulfur dioxide (see also LO 13.1.f) (iv) the catalytic role of Fe 2+ or Fe 3+ in the I – /S 2 O 82– reaction (v) the catalytic role of enzymes (including the explanation of specificity using a simple lock and key model but excluding inhibition)
23.1.a state that most chiral drugs extracted from natural sources often contain only a single optical isomer
23.1.b state reasons why the synthetic preparation of drug molecules often requires the production of a single optical isomer, e.g. better therapeutic activity, fewer side effects
12.4.a describe the types of stereoisomerism shown by complexes, including those associated with bidentate ligands: (i) cis-trans isomerism, e.g. cis-and trans-platin Pt(NH₃)₂Cl₂ (ii) optical isomerism, e.g. [Ni(NH₂CH₂CH₂NH₂)₃]²⁺
12.4.b describe the use of cisplatin as an anticancer drug and its action by binding to DNA in cancer cells, preventing cell division
21.1.a describe the formation of polyesters and polyamides
21.1.b describe the characteristics of condensation polymerisation; (i) in polyesters as exemplified by Terylene (ii) in polyamides as exemplified by polypeptides, proteins, nylon 6, nylon 6,6 and Kevlar
21.1.c deduce the repeat unit of a condensation polymer obtained from a given monomer or pair of monomers
21.1.d identify the monomer(s) present in a given section of a condensation polymer molecule
21.2.a predict the type of polymerisation reaction for a given monomer or pair of monomers
21.2.b deduce the type of polymerisation reaction which produces a given section of a polymer molecule
21.3.a discuss the properties and structure of polymers based on their methods of formation (addition or condensation, see also Section 15.2)
21.3.b discuss how the presence of side-chains and intermolecular forces affect the properties of polymeric materials (e.g. polyalkenes, PTFE (Teflon),Kevlar)
21.3.c explain the significance of hydrogen-bonding in the pairing of bases in DNA in relation to the replication of genetic information
21.4.a recognise that polyalkenes are chemically inert and can therefore be difficult to biodegrade
21.4.b recognise that a number of polymers can be degraded by the action of light
21.4.c recognise that polyesters and polyamides are biodegradable by hydrolysis
21.4.d describe the hydrolysis of proteins
23.2.a for an organic molecule containing several functional groups: (i) identify organic functional groups using the reactions in the syllabus (ii) predict properties and reactions
23.2.b devise multi-stage synthetic routes for preparing organic molecules using the reactions in the syllabus
23.2.c analyse a given synthetic route in terms of type of reaction and reagents used for each step of it, and possible by-products
21.3.d distinguish between the primary, secondary (α-helix and β-sheet) and tertiary structures of proteins and explain the stabilisation of secondary structure (through hydrogen bonding between C=O and N-H bonds of peptide groups) and tertiary structure (through interactions between R-groups)
21.3.e describe how polymers have been designed to act as: (i) non-solvent based adhesives, e.g. epoxy resins and superglues (ii) conducting polymers, e.g. polyacetylene

22.3.a deduce the molecular mass of an organic molecule from the molecular ion peak in a mass spectrum
22.3.b deduce the number of carbon atoms in a compound using the M+1 peak
22.3.c deduce the presence of bromine and chlorine atoms in a compound using the M+2 peak
22.3.d suggest the identity of molecules formed by simple fragmentation in a given mass spectrum
22.1.a explain and use the terms Rf value in thin layer chromatography and retention time in gas/liquid chromatography from chromatograms
22.1.b interpret gas/liquid chromatograms in terms of the percentage composition of a mixture
7.3.a state what is meant by partition coefficient; calculate and use a partition coefficient for a system in which the solute is in the same molecular state in the two solvents
22.4.a analyse a carbon-13 NMR spectrum of a simple molecule to deduce: (i) the different environments of the carbon atoms present (ii) the possible structures for the molecule
22.4.b predict the number of peaks in a carbon-13 NMR spectrum for a given molecule
22.5.a analyse and interpret a proton NMR spectrum of a simple molecule to deduce: (i) the different types of proton present using chemical shift values (ii) the relative numbers of each type of proton present from relative peak areas (iii) the number of non- equivalent protons adjacent to a given proton from the splitting pattern, using the n + 1 rule (iv) the possible structures for the molecule
22.5.b predict the chemical shifts and splitting patterns of the protons in a given molecule
22.5.c describe the use of tetramethylsilane, TMS, as the standard for chemical shift measurements
22.5.d state the need for deuterated solvents, e.g. CDCl3, when obtaining an NMR spectrum
22.5.e describe the identification of O-H and N-H protons by proton exchange using D2O