Chemistry 9701 (2022, 2023, & 2024): A Level

Physical chemistry

Inorganic chemistry

Organic chemistry

Analysis

Chemical energetics

Electrochemistry

Equilibria

Reaction kinetics

Group 2

Chemistry of transition elements

An introduction to A Level organic chemistry

Hydrocarbons

Halogen compounds

Hydroxy compounds

Carboxylic acids and derivatives

Nitrogen compounds

Polymerisation

Organic synthesis

Analytical techniques

Lattice energy and Born-Haber cycles

Enthalpies of solution and hydration 

Entropy change, ΔS 

Gibbs free energy change, ΔG

Electrolysis

Standard electrode potentials E ⦵ ; standard cell potentials E ⦵ cell and the Nernst equation

Acids and bases

Partition coefficients

Simple rate equations, orders of reaction and rate constants

Homogeneous and heterogeneous catalysts 

Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds

General physical and chemical properties of the first row of transition elements, titanium to copper

General characteristic chemical properties of the first set of transition elements, titanium to copper

Colour of complexes 

Stereoisomerism in transition element complexes

Stability constants, Kstab

Formulae, functional groups and the naming of organic compounds  

Characteristic organic reactions

Shapes of aromatic organic molecules; σ and π bonds

Isomerism: optical

Arenes

Halogen compounds

Alcohols

Phenol

Carboxylic acids

Esters

Acyl chlorides

Primary and secondary amines

Phenylamine and azo compounds

Amides

Amino acids

Condensation polymerisation

Predicting the type of polymerisation

Degradable polymers

Organic synthesis

Thin-layer chromatography

Gas /liquid chromatography 

Carbon-13 NMR spectroscopy

Proton (1 H) NMR spectroscopy

define and use the terms:  

define and use the term first electron affinity, EA

construct and use Born–Haber cycles for ionic solids  (limited to +1 and +2 cations, –1 and –2 anions)

carry out calculations involving Born–Haber cycles

explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy  

define and use the term enthalpy change with reference to hydration, ΔHhyd, and solution, ΔHsol

construct and use an energy cycle involving enthalpy change of solution, lattice energy and enthalpy change of hydration

carry out calculations involving the energy cycles in 23.2.2

explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of an enthalpy change of hydration 

define the term entropy, S, as the number of possible arrangements of the particles and their energy in a given system 

predict and explain the sign of the entropy changes that occur:

calculate the entropy change for a reaction, ΔS, given the standard entropies, S ⦵ , of the reactants and products, ΔS ⦵ = ΣS ⦵ (products) – ΣS ⦵ (reactants) (use of ΔS ⦵ = ΔSsurr⦵ + ΔSsys⦵ is not required)

state and use the Gibbs equation ΔG⦵ = ΔH⦵ – TΔS ⦵

perform calculations using the equation ΔG⦵ = ΔH⦵ – TΔS ⦵

state whether a reaction or process will be feasible by using the sign of ΔG

predict the effect of temperature change on the feasibility of a reaction, given standard enthalpy and entropy changes

predict the identities of substances liberated during electrolysis from the state of electrolyte (molten or aqueous), position in the redox series (electrode potential) and concentration

state and apply the relationship F = Le between the Faraday constant, F, the Avogadro constant, L, and the charge on the electron, e

calculate:

describe the determination of a value of the Avogadro constant by an electrolytic method

define the terms:

describe the standard hydrogen electrode

describe methods used to measure the standard electrode potentials of: 

calculate a standard cell potential by combining two standard electrode potentials

use standard cell potentials to:

deduce from E ⦵ values the relative reactivity of elements, compounds and ions as oxidising agents or as reducing agents 

construct redox equations using the relevant half-equations 

predict qualitatively how the value of an electrode potential, E, varies with the concentrations of the aqueous ions

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 Cu2+(aq) + 2e– ⇌ Cu(s), Fe3+(aq) + e– ⇌ Fe2+(aq) 

understand and use the equation ΔG⦵ = –nE ⦵ cell F

understand and use the terms conjugate acid and conjugate base

define conjugate acid–base pairs, identifying such pairs in reactions

define mathematically the terms pH, Ka , pKa and Kw and use them in calculations (Kb and the equation Kw = Ka × Kb will not be tested) 

calculate [H+ (aq)] and pH values for:

define a buffer solution 

calculate the pH of buffer solutions, given appropriate data

understand and use the term solubility product, K sp

write an expression for K sp

calculate K sp from concentrations and vice versa

understand and use the common ion effect to explain the different solubility of a compound in a solution containing a common ion

state what is meant by the term partition coefficient, K pc

calculate and use a partition coefficient for a system in which the solute is in the same physical state in the two solvents

understand the factors affecting the numerical value of a partition coefficient in terms of the polarities of the solute and the solvents used 

explain and use the terms rate equation, order of reaction, overall order of reaction, rate constant, half-life, rate-determining step and intermediate

understand and use rate equations of the form rate = k [A]m[B]n (for which m and n are 0, 1 or 2)

show understanding that the half-life of a first-order reaction is independent of concentration 

calculate the numerical value of a rate constant, for example by:

for a multi-step reaction:

describe qualitatively the effect of temperature change on the rate constant and hence the rate of a reaction 

explain that catalysts can be homogeneous or heterogeneous 

describe the mode of action of a heterogeneous catalyst to include adsorption of reactants, bond weakening and desorption of products, for example: 

describe the mode of action of a homogeneous catalyst by being used in one step and reformed in a later step, for example:

describe and explain qualitatively the trend in the thermal stability of the nitrates and carbonates including the effect of ionic radius on the polarisation of the large anion

describe and explain qualitatively the variation in solubility and of enthalpy change of solution, ΔH⦵ sol, of the hydroxides and sulfates in terms of relative magnitudes of the enthalpy change of hydration and the lattice energy

define a transition element as a d-block element which forms one or more stable ions with incomplete d orbitals

sketch the shape of a 3dxy orbital and 3dz2 orbital 

understand that transition elements have the following properties: 

explain why transition elements have variable oxidation states in terms of the similarity in energy of the 3d and the 4s sub-shells

explain why transition elements behave as catalysts in terms of having more than one stable oxidation state, and vacant d orbitals that are energetically accessible and can form dative bonds with ligands

explain why transition elements form complex ions in terms of vacant d orbitals that are energetically accessible 

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  

define the term ligand as a species that contains a lone pair of electrons that forms a dative covalent bond to a central metal atom/ion 

understand and use the terms

define the term complex as a molecule or ion formed by a central metal atom/ion surrounded by one or more ligands

describe the geometry (shape and bond angles) of transition element complexes which are linear, square planar, tetrahedral or octahedral 

state what is meant by coordination number 

explain qualitatively that ligand exchange can occur, including the complexes of copper(II) ions and cobalt(II) ions with water and ammonia molecules and hydroxide and chloride ions

predict, using E ⦵ values, the feasibility of redox reactions involving transition elements and their ions

describe the reactions of, and perform calculations involving:  

perform calculations involving other redox systems given suitable data 

define and use the terms degenerate and non-degenerate d orbitals  

describe the splitting of degenerate d orbitals into two non-degenerate sets of d orbitals of higher energy, and use of ΔE in:

explain why transition elements form coloured compounds in terms of the frequency of light absorbed as an electron is promoted between two non-degenerate d orbitals

describe, in qualitative terms, the effects of different ligands on ΔE, frequency of light absorbed, and hence the complementary colour that is observed

use the complexes of copper(II) ions and cobalt(II) ions with water and ammonia molecules and hydroxide and chloride ions as examples of ligand exchange affecting the colour observed

describe the types of stereoisomerism shown by complexes, including those associated with bidentate ligands:

deduce the overall polarity of complexes such as those described in 28.4.1(a) and 28.4.1(b) 

define the stability constant, Kstab, of a complex as the equilibrium constant for the formation of the complex ion in a solvent (from its constituent ions or molecules) 

write an expression for a Kstab of a complex ([H2O] should not be included)

use Kstab expressions to perform calculations

describe and explain ligand exchanges in terms of Kstab values and understand that a large Kstab is due to the formation of a stable complex ion

understand that the compounds in the table on page 42 contain a functional group which dictates their physical and chemical properties

interpret and use the general, structural, displayed and skeletal formulae of the classes of compound stated in the table on page 42

understand and use systematic nomenclature of simple aliphatic organic molecules (including cyclic compounds containing a single ring of up to six carbon atoms) with functional groups detailed in the table on page 42, up to six carbon atoms (six plus six for esters and amides, straight chains only for esters and nitriles)

understand and use systematic nomenclature of simple aromatic molecules with one benzene ring and one or more simple substituents, for example 3-nitrobenzoic acid or 2,4,6-tribromophenol 

understand and use the following terminology associated with types of organic mechanisms:

describe and explain the shape of benzene and other aromatic molecules, including sp2 hybridisation, in terms of σ bonds and a delocalised π system

understand that enantiomers have identical physical and chemical properties apart from their ability to rotate plane polarised light and their potential biological activity

understand and use the terms optically active and racemic mixture

describe the effect on plane polarised light of the two optical isomers of a single substance

explain the relevance of chirality to the synthetic preparation of drug molecules including: 

describe the chemistry of arenes as exemplified by the following reactions of benzene and methylbenzene:

describe the mechanism of electrophilic substitution in arenes:

predict whether halogenation will occur in the side-chain or in the aromatic ring in arenes depending on reaction conditions

describe that in the electrophilic substitution of arenes, different substituents direct to different ring positions (limited to the directing effects of –NH2, –OH, –R, –NO2, –COOH and –COR)

recall the reactions by which halogenoarenes can be produced: substitution of an arene with Cl 2 or Br2 in the presence of a catalyst, AlCl 3 or AlBr3 to form a halogenoarene, exemplified by benzene to form chlorobenzene and methylbenzene to form 2-chloromethylbenzene and 4-chloromethylbenzene

explain the difference in reactivity between a halogenoalkane and a halogenoarene as exemplified by chloroethane and chlorobenzene

describe the reaction with acyl chlorides to form esters using ethyl ethanoate

recall the reactions (reagents and conditions) by which phenol can be produced:  

recall the chemistry of phenol, as exemplified by the following reactions:

explain the acidity of phenol

describe and explain the relative acidities of water, phenol and ethanol

explain why the reagents and conditions for the nitration and bromination of phenol are different from those for benzene

recall that the hydroxyl group of a phenol directs to the 2-, 4- and 6-positions

apply knowledge of the reactions of phenol to those of other phenolic compounds, e.g. naphthol

recall the reaction by which benzoic acid can be produced:

describe the reaction of carboxylic acids with PCl 3 and heat, PCl 5, or SOCl 2 to form acyl chlorides

recognise that some carboxylic acids can be further oxidised:

describe and explain the relative acidities of carboxylic acids, phenols and alcohols

describe and explain the relative acidities of chlorine-substituted carboxylic acids

recall the reaction by which esters can be produced:

recall the reactions (reagents and conditions) by which acyl chlorides can be produced:

describe the following reactions of acyl chlorides:

describe the addition-elimination mechanism of acyl chlorides in reactions in 33.3.2(a) – (e) 

explain the relative ease of hydrolysis of acyl chlorides, alkyl chlorides and halogenoarenes (aryl chlorides)

recall the reactions (reagents and conditions) by which primary and secondary amines are produced:

describe the condensation reaction of ammonia or an amine with an acyl chloride at room temperature to give an amide

describe and explain the basicity of aqueous solutions of amines

describe the preparation of phenylamine via the nitration of benzene to form nitrobenzene followed by reduction with hot Sn/concentrated HCl, followed by NaOH(aq)

describe: 

describe and explain the relative basicities of aqueous ammonia, ethylamine and phenylamine

recall the following about azo compounds:

recall the reactions (reagents and conditions) by which amides are produced:

describe the reactions of amides: 

state and explain why amides are much weaker bases than amines

describe the acid/ base properties of amino acids and the formation of zwitterions, to include the isoelectric point

describe the formation of amide (peptide) bonds between amino acids to give di- and tripeptides

interpret and predict the results of electrophoresis on mixtures of amino acids and dipeptides at varying pHs (the assembling of the apparatus will not be tested)

describe the formation of polyesters: 

describe the formation of polyamides: 

deduce the repeat unit of a condensation polymer obtained from a given monomer or pair of monomers

identify the monomer(s) present in a given section of a condensation polymer molecule

predict the type of polymerisation reaction for a given monomer or pair of monomers

deduce the type of polymerisation reaction which produces a given section of a polymer molecule

recognise that poly(alkenes) are chemically inert and can therefore be difficult to biodegrade 

recognise that some polymers can be degraded by the action of light

recognise that polyesters and polyamides are biodegradable by acidic and alkaline hydrolysis  

for an organic molecule containing several functional groups:

devise multi-step synthetic routes for preparing organic molecules using the reactions in the syllabus

analyse a given synthetic route in terms of type of reaction and reagents used for each step of it, and possible by-products

describe and understand the terms

interpret Rf values

explain the differences in Rf values in terms of interaction with the stationary phase and of relative solubility in the mobile phase  

describe and understand the terms 

interpret gas/liquid chromatograms in terms of the percentage composition of a mixture

explain retention times in terms of interaction with the stationary phase

analyse and interpret a carbon-13 NMR spectrum of a simple molecule to deduce: 

predict or explain the number of peaks in a carbon-13 NMR spectrum for a given molecule

analyse and interpret a proton (1 H) NMR spectrum of a simple molecule to deduce:

predict the chemical shifts and splitting patterns of the protons in a given molecule

describe the use of tetramethylsilane, TMS, as the standard for chemical shift measurements

state the need for deuterated solvents, e.g. CDCl 3, when obtaining a proton NMR spectrum

describe the identification of O–H and N–H protons by proton exchange using D2O

enthalpy change of atomisation, ΔHat

lattice energy, ΔHlatt (the change from gas phase ions to solid lattice) 

define and use the term first electron affinity, EA

explain the factors affecting the electron affinities of elements

describe and explain the trends in the electron affinities of the Group 16 and Group 17 elements  

during a change in state, e.g. melting, boiling and dissolving (and their reverse)

during a temperature change  

during a reaction in which there is a change in the number of gaseous molecules

the quantity of charge passed during electrolysis, using Q = It

the mass and/or volume of substance liberated during electrolysis

standard electrode (reduction) potential 

standard cell potential 

metals or non-metals in contact with their ions in aqueous solution

ions of the same element in different oxidation states

 deduce the polarity of each electrode and hence explain/deduce the direction of electron flow in the external circuit of a simple cell

predict the feasibility of a reaction 

strong acids 

strong alkalis

weak acids 

define a buffer solution 

explain how a buffer solution can be made 

explain how buffer solutions control pH; use chemical equations in these explanations

describe and explain the uses of buffer solutions, including the role of HCO3 – in controlling pH in blood  

understand and use the common ion effect to explain the different solubility of a compound in a solution containing a common ion

perform calculations using K sp values and concentration of a common ion 

understand and use rate equations of the form rate = k [A]m[B]n (for which m and n are 0, 1 or 2)

deduce the order of a reaction from concentration-time graphs or from experimental data relating to the initial rates method and half-life method

interpret experimental data in graphical form, including concentration-time and rate-concentration graphs

calculate an initial rate using concentration data 

construct a rate equation  

show understanding that the half-life of a first-order reaction is independent of concentration 

use the half-life of a first-order reaction in calculations

using the initial rates and the rate equation 

using the half-life, t1/2, and the equation k = 0.693/t1

suggest a reaction mechanism that is consistent with the rate equation and the equation for the overall reaction

predict the order that would result from a given reaction mechanism and rate-determining step

deduce a rate equation using a given reaction mechanism and rate-determining step for a given reaction 

identify an intermediate or catalyst from a given reaction mechanism 

identify the rate determining step from a rate equation and a given reaction mechanism

iron in the Haber process

palladium, platinum and rhodium in the catalytic removal of oxides of nitrogen from the exhaust gases of car engines 

atmospheric oxides of nitrogen in the oxidation of atmospheric sulfur dioxide

Fe2+ or Fe3+ in the I – /S2O8 2– reaction

they have variable oxidation states

they behave as catalysts

they form complex ions

they form coloured compounds

monodentate ligand including as examples H2O, NH3, Cl – and CN–

bidentate ligand including as examples 1,2-diaminoethane, en, H2NCH2CH2NH2 and the ethanedioate ion, C2O4 2– 

polydentate ligand including as an example EDTA4–

state what is meant by coordination number 

predict the formula and charge of a complex ion, given the metal ion, its charge or oxidation state, the ligand and its coordination number or geometry  

MnO4 – /C2O4 2– in acid solution given suitable data  

MnO4 – / Fe2+ in acid solution given suitable data 

Cu2+ / I – given suitable data

octahedral complexes, two higher and three lower d orbitals

tetrahedral complexes, three higher and two lower d orbitals

 geometrical (cis-trans) isomerism, e.g. square planar such as [Pt(NH3) 2Cl 2] and octahedral such as [Co(NH3) 4(H2O)2] 2+ and [Ni(H2NCH2CH2NH2) 2(H2O)2] 2+ 

optical isomerism, e.g. [Ni(H2NCH2CH2NH2) 3] 2+ and [Ni(H2NCH2CH2NH2) 2(H2O)2] 2+ 

electrophilic substitution

addition-elimination 

the potential different biological activity of the two enantiomers 

the need to separate a racemic mixture into two pure enantiomers  

the use of chiral catalysts to produce a single pure optical isomer (Candidates should appreciate that compounds can contain more than one chiral centre, but knowledge of meso compounds and nomenclature such as diastereoisomers is not required.)

substitution reactions with Cl 2 and with Br2 in the presence of a catalyst, AlCl 3 or AlBr3, to form halogenoarenes (aryl halides) 

nitration with a mixture of concentrated HNO3 and concentrated H2SO4 at a temperature between 25°C and 60°C

Friedel–Crafts alkylation by CH3Cl and AlCl 3 and heat

Friedel–Crafts acylation by CH3COCl and AlCl 3 and heat

complete oxidation of the side-chain using hot alkaline KMnO4 and then dilute acid to give a benzoic acid

hydrogenation of the benzene ring using H2 and Pt/Ni catalyst and heat to form a cyclohexane ring

as exemplified by the formation of nitrobenzene and bromobenzene

with regards to the effect of delocalisation (aromatic stabilisation) of electrons in arenes to explain the predomination of substitution over addition

reaction of phenylamine with HNO2 or NaNO2 and dilute acid below 10°C to produce the diazonium salt; further warming of the diazonium salt with H2O to give phenol

with bases, for example NaOH(aq) to produce sodium phenoxide

with Na(s) to produce sodium phenoxide and H2(g) 

in NaOH(aq) with diazonium salts, to give azo compounds

nitration of the aromatic ring with dilute HNO3(aq) at room temperature to give a mixture of 2-nitrophenol and 4-nitrophenol

bromination of the aromatic ring with Br2(aq) to form 2,4,6-tribromophenol

reaction of an alkylbenzene with hot alkaline KMnO4 and then dilute acid, exemplified by methylbenzene

the oxidation of methanoic acid, HCOOH, with Fehling’s reagent or Tollens’ reagent or acidified KMnO4 or acidified K2Cr2O7 to carbon dioxide and water 

the oxidation of ethanedioic acid, HOOCCOOH, with warm acidified KMnO4 to carbon dioxide

reaction of alcohols with acyl chlorides using the formation of ethyl ethanoate and phenyl benzoate as examples 

reaction of carboxylic acids with PCl 3 and heat, PCl 5, or SOCl 2

hydrolysis on addition of water at room temperature to give the carboxylic acid and HCl

reaction with an alcohol at room temperature to produce an ester and HCl

reaction with phenol at room temperature to produce an ester and HCl

reaction with ammonia at room temperature to produce an amide and HCl 

reaction with a primary or secondary amine at room temperature to produce an amide and HCl 

reaction of halogenoalkanes with NH3 in ethanol heated under pressure

reaction of halogenoalkanes with primary amines in ethanol, heated in a sealed tube /under pressure

the reduction of amides with LiAlH4

the reduction of nitriles with LiAlH4 or H2 /Ni

the reaction of phenylamine with Br2(aq) at room temperature

the reaction of phenylamine with HNO2 or NaNO2 and dilute acid below 10°C to produce the diazonium salt; further warming of the diazonium salt with H2O to give phenol

describe the coupling of benzenediazonium chloride with phenol in NaOH(aq) to form an azo compound

identify the azo group

state that azo compounds are often used as dyes

that other azo dyes can be formed via a similar route

the reaction between ammonia and an acyl chloride at room temperature

the reaction between a primary amine and an acyl chloride at room temperature 

hydrolysis with aqueous alkali or aqueous acid

the reduction of the CO group in amides with LiAlH4 to form an amine

the reaction between a diol and a dicarboxylic acid or dioyl chloride  

the reaction of a hydroxycarboxylic acid

the reaction between a diamine and a dicarboxylic acid or dioyl chloride

the reaction of an aminocarboxylic acid 

the reaction between amino acids

identify organic functional groups using the reactions in the syllabus

predict properties and reactions

stationary phase, for example aluminium oxide (on a solid support)

mobile phase; a polar or non-polar solvent  

Rf value 

solvent front and baseline 

stationary phase; a high boiling point non-polar liquid (on a solid support)

mobile phase; an unreactive gas 

retention time 

the different environments of the carbon atoms present

the possible structures for the molecule

the different environments of proton present using chemical shift values

the relative numbers of each type of proton present from relative peak areas

the number of equivalent protons on the carbon atom adjacent to the one to which the given proton is attached from the splitting pattern, using the n + 1 rule (limited to singlet, doublet, triplet, quartet and multiplet)

the possible structures for the molecule