AQA Level 3 Advanced GCE in Chemistry
Official qualification regulated by Ofqual, awarded by AQA. Classified as a GCE A Level.
Course Syllabus & Study Modules
Complete breakdown of verified curriculum topics and learning objectives.
Action of anticancer drugs (A-level only)
Appreciate that society needs to assess the balance between the benefits and the adverse effects of drugs, such as the anticancer drug cisplatin.
- Explain why an effective anticancer drug can still cause serious harm, using the non-selective binding of platinum.
- Name two healthy tissues damaged by the drug and state the adverse effect each produces.
- Given survival figures for a cancer treated with the drug and the incidence of a named side effect, argue for or against its use and quote numbers from the data in support.
- Suggest one way of reducing adverse effects and explain how it would work.
Action of anticancer drugs (A-level only)
Cisplatin prevents DNA replication in cancer cells by a ligand replacement reaction with DNA in which a bond is formed between platinum and a nitrogen atom on guanine.
- Write an equation for the reaction of the drug with two guanine bases, showing the two chloride ions released.
- Explain why a lone pair on a nitrogen atom of guanine is able to bond to platinum.
- Explain why the adjacent arrangement of the two chloride ligands is essential to the drug's action.
- Link the formation of the platinum to guanine bond to the failure of DNA replication and of cell division.
Action of anticancer drugs (A-level only)
Students should be able to: explain why cisplatin prevents DNA replication explain why such drugs can have adverse effects.
- Sequence the steps from the drug entering a cell to replication being blocked.
- Explain each of those steps in terms of bonding rather than description.
- Explain adverse effects by applying the same binding chemistry to healthy tissue.
- Suggest why the dose given to a patient must be controlled carefully, using the fact that platinum binds guanine in healthy cells as readily as in cancer cells.
Action of anticancer drugs (A-level only)
The Pt(II) complex cisplatin is used as an anticancer drug.
- Draw the structure of the drug, showing square planar geometry with the chloride ligands adjacent.
- Deduce the oxidation state of platinum in the complex from the charges on its ligands.
- Distinguish the drug from its trans isomer in a diagram and state which one is used medically.
- State the coordination number of platinum here and justify it from the number of coordinate bonds.
Acylation (A-level only)
Required practical 10 Preparation of: a pure organic solid and test of its purity a pure organic liquid.
- Describe the full recrystallisation procedure for an impure organic solid and give the reason for each step.
- Interpret a measured melting point to decide whether a sample is pure, referring to both the value and the range.
- Justify each stage used to purify an organic liquid, naming the species removed at each stage.
Acylation (A-level only)
Students should be able to outline the mechanism of nucleophilic addition–elimination reactions of acyl chlorides with water, alcohols, ammonia and primary amines.
- Draw the full mechanism for ethanoyl chloride with water, an alcohol, ammonia or a primary amine, arrow by arrow.
- Draw the tetrahedral intermediate for a given acyl chloride reaction, with both formal charges in place.
- Explain why this mechanism is named addition-elimination rather than substitution.
Acylation (A-level only)
The industrial advantages of ethanoic anhydride over ethanoyl chloride in the manufacture of the drug aspirin.
- Give three properties of ethanoic anhydride - its cost, how readily moisture hydrolyses it, and the by-product it releases - and state the consequence of each for a factory making aspirin.
- Identify which functional group of 2-hydroxybenzoic acid reacts when aspirin is made, and draw the product.
- Compare the two routes on atom economy and explain which one wins on that measure.
Acylation (A-level only)
The nucleophilic addition–elimination reactions of water, alcohols, ammonia and primary amines with acyl chlorides and acid anhydrides.
- Predict and draw the organic product of an acyl chloride or anhydride with water, an alcohol, ammonia or a primary amine.
- Write balanced equations that account for both moles of ammonia or amine used per mole of acyl chloride.
- State the observation expected when ethanoyl chloride is added to water and explain what causes it.
Acylation (A-level only)
The structures of: acid anhydrides acyl chlorides amides.
- Draw and name an acyl chloride, an acid anhydride and a primary amide given the parent carboxylic acid.
- Deduce the parent acid from the structure of an acyl chloride, an anhydride or an amide.
- Place acyl chlorides, anhydrides, esters and amides in order of reactivity towards nucleophiles.
Addition polymers
Addition polymers are formed from alkenes and substituted alkenes.
- Write a balanced polymerisation equation from a given monomer, using n on both sides.
- Draw the polymer formed from a substituted alkene such as chloroethene or phenylethene.
- Explain why addition polymerisation has an atom economy of one hundred per cent.
Addition polymers
Addition polymers are unreactive.
- Explain the unreactivity of poly(ethene) by reference to bond strength and the absence of polar sites.
- Describe the intermolecular forces between polyalkene chains and relate them to chain length.
- Explain why an unreactive addition polymer creates a waste disposal problem.
Addition polymers
Appreciate that knowledge and understanding of the production and properties of polymers has developed over time.
- Explain how catalyst development changed the structure and properties of poly(ethene).
- Relate branching in a polymer chain to its density, rigidity and melting point.
- Discuss one recent development in polymer science and the problem it was designed to solve.
Addition polymers
IUPAC rules for naming addition polymers.
- Name an addition polymer from its monomer using the poly(monomer) convention.
- Convert common trade names such as PVC into the correct IUPAC polymer name.
- Deduce and name the monomer from a drawn repeating unit, then name the polymer.
Addition polymers
The repeating unit of addition polymers.
- Draw the repeating unit from a given monomer structure, with brackets and continuation bonds.
- Deduce the monomer from a section of an addition polymer chain by restoring the double bond.
- Check a proposed repeating unit by reproducing a longer section of the chain from it.
Addition polymers
Typical uses of poly(chloroethene), commonly known as PVC, and how its properties can be modified using a plasticiser.
- Explain, in terms of intermolecular forces, how a plasticiser makes poly(chloroethene) flexible.
- Give a use for rigid and for plasticised poly(chloroethene) and link each to the property involved.
- Identify the intermolecular forces between poly(chloroethene) chains and justify the choice.
Addition reactions of alkenes
Students should be able to: outline the mechanisms for these reactions explain the formation of major and minor products by reference to the relative stabilities of primary, secondary and tertiary carbocation intermediates.
- Draw the full electrophilic addition mechanism for an alkene reacting with hydrogen bromide or with sulfuric acid.
- Explain which carbocation forms preferentially and use it to predict the major product.
- Check a drawn mechanism against the arrow conventions and identify any arrow that would gain no credit.
Addition reactions of alkenes
The formation of major and minor products in addition reactions of unsymmetrical alkenes.
- Draw both carbocation intermediates possible when a named unsymmetrical alkene reacts with hydrogen bromide.
- Predict the major product and justify the choice by the relative stability of the intermediates.
- Explain, using the inductive effect, why a tertiary carbocation is more stable than a primary carbocation.
Alcohol production
Alcohols are produced industrially by hydration of alkenes in the presence of an acid catalyst.
- State the reagent, catalyst and conditions for the industrial hydration of ethene and write the equation.
- Predict the major alcohol formed when an unsymmetrical alkene is hydrated.
- Calculate the mass of alcohol obtained from a given volume of alkene at a stated percentage yield.
Alcohol production
Ethanol is produced industrially by fermentation of glucose. The conditions for this process.
- Write the balanced equation for the fermentation of glucose to ethanol.
- Justify the temperature and the exclusion of air used in fermentation, giving the consequence of departing from each.
- Compare fermentation with hydration of ethene in terms of rate, purity of product and raw material.
Alcohol production
Ethanol produced industrially by fermentation is separated by fractional distillation and can then be used as a biofuel.
- Explain how fractional distillation separates ethanol from the aqueous mixture produced by fermentation.
- Label a distillation apparatus correctly, including the direction of flow of condenser water.
- Define the term biofuel and classify ethanol from fermentation against ethanol from ethene.
Alcohol production
Students should be able to: explain the meaning of the term biofuel justify the conditions used in the production of ethanol by fermentation of glucose write equations to support the statement that ethanol produced by fermentation is a carbon neutral fuel and give reasons why this statement is not valid outline the mechanism for the formation of an alcohol by the reaction of an alkene with steam in the presence of an acid catalyst discuss the environmental (including ethical) issues linked to decision making about biofuel use.
- Explain the meaning of the term biofuel and justify the conditions used for fermentation.
- Write and combine the photosynthesis, fermentation and combustion equations to test the carbon neutral claim.
- Outline the mechanism for the acid-catalysed hydration of an alkene, showing the catalyst regenerated.
- Discuss the environmental and ethical issues linked to decision making about biofuel use, such as land use and food supply.
Aldehydes and ketones (A-level only)
Aldehydes and unsymmetrical ketones form mixtures of enantiomers when they react with KCN followed by dilute acid.
- Explain why nucleophilic addition of cyanide to most aldehydes produces equal amounts of two enantiomers.
- Deduce whether the product of a given carbonyl compound with cyanide will be chiral.
- Explain why methanal and propanone give products that show no optical isomerism.
Aldehydes and ketones (A-level only)
Aldehydes are readily oxidised to carboxylic acids.
- Write the equation using [O] for the oxidation of a named aldehyde to its carboxylic acid.
- Explain why aldehydes are readily oxidised but ketones are not, referring to the carbonyl carbon.
- State the reagent, conditions and colour change for oxidising an aldehyde with dichromate.
Aldehydes and ketones (A-level only)
Chemical tests to distinguish between aldehydes and ketones including Fehling’s solution and Tollens’ reagent.
- Describe how to carry out the Tollens' and Fehling's tests, including how the reagent is prepared and warmed.
- State the observation with an aldehyde and with a ketone for each test.
- Explain each test as a redox reaction, naming the species oxidised and the species reduced.
Aldehydes and ketones (A-level only)
Students should be able to: write overall equations for reduction reactions using [H] as the reductant outline the nucleophilic addition mechanism for reduction reactions with NaBH (the nucleophile – should be shown as H ) write overall equations for the formation of hydroxynitriles using HCN outline the nucleophilic addition mechanism for the reaction with KCN followed by dilute acid explain why nucleophilic addition reactions of KCN, followed by dilute acid, can produce a mixture of enantiomers.
- Write balanced equations using 2[H] for the reduction of a named aldehyde and a named ketone.
- Draw the nucleophilic addition mechanism with H⁻ and with CN⁻, including the final protonation.
- Explain why the cyanide reaction gives a racemate while the borohydride reduction of propanone does not give a chiral product.
Aldehydes and ketones (A-level only)
The hazards of using KCN.
- State two specific hazards of using potassium cyanide and the precaution that controls each.
- Explain why adding dilute acid to a cyanide mixture is the most dangerous step in the preparation.
- Justify why a fume cupboard rather than eye protection alone is required for this reaction.
Aldehydes and ketones (A-level only)
The nucleophilic addition reactions of carbonyl compounds with KCN, followed by dilute acid, to produce hydroxynitriles.
- Draw the nucleophilic addition mechanism for the reaction of a carbonyl compound with cyanide ion and dilute acid.
- Identify the reagent and conditions needed to convert a named aldehyde into a hydroxynitrile.
- Name the hydroxynitrile formed from a given aldehyde or ketone using correct locants.
Aldehydes and ketones (A-level only)
These reduction reactions are examples of nucleophilic addition.
- Name the mechanism by which NaBH₄ reduces a carbonyl compound and justify both words in the name.
- Explain why the carbonyl carbon is the site attacked, referring to electronegativity.
- Predict the class of alcohol formed when a given aldehyde or ketone is reduced.
Amino acids (A-level only)
Amino acids have both acidic and basic properties, including the formation of zwitterions.
- Draw the zwitterion of a named amino acid and show that the charges cancel.
- Explain why amino acids have unexpectedly high melting points and are soluble in water.
- Predict what an amino acid does when hydrochloric acid or sodium hydroxide solution is added.
Amino acids (A-level only)
Students should be able to draw the structures of amino acids as zwitterions and the ions formed from amino acids: in acid solution in alkaline solution.
- Draw glycine as a zwitterion, as its cation and as its anion, labelling the overall charge on each and showing three hydrogens on the nitrogen in the zwitterion and the cation, two in the anion.
- Deduce the overall charge on an amino acid at a stated pH relative to its isoelectric point.
- Explain which group gains or loses a proton when acid or alkali is added, and why.
Applications of Hess’s law
Students should be able to: use Hess’s law to perform calculations, including calculation of enthalpy changes for reactions from enthalpies of combustion or from enthalpies of formation.
- Construct a labelled Hess cycle linking a target reaction to a set of given enthalpies of formation or combustion.
- Calculate an enthalpy change from formation data using Σ∆fH(products) − Σ∆fH(reactants), including the balancing numbers.
- Calculate an enthalpy change from combustion data and explain why the subtraction is the other way round.
Balanced equations and associated calculations
Equations (full and ionic).
- Write the ionic equation for a precipitation or neutralisation reaction by cancelling the spectator ions from the full equation.
- Check an ionic equation by confirming that both atoms and total charge balance.
- Decide which species in a full equation may be written as separate ions and justify the choice.
Balanced equations and associated calculations
Required practical 1 Make up a volumetric solution and carry out a simple acid–base titration.
- Describe, in order, the steps needed to make up 250 cm³ of a solution of accurately known concentration from a weighed solid.
- Explain why the beaker washings must be transferred to the volumetric flask.
- Select concordant titres from a results table and calculate the mean titre from them alone.
Balanced equations and associated calculations
Students should be able to: write balanced equations for reactions studied balance equations for unfamiliar reactions when reactants and products are specified.
- Balance an unfamiliar ionic equation for both atoms and total charge when the reactants and products are supplied.
- Construct a half-equation by adding water, hydrogen ions and electrons in the correct order.
- Explain why balancing must be done by changing coefficients and never by changing a formula.
Base properties (A-level only)
Amines are weak bases.
- Write the equilibrium equation for a named amine in water and state where the equilibrium lies.
- Explain the difference between a weak base and a dilute base using an amine as the example.
- Explain why a titration of a carboxylic acid with an amine cannot be followed using an indicator.
Base properties (A-level only)
The difference in base strength between ammonia, primary aliphatic and primary aromatic amines.
- Name the effect the benzene ring has on the nitrogen lone pair in phenylamine and the opposite effect the alkyl group has in ethylamine, and use the two to place each amine either side of ammonia.
- Explain the position of each using the availability of the nitrogen lone pair.
- Predict which of two given amines is the stronger base and justify the prediction.
Biodegradability and disposal of polymers (A-level only)
Polyalkenes are chemically inert and non-biodegradable.
- Explain, in terms of bond polarity, why a polyalkene chain resists attack by water, acid and alkali.
- State one advantage and one disadvantage that follow from a polyalkene being chemically inert.
- Identify from a structure whether a polymer chain contains a site open to nucleophilic attack.
Biodegradability and disposal of polymers (A-level only)
Polyesters and polyamides can be broken down by hydrolysis and are biodegradable.
- Draw the products formed when a given polyester or polyamide is hydrolysed in acid and in alkali.
- Explain why the carbonyl carbon of the link is the site attacked during hydrolysis.
- Account for the biodegradability of a condensation polymer in terms of the bonds in its backbone.
Biodegradability and disposal of polymers (A-level only)
Students should be able to: explain why polyesters and polyamides can be hydrolysed but polyalkenes cannot.
- Give the products of hydrolysing a named polyamide under acid conditions, and state what happens to poly(ethene) treated the same way.
- Label the partial charges on the carbonyl carbon and the neighbouring oxygen of a polyester chain.
- Predict from an unfamiliar polymer structure whether it will be biodegradable, and justify the prediction.
Biodegradability and disposal of polymers (A-level only)
The advantages and disadvantages of different methods of disposal of polymers, including recycling.
- Give one gain and one cost for landfill, for incineration and for recycling, each one naming a substance, a resource or a practical difficulty rather than saying it harms the environment.
- Explain how the hydrogen chloride produced by burning PVC is dealt with, and write an equation for it.
- Evaluate which disposal route is most appropriate for a given waste stream and justify the choice.
Bond enthalpies
Mean bond enthalpy.
- Define mean bond enthalpy in a form that includes one mole of bonds, the gaseous state and the averaging over different compounds.
- Explain why every mean bond enthalpy is positive by reference to what happens to the bond.
- Justify why the O−H mean bond enthalpy does not exactly describe the O−H bond in any one named molecule.
- Write an equation, with state symbols, whose enthalpy change is the bond enthalpy of a stated diatomic bond.
Bond enthalpies
Students should be able to: define the term mean bond enthalpy use mean bond enthalpies to calculate an approximate value of ∆H for reactions in the gaseous phase explain why values from mean bond enthalpy calculations differ from those determined using Hess’s law.
- Define mean bond enthalpy as an average value for breaking one mole of a bond in the gaseous state.
- Calculate an approximate ∆H for a gas phase reaction from mean bond enthalpies, counting each bond type from displayed formulae.
- Explain two distinct reasons why a mean bond enthalpy result differs from a value obtained using Hess's law.
Bond polarity
Electronegativity as the power of an atom to attract the pair of electrons in a covalent bond.
- Define electronegativity in terms of the bonding pair in a covalent bond.
- Explain why electronegativity increases from sodium to chlorine but decreases from fluorine to iodine.
- Rank a set of atoms by electronegativity using supplied data and justify the order.
Bond polarity
Students should be able to: use partial charges to show that a bond is polar explain why some molecules with polar bonds do not have a permanent dipole.
- Annotate a displayed structure with δ⁺ and δ⁻ and justify each label using electronegativity.
- Explain why tetrachloromethane has four polar bonds but no permanent dipole.
- Rank several molecules by polarity using both their shapes and the electronegativities of the atoms involved.
Bond polarity
The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical.
- Deduce which atom in a given bond carries the partial negative charge and justify the choice with electronegativity data.
- Explain why a partial charge is written as δ⁺ or δ⁻ rather than as a whole charge.
- Identify which bonds in a molecule are polar and which are non-polar.
Bond polarity
This produces a polar covalent bond, and may cause a molecule to have a permanent dipole.
- Predict whether a molecule has a permanent dipole by combining its bond polarities with its shape.
- Explain why carbon dioxide has polar bonds but no permanent dipole while water has both.
- Identify the structural feature that prevents bond dipoles from cancelling in a named molecule.
Bonding (A-level only)
Delocalisation of p electrons makes benzene more stable than the theoretical molecule cyclohexa-1,3,5-triene.
- Calculate the delocalisation energy of benzene from the enthalpies of hydrogenation of benzene and cyclohexene.
- Explain why the theoretical value is three times the cyclohexene figure and why the same product must be used.
- Draw and label an energy level diagram showing benzene, cyclohexa-1,3,5-triene and cyclohexane.
Bonding (A-level only)
Students should be able to: use thermochemical evidence from enthalpies of hydrogenation to account for this extra stability explain why substitution reactions occur in preference to addition reactions.
- Decide from two enthalpies of hydrogenation which compound is the more stable, explaining why the less exothermic value belongs to the more stable compound.
- State what an addition reaction does to the delocalised ring and what a substitution restores, then use the delocalisation energy of about 150 kJ mol⁻¹ to say which product is the more stable.
- Predict the result of adding bromine water to benzene and to cyclohexene, and account for the difference.
Bonding (A-level only)
The nature of the bonding in a benzene ring, limited to planar structure and bond length intermediate between single and double.
- Describe the shape and bond angles of benzene and state the evidence that all six carbon-carbon bonds are identical.
- Explain how the p orbitals on the six carbon atoms give rise to a delocalised ring of electrons.
- Compare the carbon-carbon bond length in benzene with those in ethane and ethene, and account for the value.
Bonding and physical properties
The four types of crystal structure: ionic metallic macromolecular (giant covalent) molecular.
- Classify a named solid into one of the four crystal types from data on its melting point and conductivity.
- Name the force overcome when each of the four structure types melts.
- Explain why melting a molecular solid does not break any covalent bonds.
Bonding and physical properties
The structures of the following crystals as examples of these four types of crystal structure: diamond graphite ice iodine magnesium sodium chloride.
- Explain why graphite conducts electricity and is soft while diamond does neither, using the bonding in each.
- Describe the arrangement of particles in the sodium chloride lattice and in a lattice of iodine molecules.
- Predict the relative melting points of magnesium, iodine and diamond and justify the order.
Born–Haber cycles (A-level only)
Born–Haber cycles are used to calculate lattice enthalpies using the following data: enthalpy of formation ionisation energy enthalpy of atomisation bond enthalpy electron affinity.
- Construct a Born–Haber cycle for a named ionic compound with every level labelled with formulae, charges and state symbols.
- Calculate a lattice enthalpy from tabulated data by summing the indirect route and reversing the sign of every step travelled against its arrow, then check that lattice formation comes out negative and lattice dissociation positive.
- Select the correct number of moles for each step from the formula of the compound.
- Explain why the second electron affinity of oxygen is endothermic while the first is exothermic.
Born–Haber cycles (A-level only)
Cycles are used to calculate enthalpies of solution for ionic compounds from lattice enthalpies and enthalpies of hydration.
- Calculate an enthalpy of solution from a lattice enthalpy and the hydration enthalpies of each ion present.
- Define enthalpy of hydration with reference to one mole of gaseous ions and an aqueous solution.
- Deduce from a calculated enthalpy of solution whether the temperature of the water rises or falls.
- Explain why the hydration enthalpy of a small, highly charged ion is more exothermic than that of a large singly charged ion.
Born–Haber cycles (A-level only)
Lattice enthalpy can be defined as either enthalpy of lattice dissociation or enthalpy of lattice formation.
- Define lattice dissociation enthalpy and lattice formation enthalpy so that the two differ only in direction and sign.
- Write an equation with state symbols corresponding to the lattice enthalpy of a stated ionic compound.
- Deduce the sign of a lattice enthalpy from the direction of the equation supplied.
- Explain why a definition referring to one mole of gaseous ions does not describe lattice enthalpy.
Born–Haber cycles (A-level only)
Students should be able to: define each of the above terms and lattice enthalpy construct Born–Haber cycles to calculate lattice enthalpies using these enthalpy changes construct Born–Haber cycles to calculate one of the other enthalpy changes compare lattice enthalpies from Born–Haber cycles with those from calculations based on a perfect ionic model to provide evidence for covalent character in ionic compounds.
- Define lattice enthalpy, enthalpy of formation, atomisation, ionisation energy, and electron affinity.
- Construct Born-Haber cycles to calculate lattice enthalpies or other enthalpy changes.
- Compare experimental and theoretical lattice enthalpies to deduce covalent character.
Brønsted–Lowry acid–base equilibria in aqueous solution (A-level only)
A base is a proton acceptor.
- Identify the proton acceptor in NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ and mark the lone pair involved.
- Write the conjugate acid of OH⁻, of NH₃ and of CH₃COO⁻.
- Explain why water is described as amphoteric, using one equation in which it donates a proton and one in which it accepts one.
Brønsted–Lowry acid–base equilibria in aqueous solution (A-level only)
Acid–base equilibria involve the transfer of protons.
- Label both conjugate acid and base pairs in CH₃COOH + NH₃ ⇌ CH₃COO⁻ + NH₄⁺.
- Predict which side the equilibrium HCl + CH₃COO⁻ ⇌ Cl⁻ + CH₃COOH lies on and justify it from acid strength.
- Explain, using an equation, why the conjugate base of a strong acid is a very weak base.
Brønsted–Lowry acid–base equilibria in aqueous solution (A-level only)
An acid is a proton donor.
- Identify the proton donor in the equilibrium CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻.
- Write the conjugate base of HNO₃, of H₂SO₄ and of NH₄⁺.
- Distinguish a strong acid from a concentrated acid using the extent of dissociation.
Buffer action (A-level only)
A buffer solution maintains an approximately constant pH, despite dilution or addition of small amounts of acid or base.
- Write the two equations showing how an ethanoic acid and ethanoate buffer removes added H⁺ and added OH⁻.
- Explain why diluting a buffer with water leaves its pH almost unchanged.
- Predict what happens to the pH when acid is added beyond the capacity of the buffer.
Buffer action (A-level only)
Acidic buffer solutions contain a weak acid and the salt of that weak acid.
- Describe two ways of preparing an ethanoic acid and ethanoate buffer from reagents available in a school laboratory.
- Write the equation for the partial neutralisation of ethanoic acid by potassium hydroxide.
- Identify which species in an acidic buffer removes added H⁺ and write the equation for that reaction.
Buffer action (A-level only)
Applications of buffer solutions.
- Write the equilibrium that buffers blood plasma and state the pH range it maintains.
- Explain, with an equation, how the hydrogencarbonate ion removes acid produced by exercising muscle.
- Suggest why a buffer is added to a shampoo and state the approximate pH chosen.
Buffer action (A-level only)
Basic buffer solutions contain a weak base and the salt of that weak base.
- Select, from a list of pairs of solutions, the one that forms a basic buffer and justify the choice.
- Write the two equations showing how an ammonia and ammonium chloride buffer resists added acid and added alkali.
- Explain why mixing equal volumes of hydrochloric acid and sodium hydroxide of equal concentration does not give a buffer.
Buffer action (A-level only)
Students should be able to: explain qualitatively the action of acidic and basic buffers calculate the pH of acidic buffer solutions.
- Calculate the pH of the buffer formed when 20.0 cm³ of 0.100 mol dm⁻³ sodium hydroxide is added to 25.0 cm³ of 0.100 mol dm⁻³ butanoic acid.
- Explain why the volumes cancel in the buffer expression, so that moles may be used in place of concentrations.
- State which species removes added acid and which removes added alkali in a butanoic acid and butanoate buffer, giving an equation for each.
Calorimetry
Required practical 2 Measurement of an enthalpy change.
- Describe a procedure that measures the enthalpy change of a reaction in an insulated cup.
- Determine a corrected temperature change by extrapolating the cooling line of a temperature against time graph back to the time of mixing.
- Calculate the percentage uncertainty in ∆T from the resolution of the thermometer.
Calorimetry
Students should be able to: use this equation to calculate the molar enthalpy change for a reaction use this equation in related calculations.
- Calculate a molar enthalpy change from raw calorimetry data by combining q = mc∆T with the amount in moles of the limiting reactant.
- Rearrange q = mc∆T to find an unknown mass, temperature change or specific heat capacity from a quoted enthalpy change.
- Calculate an enthalpy of vaporisation from the electrical energy supplied by a heater of stated power and the mass of liquid lost.
- Identify the limiting reactant in a calorimetry experiment and justify the choice before dividing.
Calorimetry
The heat change, q, in a reaction is given by the equation q = mc∆T where m is the mass of the substance that has a temperature change ∆T and a specific heat capacity c.
- Identify which substance's mass and specific heat capacity belong in q = mc∆T for a dissolving, neutralisation or displacement experiment.
- Calculate q in joules from a recorded volume, temperature change and given specific heat capacity, then express it in kilojoules.
- Justify the sign given to ∆H by stating the direction of heat transfer between system and surroundings.
- Suggest a change of method that increases ∆T and explain why this lowers the percentage uncertainty in the temperature reading.
Carboxylic acids and esters (A-level only)
Biodiesel is a mixture of methyl esters of long-chain carboxylic acids.
- Deduce the long-chain carboxylic acid a given biodiesel molecule came from by splitting it at the ester linkage, and confirm the alcohol was methanol.
- Explain why biodiesel from one vegetable oil is a mixture of esters rather than a single compound.
- Evaluate the claim that biodiesel is carbon neutral, giving one point on each side.
Carboxylic acids and esters (A-level only)
Biodiesel is produced by reacting vegetable oils with methanol in the presence of a catalyst.
- Write a balanced equation for the production of biodiesel from a triglyceride and methanol.
- Explain, in terms of equilibrium, why a large excess of methanol is used and why the glycerol is removed.
- Identify the catalyst used and state what happens to it by the end of the process.
Carboxylic acids and esters (A-level only)
Carboxylic acids and alcohols react, in the presence of an acid catalyst, to give esters.
- Write a balanced equation for a named alcohol reacting with a named carboxylic acid, including water and the ⇌ sign.
- Justify each purification step used on the crude ester, naming the species removed at each stage.
- Calculate a percentage yield of an ester by identifying the limiting reagent from the masses supplied.
- Explain why distilling the ester out of the mixture increases the amount of ester obtained.
Carboxylic acids and esters (A-level only)
Common uses of esters (eg in solvents, plasticisers, perfumes and food flavourings).
- Explain the volatility of a short-chain ester in terms of the intermolecular forces its molecules can and cannot form.
- Describe how a plasticiser changes the properties of a polymer, at the level of the forces between chains.
- Explain why an ester is used as a solvent, referring to the polar C=O group and the non-polar alkyl chains that flank it.
Carboxylic acids and esters (A-level only)
Esters can be hydrolysed in acid or alkaline conditions to form alcohols and carboxylic acids or salts of carboxylic acids.
- Predict and name the organic products of hydrolysing a given ester under acidic and under alkaline conditions.
- Explain why alkaline hydrolysis goes to completion while acid hydrolysis reaches an equilibrium.
- Describe how the free carboxylic acid is recovered after alkaline hydrolysis.
Carboxylic acids and esters (A-level only)
The structures of: carboxylic acids esters.
- Draw displayed formulae for a named carboxylic acid and a named ester, showing where the C=O, the O-H and the C-O-C sit.
- Name an ester from its structure and work backwards to the alcohol and the acid that would make it.
- Explain why an ester boils at a lower temperature than the carboxylic acid of similar relative molecular mass.
- Identify a carboxylic acid in an unknown by the reagent used and the observation seen with a carbonate.
Carboxylic acids and esters (A-level only)
Vegetable oils and animal fats are esters of propane-1,2,3-triol (glycerol).
- Draw the structure of a triglyceride from glycerol and a named long-chain carboxylic acid, showing all three linkages.
- Explain why a saturated fat is a solid and an unsaturated oil a liquid at the same temperature.
- Deduce the number of C=C bonds per molecule of an oil from the moles of bromine that react with it.
Carboxylic acids and esters (A-level only)
Vegetable oils and animal fats can be hydrolysed in alkaline conditions to give soap (salts of long-chain carboxylic acids) and glycerol.
- Write a balanced equation for the saponification of a named triglyceride, with the correct stoichiometry of alkali.
- Draw the structure of a soap molecule and label the part attracted to water and the part attracted to grease.
- Explain how soap removes grease from a surface in terms of the two ends of the molecule.
Catalysts
A catalyst is a substance that increases the rate of a chemical reaction without being changed in chemical composition or amount.
- Decide whether a named species in a described reaction is acting as a catalyst, a reactant or an intermediate, justifying the choice from whether it is present at the start, at the end, or only part way through.
- Classify a given catalysed reaction as heterogeneous or homogeneous and justify the classification.
- Explain how a catalyst can take part in a reaction and still be present unchanged at the end.
- Identify autocatalysis from a concentration against time curve that steepens partway through the reaction, or from a rate against time curve that rises to a maximum before falling.
Catalysts
Catalysts work by providing an alternative reaction route of lower activation energy.
- Explain how a catalyst increases rate by reference to an alternative route with a lower activation energy.
- Sketch an enthalpy profile showing catalysed and uncatalysed routes with the same reactants, products and ∆H.
- Describe the steps of heterogeneous catalysis, from adsorption at active sites to desorption of products.
- Explain why a reaction between two ions of like charge has a high activation energy and how a catalyst of opposite charge overcomes it.
Catalysts
Students should be able to: use a Maxwell–Boltzmann distribution to help explain how a catalyst increases the rate of a reaction involving a gas.
- Sketch one Maxwell–Boltzmann distribution marked with both catalysed and uncatalysed activation energies.
- Explain a catalytic rate increase using the change in the area beyond the activation energy on an unchanged curve.
- Distinguish the diagram required for a catalyst from that required for a temperature rise, and justify the difference.
- Deduce what happens to the mean molecular energy and the collision frequency when a catalyst is added.
Catalysts (A-level only)
A heterogeneous catalyst is in a different phase from the reactants and the reaction occurs at active sites on the surface.
- State the meaning of heterogeneous in terms of the phases of the catalyst and the reactants.
- Describe the adsorption, reaction and desorption stages that occur at an active site.
- Explain why adsorption that is too strong or too weak reduces a catalyst's effectiveness.
Catalysts (A-level only)
A homogeneous catalyst is in the same phase as the reactants.
- Classify a catalysed reaction as homogeneous from the state symbols in its equation, as when Mn²⁺(aq) catalyses MnO₄⁻(aq) with C₂O₄²⁻(aq).
- Name a homogeneous catalyst for a specified reaction and write the species involved.
- Define an autocatalyst as a product that catalyses its own reaction and give an example.
Catalysts (A-level only)
Fe is used as a heterogeneous catalyst in the Haber process.
- Write the equilibrium equation for the Haber process and state the conditions used.
- Explain why the iron catalyst changes the rate but not the equilibrium yield of ammonia.
- Justify the classification of this catalysis as heterogeneous by naming the phases present.
Catalysts (A-level only)
Heterogeneous catalysts can become poisoned by impurities that block the active sites and consequently have reduced efficiency; this has a cost implication.
- Explain catalyst poisoning in terms of impurities blocking active sites on the surface.
- Name a poison for a specified industrial catalyst and state the precaution taken against it.
- Describe the economic consequences of poisoning for an industrial process.
Catalysts (A-level only)
The use of a support medium to maximise the surface area of a heterogeneous catalyst and minimise the cost.
- Explain how a support medium increases the number of active sites available per gram of catalyst.
- State how using a support reduces the cost of a catalyst and give an example of one in use.
- Distinguish the role of the catalyst from that of the inert support in a catalytic converter.
Catalysts (A-level only)
Transition metals and their compounds can act as heterogeneous and homogeneous catalysts.
- Define a catalyst in terms of an alternative route of lower activation energy and being chemically unchanged.
- Classify named industrial and laboratory catalysts as heterogeneous or homogeneous, stating the phases involved.
- Explain why variable oxidation states make transition metals effective catalysts.
Catalysts (A-level only)
When catalysts and reactants are in the same phase, the reaction proceeds through an intermediate species.
- Write the two step equations for a homogeneous catalytic mechanism, showing the intermediate and the regenerated catalyst.
- Explain why an uncatalysed reaction between two anions is slow and why the catalysed route is faster.
- Interpret a rate curve that steepens as a reaction proceeds as evidence of autocatalysis.
Chemical equilibria and Le Chatelier's principle
A catalyst does not affect the position of equilibrium.
- Explain why speeding up both the forward and reverse reactions equally leaves the position of equilibrium unchanged.
- State the effect of a catalyst on the equilibrium yield, on the equilibrium constant and on the time taken to reach equilibrium.
- Identify the incorrect statement in a set of claims about catalysts in equilibrium systems and justify the choice.
- Explain why a catalyst allows an industrial equilibrium process to run at a lower temperature.
Chemical equilibria and Le Chatelier's principle
In a reversible reaction at equilibrium: forward and reverse reactions proceed at equal rates the concentrations of reactants and products remain constant
- Deduce from a graph of concentrations against time whether a mixture has reached equilibrium, and identify the time at which it did so.
- Explain why constant concentrations do not imply equal concentrations of reactants and products.
- Describe evidence that the reactions continue after equilibrium has been reached.
- Explain why a system open to the atmosphere cannot reach equilibrium when a gaseous product is formed.
Chemical equilibria and Le Chatelier's principle
Le Chatelier's principle can be used to predict the effects of changes in temperature, pressure and concentration on the position of equilibrium in homogeneous reactions.
- Predict the direction of shift and the effect on yield when the temperature of a named equilibrium is raised or lowered.
- Deduce the sign of ∆H for a forward reaction from reported changes in yield with temperature.
- Predict the effect of a pressure change by comparing the number of moles of gas on each side of the equation.
- Explain the effect on the position of equilibrium of removing a product as it is formed.
Chemical equilibria and Le Chatelier's principle
Le Chatelier’s principle.
- State Le Chatelier's principle in a form that includes opposing the change imposed on the system.
- Identify the imposed change in a described scenario before predicting the direction of any shift.
- Explain why the principle predicts direction only and says nothing about how quickly equilibrium is restored.
- Distinguish between a change that alters the position of equilibrium and one that alters only the rate.
Chemical equilibria and Le Chatelier's principle
Many chemical reactions are reversible.
- Explain why the percentage yield of a reversible reaction cannot reach 100 per cent even under ideal handling.
- Describe how the forward and reverse rates change from the moment of mixing until equilibrium is reached.
- Suggest two ways of increasing the yield of a named reversible reaction and justify each.
- Distinguish between yield lost to the position of equilibrium and yield lost during isolation and purification.
Chemical equilibria and Le Chatelier's principle
Students should be able to: use Le Chatelier’s principle to predict qualitatively the effect of changes in temperature, pressure and concentration on the position of equilibrium explain why, for a reversible reaction used in an industrial process, a compromise temperature and pressure may be used.
- Predict the temperature and pressure that would maximise the equilibrium yield of a named industrial reaction.
- Explain why the conditions actually used differ from those predicted by yield alone.
- Justify the use of a catalyst in an industrial equilibrium process in terms of rate and operating temperature.
- Explain how recycling unreacted gases compensates for the yield sacrificed by a compromise condition.
Chlorination of alkanes
Students should be able to: explain this reaction as a free-radical substitution mechanism involving initiation, propagation and termination steps.
- Write the initiation, both propagation and at least two termination steps for the chlorination of methane.
- Classify an unfamiliar step as initiation, propagation or termination by counting the radicals on each side.
- Explain why a trace of ultraviolet light produces a large amount of chloromethane, referring to radical regeneration.
Chlorination of alkanes
The reaction of methane with chlorine.
- Write balanced equations for each successive chlorination of methane, from CH₃Cl through to CCl₄.
- Predict which product dominates when methane is in large excess and when chlorine is in large excess, giving a reason.
- Explain why this reaction is unsuitable for preparing a pure sample of chloromethane.
Chromatography (A-level only)
Chromatography can be used to separate and identify the components in a mixture.
- Identify the stationary phase and the mobile phase in a described experiment.
- Explain why the components of a mixture travel at different speeds through a chromatography system.
- Calculate an Rf value from a chromatogram and state why it carries no unit.
- Explain why a retention time identifies a component only when a standard is run under the same conditions.
Chromatography (A-level only)
Required practical 12 Separation of species by thin-layer chromatography.
- Describe the steps of a thin-layer chromatography run in the order they are carried out.
- Explain why the tank must be covered and why the solvent must start below the baseline.
- Calculate an Rf value from measurements taken on a developed plate.
- Identify a fault in a diagram of the apparatus and state the effect it would have on the chromatogram.
Chromatography (A-level only)
Separation depends on the balance between solubility in the moving phase and retention by the stationary phase.
- Explain separation in chromatography as a balance between solubility in the mobile phase and retention by the stationary phase.
- Distinguish between adsorption on a solid stationary phase and partition in a liquid stationary phase.
- Predict the relative Rf values or retention times of two components based on their polarities and affinities for the stationary phase.
Chromatography (A-level only)
The use of mass spectrometry to analyse the components separated by GC.
- Explain why a fragmentation pattern can tell apart two compounds that share a relative molecular mass, when the molecular ion peak alone cannot.
- Explain why a retention time on its own is not enough to identify a component.
- Use a molecular ion peak to state the relative molecular mass of a separated component.
- Suggest an application of the combined technique and justify it from a property of the method.
Chromatography (A-level only)
Types of chromatography include: thin-layer chromatography (TLC) – a plate is coated with a solid and a solvent moves up the plate column chromatography (CC) – a column is packed with a solid and a solvent moves down the column gas chromatography (GC) – a column is packed with a solid or with a solid coated by a liquid, and a gas is passed through the column under pressure at high temperature.
- Choose the technique suited to a stated sample, such as a volatile liquid mixture or a mixture of amino acids, and give the property of the sample that decides it.
- Explain why the solvent climbs a thin-layer plate but runs down a packed column, naming what drives it in each case.
- Explain why gas chromatography is carried out at high temperature and under pressure.
- Select the appropriate technique when a question needs the separated components to be collected rather than only detected.
Classification
An element is classified as s, p, d or f block according to its position in the Periodic Table, which is determined by its proton number.
- Deduce the block, period and group of an element from the configuration 1s² 2s² 2p⁶ 3s² 3p⁴.
- Explain why helium is classified as an s-block element despite being printed above the noble gases.
- Write the electron configuration of copper and identify the sub-shell that names its block.
Collision theory
Reactions can only occur when collisions take place between particles having sufficient energy.
- State the two conditions a collision must satisfy before it can lead to reaction.
- Explain why a reaction mixture in which billions of collisions occur each second may still react slowly.
- Distinguish between collision frequency and the proportion of successful collisions when accounting for a rate change.
- Apply collision theory to explain why a reaction is fastest at the start and slows as reactants are used up.
Collision theory
Students should be able to: define the term activation energy explain why most collisions do not lead to a reaction.
- Define activation energy and use the definition to explain why most collisions in a reacting mixture are unsuccessful.
- Explain, in terms of an energy distribution, why only a small fraction of particles can react at a given temperature.
- Give two distinct reasons a collision between reactant particles may fail to produce products.
- Contrast the effect on rate of changing the collision frequency with that of changing the proportion of successful collisions.
Collision theory
This energy is called the activation energy.
- Define activation energy using the word minimum and referring to colliding particles.
- Label Eₐ and ∆H correctly on an enthalpy profile for an exothermic and an endothermic reaction.
- Explain why a mixture that is thermodynamically unstable can persist unchanged until it is ignited.
- Predict which of two reactions with quoted activation energies is faster at the same temperature, and justify the answer.
Combustion of alkanes
Alkanes are used as fuels.
- Explain why alkane combustion is exothermic by comparing the bonds broken with the bonds made.
- Write a balanced equation for the complete combustion of a named alkane.
- Evaluate the use of alkanes as fuels by giving one advantage and one specific disadvantage.
Combustion of alkanes
Combustion of alkanes and other organic compounds can be complete or incomplete.
- Write balanced equations for both complete and incomplete combustion of a named alkane.
- Explain why an experimental enthalpy of combustion is less exothermic than the data book value.
- Name the products of incomplete combustion and state one hazard of each.
Combustion of alkanes
Combustion of hydrocarbons containing sulfur leads to sulfur dioxide that causes air pollution.
- Write equations for the formation of sulfur dioxide during combustion and for its reaction with water.
- Describe two specific environmental effects of acid rain.
- State one measure used to reduce sulfur dioxide emissions and explain where in the process it acts.
Combustion of alkanes
Students should be able to: explain why sulfur dioxide can be removed from flue gases using calcium oxide or calcium carbonate.
- Explain the removal of sulfur dioxide from flue gases as an acid-base reaction between an acidic oxide and a base.
- Write balanced equations for the reaction of sulfur dioxide with calcium oxide and with calcium carbonate.
- Evaluate the carbonate route by identifying the carbon dioxide released as a disadvantage.
Combustion of alkanes
These gaseous pollutants from internal combustion engines can be removed using catalytic converters.
- Name the pollutants a catalytic converter removes and state where each one comes from.
- Write balanced equations for the removal of carbon monoxide and of nitrogen monoxide in a converter.
- Explain the operation of a converter as heterogeneous catalysis on a coated honeycomb.
Commercial applications of electrochemical cells (A-level only)
Electrochemical cells can be used as a commercial source of electrical energy.
- Deduce the EMF and the polarity of a commercial cell from two given electrode potentials.
- Explain why a lithium-ion cell can be recharged whereas an alkaline zinc cell cannot.
- Describe the role of the electrolyte in allowing a cell to deliver current, without referring to electron flow through it.
Commercial applications of electrochemical cells (A-level only)
Fuel cells are used to generate an electric current and do not need to be electrically recharged.
- Explain why a hydrogen and oxygen fuel cell never needs recharging, in terms of the supply of reactants.
- Compare a hydrogen fuel cell with a petrol engine for efficiency and for what is released at the point of use.
- Suggest two practical difficulties in storing hydrogen for use in a vehicle.
Commercial applications of electrochemical cells (A-level only)
Students should be able to: use given electrode data to deduce the reactions occurring in non-rechargeable and rechargeable cells deduce the EMF of a cell explain how the electrode reactions can be used to generate an electric current.
- Deduce from two given E⦵ values which electrode of a cell is negative and write the reaction occurring there.
- Write the charging equations for a rechargeable cell as the reverse of its discharge equations.
- Explain, in terms of electrons and ions, how the electrode reactions produce a current in an external circuit.
Commercial applications of electrochemical cells (A-level only)
The benefits and risks to society associated with using these cells.
- Evaluate the use of hydrogen fuel cells in buses, naming one benefit and one risk.
- Explain why the carbon footprint of a fuel-cell vehicle depends on how its hydrogen was manufactured.
- Justify why used lithium cells are collected separately rather than sent to landfill.
Commercial applications of electrochemical cells (A-level only)
The electrode reactions in an alkaline hydrogen–oxygen fuel cell.
- Write the half-equation occurring at the negative electrode of an alkaline hydrogen and oxygen fuel cell and check that charge balances.
- Combine the two electrode half-equations to show that water is the only product.
- Calculate the EMF of the alkaline fuel cell from E⦵ values of −0.83 V and +0.40 V.
Condensation polymers (A-level only)
Condensation polymers are formed by reactions between: dicarboxylic acids and diols dicarboxylic acids and diamines amino acids.
- Identify from a pair of monomer structures whether a polyester or a polyamide would form, and name the small molecule lost.
- Draw the link formed between two named monomers and show where the water molecule came from.
- Explain why each monomer in a condensation polymerisation must carry two functional groups.
Condensation polymers (A-level only)
Students should be able to: draw the repeating unit from monomer structure(s) draw the repeating unit from a section of the polymer chain draw the structure(s) of the monomer(s) from a section of the polymer explain the nature of the intermolecular forces between molecules of condensation polymers.
- Write the polymerisation equation using n molecules of each monomer, giving the number of water molecules released alongside the polymer.
- Break a drawn chain at its links and add back H and OH to the right atoms, so the monomers recovered are complete acids, diols or diamines and not fragments.
- Explain the difference in intermolecular forces between a polyester and a polyamide and its consequence for strength.
Condensation polymers (A-level only)
The repeating units in polyesters (eg Terylene) and polyamides (eg nylon 6,6 and Kevlar) and the linkages between these repeating units.
- Draw the repeating unit of Terylene, nylon 6,6 or Kevlar from the structures of its monomers.
- Decide from a drawn section of chain whether the link is an ester or an amide, and state the one atom that tells them apart.
- Explain what the numbers in the name nylon 6,6 refer to.
Condensation polymers (A-level only)
Typical uses of these polymers.
- Match Terylene, nylon 6,6 and Kevlar to a typical use and justify each match by a named property.
- Explain the strength of a named polyamide in terms of hydrogen bonding between its chains.
- Suggest why a polyester is chosen over a polyamide for a given application.
Definition and determination of pH (A-level only)
+ pH = –log [H ] Students should be able to: convert concentration of hydrogen ions into pH and vice versa calculate the pH of a solution of a strong acid from its concentration.
- Calculate the pH of 0.0500 mol dm⁻³ hydrochloric acid as 1.30.
- Convert a pH of 3.40 into a hydrogen ion concentration of 3.98 × 10⁻⁴ mol dm⁻³.
- Calculate the pH of 0.0200 mol dm⁻³ sulfuric acid, doubling the concentration before taking the logarithm.
Definition and determination of pH (A-level only)
The concentration of hydrogen ions in aqueous solution covers a very wide range.
- Compare the hydrogen ion concentrations of solutions at pH 1 and pH 3 as a factor of 100.
- Explain why 2.0 mol dm⁻³ hydrochloric acid can have a pH below zero.
- Deduce, from measured pH values of 0 to 1 and 2 to 3, which of two aqueous oxides is the stronger acid.
Definition and determination of pH (A-level only)
Therefore, a logarithmic scale, the pH scale, is used as a measure of hydrogen ion concentration.
- Calculate the pH of a solution in which [H⁺] is 2.5 × 10⁻³ mol dm⁻³, giving the answer to two decimal places.
- Explain why pure water at 323 K has a pH below 7 and is still neutral.
- Deduce the change in [H⁺] when the pH of a solution falls from 4.0 to 2.0.
Determination of rate equation (A-level only)
Required practical 7 Measuring the rate of reaction: by an initial rate method by a continuous monitoring method.
- Describe an initial rate experiment in which a fixed small extent of reaction is timed, and justify why rate is proportional to 1/t.
- Select a continuous monitoring technique appropriate to a given reaction and justify the choice.
- Determine the initial rate from a continuous monitoring graph by drawing a tangent at time zero.
- Explain why the total volume of the reaction mixture is kept constant when one concentration is varied.
Determination of rate equation (A-level only)
Students should be able to: use concentration–time graphs to deduce the rate of a reaction use initial concentration–time data to deduce the initial rate of a reaction use rate–concentration data or graphs to deduce the order (0, 1 or 2) with respect to a reactant derive the rate equation for a reaction from the orders with respect to each of the reactants use the orders with respect to reactants to provide information about the rate determining/limiting step of a reaction.
- Determine the initial rate of a reaction from the tangent at the start of a concentration against time graph.
- Deduce the order with respect to a reactant from the shape of a rate against concentration graph and state the graphical evidence.
- Derive the full rate equation from the individual orders and calculate the rate constant with its units.
- Use the deduced orders to identify which species appear in the rate determining step.
Determination of rate equation (A-level only)
The orders with respect to reactants can provide information about the mechanism of a reaction.
- Deduce which species are involved in the rate determining step from a given rate equation.
- Suggest a two step mechanism consistent with a stated rate equation and with the overall balanced equation.
- Explain what the zero order of a reactant reveals about where it is consumed in the mechanism.
- Justify why rate data can support a proposed mechanism without proving it.
Determination of rate equation (A-level only)
The rate equation is an experimentally determined relationship.
- Deduce the order with respect to each reactant from a table of initial rates by comparing experiments that differ in one concentration.
- Explain why the rate equation cannot be predicted from the stoichiometric equation, using a zero order reactant as evidence.
- Explain how a catalyst can appear in the rate equation of a reaction in which it does not appear in the overall equation.
- Describe an experimental procedure that isolates the effect of one reactant on the rate.
DNA (A-level only)
A nucleotide is made up from a phosphate ion bonded to 2-deoxyribose which is in turn bonded to one of the four bases adenine, cytosine, guanine and thymine.
- Construct the structure of a named nucleotide from the phosphate ion, 2-deoxyribose and a base given in the Data Booklet.
- Identify the two condensation reactions in that assembly and state the small molecule eliminated at each.
- Predict which atom of a base forms the bond to the sugar and explain why it is nitrogen rather than carbon.
- Compare the four DNA nucleotides and state the only part of the structure that differs between them.
DNA (A-level only)
A single strand of DNA (deoxyribonucleic acid) is a polymer of nucleotides linked by covalent bonds between the phosphate group of one nucleotide and the 2- deoxyribose of another nucleotide.
- Draw the link formed when two nucleotides join, showing one phosphate bridging the two sugars.
- Explain why the backbone is a condensation polymer and name the molecule eliminated.
- Compare the strength of the backbone bonds with the hydrogen bonds between strands and relate that difference to replication.
- Deduce how many water molecules are eliminated when a stated number of nucleotides polymerise.
DNA (A-level only)
DNA exists as two complementary strands arranged in the form of a double helix.
- State the two base pairs and give the number of hydrogen bonds in each.
- Draw the hydrogen bonds between a named base pair, showing the lone pair and the polar hydrogen involved.
- Deduce the base sequence of the partner strand from a sequence you are given.
- Explain why a region rich in cytosine and guanine needs more energy to separate than one rich in adenine and thymine.
DNA (A-level only)
The structures of the phosphate ion, 2-deoxyribose (a pentose sugar) and the four bases adenine, cytosine, guanine and thymine are given in the Chemistry Data Booklet.
- Use the booklet structure of guanine to pick out the ring nitrogen that keeps an unused lone pair, and distinguish it from the nitrogen that bonds to the sugar.
- Distinguish 2-deoxyribose from ribose by the missing OH group on carbon two, and say why that difference names the sugar.
- Classify adenine and guanine as purines and cytosine and thymine as pyrimidines, using the number of rings in the booklet structures.
- Identify, on a booklet structure, the OH group and the nitrogen atom that react when a nucleotide is assembled.
DNA (A-level only)
This results in a sugar- phosphate-sugar-phosphate polymer chain with bases attached to the sugars in the chain.
- Sketch three repeat units of the backbone, placing each base on a sugar.
- Explain why the backbone carries no genetic information while the base sequence does.
- Relate the negative charge on the phosphate groups to the position of the backbone on the outside of the helix.
- Identify an incorrect DNA diagram by checking what each base is attached to.
Effect of concentration and pressure
Students should be able to: explain how a change in concentration or a change in pressure influences the rate of a reaction.
- Explain how an increase in gas pressure increases concentration and collision frequency.
- Explain the effect of concentration on reaction rate in solutions.
- Justify why the Maxwell-Boltzmann distribution is unchanged by pressure.
Effect of concentration and pressure
The qualitative effect of a change in the pressure of a gas on collision frequency.
- Explain the effect of increasing the pressure of a gaseous reaction mixture on the rate, working from volume to concentration to collision frequency.
- Predict the effect on rate of adding an unreactive gas at constant volume and justify the prediction.
- Identify when a pressure argument is inappropriate because the reactants are not gases.
- Contrast the effect of pressure on the rate of a reaction with its effect on the position of equilibrium.
Effect of concentration and pressure
The qualitative effect of changes in concentration on collision frequency.
- Explain an increase in rate on increasing concentration using a chain from particle density to successful collision frequency.
- State what a concentration change does not alter about the distribution of molecular energies.
- Explain why a reaction is fastest at the start and slows as time passes.
- Distinguish the reasoning used for a concentration change from that used for a temperature change.
Effect of temperature on reaction rate
Meaning of the term rate of reaction.
- Define rate of reaction with reference to concentration and time and state its units.
- Determine the rate at a stated time by drawing a tangent to a concentration against time graph and calculating its gradient.
- Select an appropriate measurable quantity to follow a given reaction and justify why it tracks concentration.
- Explain why 1/t is a valid measure of rate when the time for a fixed visible change is recorded.
Effect of temperature on reaction rate
Required practical 3 Investigation of how the rate of a reaction changes with temperature.
- Describe a method for measuring how the rate of a named reaction changes with temperature, using a fixed visible endpoint.
- Identify the control variables in the experiment and explain the effect of failing to control each one.
- Plot 1/t against temperature, describe the shape obtained and relate it to the Maxwell–Boltzmann distribution.
- Suggest an improvement that removes the subjective judgement of the endpoint and explain why it improves the data.
Effect of temperature on reaction rate
Students should be able to: use the Maxwell–Boltzmann distribution to explain why a small temperature increase can lead to a large increase in rate.
- Explain, using the shape of the energy tail, why a small rise in temperature produces a large rise in rate.
- Sketch two labelled distribution curves and use the areas beyond Eₐ to justify a stated change in rate.
- Compare the contribution of increased collision frequency with that of the increased proportion of energetic molecules.
- Criticise an answer that explains a temperature effect using collision frequency alone, and correct it.
Effect of temperature on reaction rate
The qualitative effect of temperature changes on the rate of reaction.
- Describe the shape of a graph of rate against temperature and state what the shape shows about how rate changes.
- Explain a temperature induced rate increase using both collision frequency and the proportion of collisions exceeding Eₐ, and say which dominates.
- Predict the effect on rate of a stated temperature rise for a reaction and justify the prediction.
- Identify why an answer referring only to more frequent collisions is incomplete for a temperature change.
Electrode potentials and cells (A-level only)
Cells are used to measure electrode potentials by reference to the standard hydrogen electrode.
- Draw and label a standard hydrogen electrode, marking the platinum, the gas pressure and the acid concentration.
- Explain why potassium nitrate rather than potassium chloride is chosen for the salt bridge in a cell containing silver ions.
- Deduce which electrode is negative when a standard zinc half-cell is connected to a standard hydrogen electrode.
Electrode potentials and cells (A-level only)
IUPAC convention for writing half-equations for electrode reactions.
- Write the Fe³⁺/Fe²⁺ half-equation as a reduction and show that the charge is +2 on each side.
- Balance the half-equation for MnO₄⁻ in acid by adding 8H⁺, 4H₂O and 5 electrons.
- Deduce what happens to the sign of E⦵ when a half-equation is written in the reverse direction.
Electrode potentials and cells (A-level only)
Required practical 8 Measuring the EMF of an electrochemical cell.
- Describe how to set up a zinc and copper cell and measure its EMF with a high-resistance voltmeter.
- Justify the choice of potassium nitrate over potassium chloride for the salt bridge in a cell containing a silver half-cell.
- Suggest why a measured EMF differs from the value calculated from standard electrode potentials.
Electrode potentials and cells (A-level only)
Standard electrode potentials can be listed as an electrochemical series.
- Rank zinc, iron and copper as reducing agents from E⦵ values of −0.76 V, −0.44 V and +0.34 V.
- Deduce from a table of E⦵ values whether acidified manganate(VII) will oxidise chloride ions to chlorine.
- Explain why a redox reaction with a positive EMF may still not be observed in the laboratory.
Electrode potentials and cells (A-level only)
Students should be able to: Ɵ use E values to predict the direction of simple redox reactions calculate the EMF of a cell write and apply the conventional representation of a cell.
- Calculate the EMF of the cell Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s) from E⦵ values of −0.76 V and +0.34 V.
- Combine two half-equations so that the electrons cancel and write the balanced overall redox equation.
- Write the conventional representation of a cell, including phase boundaries, salt bridge, and Pt electrodes where necessary.
Electrode potentials and cells (A-level only)
The conventional representation of cells.
- Write the cell diagram for a zinc half-cell joined to a copper half-cell.
- Justify the inclusion of a platinum electrode in the cell diagram for an Fe³⁺/Fe²⁺ half-cell.
- Calculate the EMF of a cell directly from its cell diagram using E⦵(right) − E⦵(left).
Electrode potentials and cells (A-level only)
The importance of the conditions when measuring the electrode potential, E (Nernst equation not required).
- Predict how E for Cu²⁺(aq) + 2e⁻ ⇌ Cu(s) changes when the copper ion concentration is halved.
- Identify, from a list of changes to a hydrogen electrode, the one that leaves the electrode potential unchanged.
- Explain why a potential measured with 0.10 mol dm⁻³ solutions may not be quoted as a standard value.
Electron configuration
Electron configurations of atoms and ions up to Z = 36 in terms of shells and sub-shells (orbitals) s, p and d.
- Write the full sub-shell electron configuration of any atom or ion up to krypton, adjusting the electron count for charge.
- Explain why iron(II) is written with no 4s electrons despite 4s filling before 3d in the atom.
- Complete an electrons-in-boxes diagram showing single occupancy before pairing and opposite spins in a filled orbital.
Electron configuration
Ionisation energies.
- Define first ionisation energy in a form that specifies one mole, the gaseous state and the charge on the product ion.
- Explain, using nuclear charge, radius and shielding, why one element has a higher first ionisation energy than another.
- Justify why ionisation is always an endothermic change.
Electron configuration
Students should be able to: define first ionisation energy write equations for first and successive ionisation energies explain how first and successive ionisation energies in Period 3 (Na–Ar) and in Group 2 (Be–Ba) give evidence for electron configuration in sub-shells and in shells.
- Define first ionisation energy and write equations, with state symbols, for first and successive ionisation energies.
- Deduce the group of an element from the position of the large jump in its successive ionisation energies.
- Explain the dips at aluminium and at sulfur in terms of sub-shell energy and electron pairing respectively.
Electrophilic substitution (A-level only)
Electrophilic attack on benzene rings results in substitution, limited to monosubstitutions.
- Draw the general two-step mechanism for electrophilic substitution on benzene using curly arrows.
- Draw the positively charged intermediate correctly, showing the horseshoe, the charge and the hydrogen atom.
- Explain why the ring loses H⁺ rather than adding the electrophile permanently.
Electrophilic substitution (A-level only)
Friedel–Crafts acylation reactions are also important steps in synthesis.
- Write equations for the formation of the acylium ion and for the regeneration of aluminium chloride.
- Predict and name the product of a named acyl chloride reacting with benzene in the presence of AlCl₃.
- Point to the new carbon-carbon bond formed in a Friedel-Crafts acylation and state what it puts onto the ring that nitration cannot.
Electrophilic substitution (A-level only)
Nitration is an important step in synthesis, including the manufacture of explosives and formation of amines.
- State the reagents and conditions for the nitration of benzene and explain why the temperature is limited.
- Write a balanced equation for the reduction of nitrobenzene to phenylamine using [H].
- Outline a two-stage synthesis that converts benzene into an aromatic amine, naming the reagents at each stage.
Electrophilic substitution (A-level only)
Students should be able to outline the electrophilic substitution mechanisms of: nitration, including the generation of the nitronium ion acylation using AlCl as a catalyst.
- Write the equations that generate the nitronium ion and the acylium ion from the reagents given.
- Draw the complete electrophilic substitution mechanism for the nitration of benzene, with all charges and arrows.
- Draw the mechanism for the acylation of benzene and write the equation regenerating the aluminium chloride.
Elimination
Alkenes can be formed from alcohols by acid-catalysed elimination reactions.
- Write the equation for the acid-catalysed dehydration of a named alcohol, showing water as a product.
- Deduce every alkene that can be formed from an unsymmetrical alcohol, including E and Z isomers.
- Explain why the acid used is described as a catalyst and does not appear in the overall equation.
Elimination
Alkenes produced by this method can be used to produce addition polymers without using monomers derived from crude oil.
- Describe the three-stage route from plant sugar to poly(ethene), naming the reagents and conditions at each stage.
- Explain why using this route does not make the resulting poly(ethene) biodegradable.
- Discuss one benefit and one drawback of making addition polymers from crop-derived ethanol.
Elimination
Required practical 5 Distillation of a product from a reaction.
- Label a distillation apparatus, placing the thermometer and the condenser water flow correctly.
- Explain why the distillate is collected only over a narrow temperature range.
- Justify each step of the work-up, including washing with sodium carbonate and drying with an anhydrous salt.
Elimination
Students should be able to: outline the mechanism for the elimination of water from alcohols.
- Draw the three-step elimination mechanism for the dehydration of a named alcohol.
- Identify which curly arrow is responsible for forming the carbon–carbon double bond.
- Explain why the alcohol is protonated first, referring to the quality of the leaving group.
Elimination
The concurrent substitution and elimination reactions of a halogenoalkane (eg 2-bromopropane with potassium hydroxide).
- Name both mechanisms operating when 2-bromopropane is warmed with potassium hydroxide and give the product of each.
- Explain how changing from aqueous to ethanolic solvent alters the ratio of alkene to alcohol.
- Draw both mechanisms side by side, labelling hydroxide as nucleophile in one and base in the other.
Empirical and molecular formula
Empirical formula is the simplest whole number ratio of atoms of each element in a compound.
- Calculate an empirical formula from percentage composition by mass, including an element given only as the remainder.
- Convert a ratio containing a half or a third into whole numbers by multiplying the whole set.
- Explain why two different compounds can share the same empirical formula.
Empirical and molecular formula
Molecular formula is the actual number of atoms of each element in a compound.
- Deduce a molecular formula by counting the atoms shown in a displayed or skeletal structure.
- Explain why a molecular formula cannot distinguish between two isomers.
- Identify cases in which the molecular and empirical formulae of a compound are identical.
Empirical and molecular formula
Students should be able to: calculate empirical formula from data giving composition by mass or percentage by mass calculate molecular formula from the empirical formula and relative molecular mass.
- Calculate an empirical formula from percentage composition where one element is given only as the remainder.
- Determine a molecular formula from an empirical formula and a supplied relative molecular mass.
- Decide when a ratio containing 1.5 must be doubled rather than rounded.
Empirical and molecular formula
The relationship between empirical formula and molecular formula.
- Calculate the multiplier linking an empirical formula to a molecular formula from the relative molecular mass.
- Deduce a molecular formula by scaling every subscript of an empirical formula.
- Verify a deduced molecular formula by recalculating its relative molecular mass.
Enthalpy change
Enthalpy change (∆H) is the heat energy change measured under conditions of constant pressure.
- Define ∆H in a form that includes heat energy, the constant pressure condition and the per mole basis.
- Convert a measured heat change q in joules into an enthalpy change in kJ mol⁻¹ for a stated substance.
- Explain why a reaction run in an open polystyrene cup gives ∆H directly, while the same reaction in a sealed rigid container does not.
- Distinguish between the total heat released by a sample and the enthalpy change per mole, using an example in which the amounts are doubled.
Enthalpy change
Reactions can be endothermic or exothermic.
- Deduce the sign of ∆H from a measured temperature change in a reaction mixture, stating whether heat left or entered the system.
- Explain the sign of ∆H for a named reaction by comparing the energy needed to break the reactant bonds with the energy released forming the product bonds.
- Sketch an enthalpy profile for an exothermic and an endothermic reaction, labelling reactants, products, ∆H and the activation energy.
- Account for a measured enthalpy of combustion being less exothermic than the data book value, quoting incomplete combustion or heat loss.
Enthalpy change
Standard enthalpy changes refer to standard conditions, Ɵ ie 100 kPa and a stated temperature (eg ∆H ).
- State the conditions AQA treats as standard and identify the standard state of water, bromine, iodine and carbon at 298 K.
- Write an equation with state symbols whose enthalpy change is the standard enthalpy of formation of a named compound, using fractions where needed.
- Explain why the standard enthalpy of formation of an element in its standard state is zero.
- Identify which of a set of given equations does not represent a standard enthalpy of combustion, and say why.
Enzymes (A-level only)
Computers can be used to help design such drugs.
- Describe how a computer model of an active site is used to identify possible drug molecules.
- Explain two advantages of modelling a drug before synthesising it.
- State why the model must take account of the chirality of the candidate molecule.
Enzymes (A-level only)
Enzymes are proteins.
- State what an enzyme is made of and which level of protein structure creates its active site.
- Explain what happens to an enzyme when it is heated strongly, in terms of the bonds broken.
- Explain why a denatured enzyme no longer works even though its primary structure is unchanged.
Enzymes (A-level only)
Students should be able to: explain why a stereospecific active site can only bond to one enantiomeric form of a substrate or drug.
- Explain what makes a molecule chiral and identify a chiral carbon in a given structure.
- Judge the claim that enzymes work equally well on both optical isomers of a substrate, and give the reason it is false in terms of the shape of the active site.
- Deduce the consequence for the effectiveness of a drug supplied as a racemic mixture.
Enzymes (A-level only)
The action of enzymes as catalysts, including the concept of a stereospecific active site that binds to a substrate molecule.
- Explain how forming an enzyme-substrate complex speeds a reaction up, naming the alternative route of lower activation energy rather than a lowering of the original route's activation energy.
- Explain what 'stereospecific active site' means and what follows from it for a pair of optical isomers.
- State two properties of a catalyst and show that an enzyme has both of them.
Enzymes (A-level only)
The principle of a drug acting as an enzyme inhibitor by blocking the active site.
- Explain why an inhibitor that occupies the active site stops the enzyme working, even though the inhibitor is not itself changed by the enzyme.
- State the structural requirement an inhibitor must meet and explain why it is necessary.
- Suggest why a drug that inhibits a bacterial enzyme may be safe for the patient.
Forces between molecules
Forces between molecules: permanent dipole–dipole forces induced dipole–dipole (van der Waals, dispersion, London) forces hydrogen bonding.
- Identify every intermolecular force acting in a named substance and justify each from its structure.
- Draw a diagram of a hydrogen bond between two molecules, showing partial charges and the lone pair involved.
- Explain why induced dipole-dipole forces increase in strength as a homologous series is ascended.
Forces between molecules
Students should be able to: explain the existence of these forces between familiar and unfamiliar molecules explain how melting and boiling points are influenced by these intermolecular forces.
- Deduce, from the displayed structure of an unfamiliar molecule, which intermolecular forces act between its molecules.
- Explain a trend in boiling point across three related compounds by naming the strongest force in each.
- Justify which of two hydrogen bonded liquids has the higher boiling point using the electronegativity of the donor atom.
Forces between molecules
The importance of hydrogen bonding in the low density of ice and the anomalous boiling points of compounds.
- Explain why ice is less dense than liquid water in terms of the number and arrangement of its hydrogen bonds.
- Account for the boiling point of water lying far above that of the other Group 6 hydrides.
- Draw the hydrogen bonding between two water molecules, showing lone pairs and partial charges.
Forces between molecules
The melting and boiling points of molecular substances are influenced by the strength of these intermolecular forces.
- Explain a difference in boiling point between two substances by naming and comparing their intermolecular forces.
- State precisely what is overcome when a molecular solid melts.
- Predict which of two isomers with the same molecular formula boils at the higher temperature and justify the prediction.
Formation of coloured ions (A-level only)
A simple colorimeter can be used to determine the concentration of coloured ions in solution.
- Outline a colorimetric method for finding the concentration of a coloured ion, from filter choice to final reading.
- Plot and use a calibration curve of absorbance against concentration.
- Explain why a blank is used and why the same filter and cell must be used throughout.
Formation of coloured ions (A-level only)
Colour arises when some of the wavelengths of visible light are absorbed and the remaining wavelengths of light are transmitted or reflected.
- Explain the colour of a named complex by stating which light is absorbed and which is transmitted.
- Distinguish absorption from emission when accounting for the colour of a solution.
- Predict how the observed colour relates to the wavelengths removed from white light.
Formation of coloured ions (A-level only)
d electrons move from the ground state to an excited state when light is absorbed.
- Describe the splitting of the d sub-level by ligands and label the gap between the two levels.
- Explain absorption of light as promotion of a d electron from the ground state to the excited state.
- Deduce whether a given ion will be coloured from the number of electrons in its d sub-level.
Formation of coloured ions (A-level only)
The absorption of visible light is used in spectroscopy.
- Interpret an absorption spectrum to find λmax for a coloured complex.
- Calculate ΔE for a complex from its λmax using ΔE = hc/λ.
- Justify the choice of colour filter for a colorimetric measurement on a given coloured solution.
Formation of coloured ions (A-level only)
The energy difference between the ground state and the excited state of the d electrons is given by: ∆E = hν = hc/λ Changes in oxidation state, co-ordination number and ligand alter ∆E and this leads to a change in colour.
- Calculate ΔE from a given wavelength or frequency, converting units to metres first.
- State the three changes to a complex that alter ΔE and give an example of each.
- Explain why two copper(II) complexes with different ligands are different colours.
Formation of coloured ions (A-level only)
Transition metal ions can be identified by their colour.
- Recall the colours of the common hexaaqua ions of copper, iron, cobalt, chromium and manganese.
- Identify an unknown transition metal ion from an observed colour and its change on adding a named reagent.
- Explain why two solutions containing the same metal can differ in colour by referring to ligand or oxidation state.
Fractional distillation of crude oil
Alkanes are saturated hydrocarbons.
- Define saturated hydrocarbon, explaining both words separately.
- Deduce the molecular formula of any open-chain alkane or cycloalkane from its number of carbon atoms.
- Explain the low reactivity of alkanes in terms of bond strength and the absence of bond polarity.
Fractional distillation of crude oil
Petroleum is a mixture consisting mainly of alkane hydrocarbons that can be separated by fractional distillation.
- Describe how a fractionating column separates crude oil using a temperature gradient.
- Explain the order in which fractions are collected by relating chain length to London forces and boiling point.
- State why each fraction is a mixture rather than a single compound.
Fundamental particles
An atom consists of a nucleus containing protons and neutrons surrounded by electrons.
- Compare the diameter of the nucleus with that of the whole atom and use the comparison to explain where the mass is concentrated.
- Explain in terms of electrostatic attraction why electrons remain bound to the nucleus.
- Deduce the charge of a species from the numbers of protons and electrons given for it.
Fundamental particles
Appreciate that knowledge and understanding of atomic structure has evolved over time.
- Describe the gold foil observation of alpha particles deflected through large angles and deduce from it that the nucleus is tiny, dense and positively charged.
- Explain which observation the plum pudding model failed to account for and what Rutherford changed as a result.
- State which particle Chadwick added and why the mass of an atom could not be explained by protons and electrons alone.
Fundamental particles
Protons, neutrons and electrons: relative charge and relative mass.
- Complete a table of relative charge and relative mass for the proton, neutron and electron without reference to data.
- Explain why the mass of an atom can be treated as the sum of its protons and neutrons alone.
- Deduce the overall charge on a particle from the numbers of protons and electrons it contains.
General properties of transition metals (A-level only)
A complex is a central metal atom or ion surrounded by ligands.
- State the meaning of the term complex, covering both a central metal atom and a central metal ion.
- Deduce the oxidation state of the metal in a complex from the overall charge and the ligand charges.
- Write the formula of a complex with square brackets placed correctly around the co-ordination sphere.
General properties of transition metals (A-level only)
A ligand is a molecule or ion that forms a co-ordinate bond with a transition metal by donating a pair of electrons.
- Define a ligand in terms of lone pair donation and name the type of bond it forms.
- Identify the donor atom in H₂O, NH₃, Cl⁻, CN⁻, C₂O₄²⁻ and EDTA⁴⁻ and classify each ligand by denticity.
- Draw a co-ordinate bond with the arrow pointing in the correct direction, from ligand to metal.
General properties of transition metals (A-level only)
Co-ordination number is number of co-ordinate bonds to the central metal atom or ion.
- Use a co-ordination number of 6 or 4 to predict the shape and bond angle of a complex: 6 gives octahedral at 90°, 4 usually gives tetrahedral at 109.5°.
- Deduce the co-ordination number of complexes containing monodentate, bidentate and multidentate ligands.
- Predict whether a given ligand produces a co-ordination number of 4 or 6 and justify the prediction by ligand size.
General properties of transition metals (A-level only)
The characteristic properties include: complex formation formation of coloured ions variable oxidation state catalytic activity.
- List the four characteristic properties and trace colour, variable oxidation states, and catalytic activity back to the partially filled d sub-level.
- State the three changes to a complex that alter its colour and give an example of each.
- Contrast Cu²⁺ with Zn²⁺, identifying which characteristic properties Zn²⁺ lacks and explaining why.
General properties of transition metals (A-level only)
Transition metal characteristics of elements Ti–Cu arise from an incomplete d sub-level in atoms or ions.
- Define a transition metal in terms of forming at least one stable ion with a partially filled d sub-level.
- Write the electron configuration of any atom or ion from titanium to copper, removing 4s electrons before 3d.
- Justify the exclusion of scandium and zinc by quoting the configurations of Sc³⁺ and Zn²⁺.
Gibbs free-energy change, ∆G, and entropy change, ∆S (A-level only)
∆H, whilst important, is not sufficient to explain feasible change.
- Explain, with a named example, why the sign of ∆H alone cannot decide whether a change is feasible.
- Identify the entropy change as the factor that allows an endothermic process to be spontaneous.
- Predict which of two endothermic processes is more likely to be feasible at room temperature and justify the choice.
- Distinguish thermodynamic feasibility from the observed rate of a reaction.
Gibbs free-energy change, ∆G, and entropy change, ∆S (A-level only)
∆S accounts for the above deficiency, illustrated by physical changes and chemical changes.
- Final objective: Explain why dissolving certain ionic compounds can result in a negative entropy change due to hydration.
- Established context: Predict and justify the sign of ∆S for melting and boiling.
- Established context: Predict the sign of ∆S for a chemical reaction based on moles of gas.
Gibbs free-energy change, ∆G, and entropy change, ∆S (A-level only)
For a reaction to be feasible, the value of ∆G must be zero or negative.
- Deduce, from a quoted ∆G for a forward reaction at a stated temperature, whether the reverse reaction is feasible at that temperature.
- Calculate ∆G for an endothermic process and deduce whether it is feasible at the stated temperature.
- Explain why a reaction with a large negative ∆G may not be observed to occur.
- Justify a conclusion about feasibility by quoting the sign of the calculated ∆G rather than the sign of ∆H.
Gibbs free-energy change, ∆G, and entropy change, ∆S (A-level only)
Students should be able to: calculate entropy changes from absolute entropy values use the relationship ∆G = ∆H – T∆S to determine how ∆G varies with temperature use the relationship ∆G = ∆H – T∆S to determine the temperature at which a reaction becomes feasible.
- Calculate ∆S for a reaction from a table of absolute entropies, including the balancing numbers.
- Determine the temperature at which a reaction becomes feasible by setting ∆G to zero and rearranging.
- Deduce whether a reaction becomes feasible above or below that temperature from the sign of ∆S.
- Explain how ∆G changes with temperature by treating ∆G = ∆H − T∆S as a straight line graph.
Gibbs free-energy change, ∆G, and entropy change, ∆S (A-level only)
The balance between entropy and enthalpy determines the feasibility of a reaction given by the relationship: ∆G = ∆H – T∆S (derivation not required).
- Calculate ∆G at a stated temperature from ∆H and ∆S, converting the units so the two terms are compatible.
- Predict, from the signs of ∆H and ∆S alone, the temperature range over which a reaction is feasible.
- Explain why the entropy term becomes more important as the temperature is raised.
- Identify the unit error in a worked calculation whose T∆S term is a thousand times too large.
Gibbs free-energy change, ∆G, and entropy change, ∆S (A-level only)
The concept of increasing disorder (entropy change, ∆S).
- Rank a set of substances in order of increasing entropy and justify the order using their states.
- Predict the sign of ∆S for a reaction by comparing the moles of gas on each side.
- Calculate ∆S from a table of absolute entropies, quoting the correct units.
- Explain why a perfect crystal at absolute zero is assigned an entropy of zero.
Group 2, the alkaline earth metals
The reactions of the elements Mg–Ba with water.
- Write balanced equations for magnesium with cold water and with steam, naming the different products.
- Explain why barium reacts with cold water faster than calcium does, using ionisation energies.
- Explain why magnesium hydroxide can be swallowed as an antacid while barium hydroxide solution cannot, referring to solubility.
Group 2, the alkaline earth metals
The relative solubilities of the sulfates of the elements Mg–Ba in water.
- Describe the test for a sulfate ion, naming the acid used and giving the ionic equation for the precipitate.
- Explain why hydrochloric acid rather than sulfuric or nitric acid is used to acidify a solution before barium chloride is added.
- Explain why barium sulfate can be swallowed as a barium meal even though barium ions are toxic.
Group 2, the alkaline earth metals
The trends in atomic radius, first ionisation energy and melting point of the elements Mg–Ba Students should be able to: explain the trends in atomic radius and first ionisation energy explain the melting point of the elements in terms of their structure and bonding.
- Explain why the first ionisation energy of barium is lower than that of magnesium, using distance and shielding.
- Explain why the second ionisation energy of calcium is greater than its first, referring to the charge and size of the particle involved.
- Predict the order of melting points of calcium, strontium and barium and justify it from ionic size.
Identification of functional groups by test-tube reactions
Required practical 6 Tests for alcohol, aldehyde, alkene and carboxylic acid.
- Carry out and describe the four specified tests, giving the observation for each functional group.
- Identify errors in a given practical method and state the correction for each.
- Justify heating in a water bath rather than over a flame when testing organic liquids.
Identification of functional groups by test-tube reactions
Students should be able to: identify the functional groups using reactions in the specification.
- Select a single test that distinguishes two given organic compounds and state both observations.
- Order a set of tests so that one result does not interfere with the next.
- Deduce the functional group present in an unknown from a set of given test results.
Identification of functional groups by test-tube reactions
The reactions of functional groups listed in the specification.
- Match each functional group on the specification to its test reagent, condition and positive observation.
- Describe the observation for a compound that gives a negative result in a named test.
- Select a single test that would distinguish two named compounds and justify the choice.
Infrared spectroscopy
‘Fingerprinting’ allows identification of a molecule by comparison of spectra.
- State where the fingerprint region lies and explain why its pattern is unique to a compound.
- Explain how a compound is identified from its fingerprint region, naming what the spectrum is compared with.
- Justify choosing infrared fingerprinting over mass spectrometry to distinguish two positional isomers.
Infrared spectroscopy
Bonds in a molecule absorb infrared radiation at characteristic wavenumbers.
- Explain why different bonds absorb infrared radiation at different wavenumbers.
- Use the data sheet to quote the wavenumber range for a named bond, with the correct unit.
- Identify the bond responsible for a labelled absorption on a given infrared spectrum.
Infrared spectroscopy
Students should be able to: use infrared spectra and the Chemistry Data Sheet or Booklet to identify particular bonds, and therefore functional groups, and also to identify impurities.
- Deduce the functional group present in a compound from the peaks present and absent in its infrared spectrum.
- Distinguish an alcohol from a carboxylic acid using the position and shape of the O–H absorption.
- Identify an impurity in a product from an unexpected absorption and name the likely substance.
Ionic bonding
Ionic bonding involves electrostatic attraction between oppositely charged ions in a lattice.
- Explain the difference in melting point between magnesium oxide and sodium chloride using ionic charge and radius.
- Describe the arrangement of ions in the sodium chloride lattice and the direction in which the attraction acts.
- Justify why an ionic solid does not conduct electricity but its molten form does.
Ionic bonding
Students should be able to: predict the charge on a simple ion using the position of the element in the Periodic Table construct formulas for ionic compounds.
- Predict the charge on the ion formed by a main-group element from its position in the Periodic Table.
- Construct the formula of an ionic compound by balancing total positive against total negative charge.
- Deduce the charge on a transition metal ion from the Roman numeral in the compound's name.
Ionic bonding
The formulas of compound ions, eg sulfate, hydroxide, nitrate, carbonate and ammonium.
- Write the formula, including charge, for sulfate, hydroxide, nitrate, carbonate and ammonium from memory.
- Construct the formula of a salt that requires more than one compound ion, using brackets correctly.
- Calculate the relative formula mass of a compound whose formula contains a bracketed compound ion.
Isomerism
Cahn–Ingold–Prelog (CIP) priority rules.
- Apply CIP rules to assign priorities to the two groups on each carbon of a C=C.
- Deduce whether a drawn alkene is the E or the Z isomer and name it in full.
- Resolve a priority tie by comparing the next set of atoms along each branch.
Isomerism
E–Z isomerism is a form of stereoisomerism and occurs as a result of restricted rotation about the planar carbon– carbon double bond.
- Explain restricted rotation about a C=C in terms of the π bond formed by sideways p orbital overlap.
- Deduce whether a named alkene can show E-Z isomerism by inspecting the groups on each double bond carbon.
- Name both isomers of a given alkene and draw them showing the planar arrangement.
Isomerism
Stereoisomerism.
- Define stereoisomerism and distinguish it from structural isomerism.
- Identify whether a given molecule can show E-Z isomerism, optical isomerism or neither.
- Draw two stereoisomers of a named compound so that the spatial difference is clear.
Isomerism
Structural isomerism.
- Define a structural isomer in terms of molecular formula and connectivity.
- Draw a pair of chain, a pair of position and a pair of functional group isomers for a given formula.
- Classify a given pair of structures as chain, position or functional group isomers.
Isomerism
Students should be able to: define the term structural isomer draw the structures of chain, position and functional group isomers define the term stereoisomer draw the structural formulas of E and Z isomers apply the CIP priority rules to E and Z isomers.
- Define structural isomer and stereoisomer in ways that clearly distinguish the two.
- Draw chain, position and functional group isomers for a given molecular formula such as C₄H₈.
- Apply CIP priority rules to label a drawn alkene as the E or the Z isomer.
Mass number and isotopes
Mass number (A) and atomic (proton) number (Z).
- Deduce the number of neutrons in a nuclide from its mass number and atomic number.
- Explain why a mass number must be a whole number whereas a relative atomic mass usually is not.
- Identify an element from its atomic number alone and justify why the neutron count does not change that identification.
Mass number and isotopes
Mass spectrometry can be used to determine relative molecular mass.
- Identify the molecular ion peak in a spectrum and use it to state the relative molecular mass of the compound.
- Explain why the base peak cannot be used to obtain relative molecular mass.
- Account for the small peak one mass unit above the molecular ion peak.
Mass number and isotopes
Mass spectrometry can be used to identify elements.
- Deduce the identity of an element from the mass-to-charge values and relative heights of the peaks in its spectrum.
- Explain why one peak alone is insufficient evidence to name an element.
- Predict the number and relative heights of the peaks expected from a sample of chlorine atoms.
Mass number and isotopes
Students should be able to: determine the number of fundamental particles in atoms and ions using mass number, atomic number and charge explain the existence of isotopes.
- Determine the numbers of protons, neutrons and electrons in a given ion by subtracting the signed charge from the proton count.
- Explain why two nuclides of the same element differ in mass number but not in chemistry.
- Justify, in terms of electron configuration rather than nuclear composition, why isotopes react identically.
Mass number and isotopes
Students should be able to: interpret simple mass spectra of elements calculate relative atomic mass from isotopic abundance, limited to mononuclear ions.
- Calculate the relative atomic mass of an element from a table of isotopic masses and abundances.
- Rearrange the weighted-mean expression to find a missing isotopic mass from a known relative atomic mass.
- Check a calculated relative atomic mass against the abundances by confirming it lies closest to the most abundant isotope.
Mass number and isotopes
The mass spectrometer gives accurate information about relative isotopic mass and also about the relative abundance of isotopes.
- Calculate an unknown relative isotopic mass from the relative atomic mass and the abundances of the other isotopes.
- Explain how the height of a peak is generated physically inside the instrument.
- Convert a set of relative abundances into fractions of the total before using them in a mean.
Mass number and isotopes
The principles of a simple time of flight (TOF) mass spectrometer, limited to ionisation, acceleration to give all ions constant kinetic energy, ion drift, ion detection, data analysis.
- Calculate the speed of an ion from its kinetic energy and mass, then the flight time from the length of the drift region.
- Explain why two ions of different mass given identical kinetic energy arrive at the detector at different times.
- Describe what happens at the detector and how the current produced relates to isotopic abundance.
Mass spectrometry
Mass spectrometry can be used to determine the molecular formula of a compound.
- Identify the molecular ion peak on a mass spectrum and state the relative molecular mass it gives.
- Explain why high resolution is needed to choose between formulae of the same whole-number mass.
- Explain why mass spectrometry alone cannot distinguish two structural isomers.
Mass spectrometry
Students should be able to: use precise atomic masses and the precise molecular mass to determine the molecular formula of a compound.
- Calculate the precise molecular mass of a candidate formula from precise atomic masses.
- Reject a candidate formula by quoting its calculated precise mass beside the measured value and stating the difference that rules it out.
- Explain why average relative atomic masses cannot be used in this calculation.
Maxwell–Boltzmann distribution
Maxwell–Boltzmann distribution of molecular energies in gases.
- Sketch a Maxwell–Boltzmann distribution with correctly labelled axes, starting at the origin and with an asymptotic tail.
- Identify the most probable energy and the mean energy on a distribution curve and explain why they differ.
- Interpret the area under a region of the curve in terms of numbers of molecules.
- Explain what the shaded area beyond Eₐ represents and why it determines the rate.
Maxwell–Boltzmann distribution
Students should be able to: draw and interpret distribution curves for different temperatures.
- Draw two Maxwell–Boltzmann distributions for the same gas at different temperatures on one pair of labelled axes.
- Explain why the two curves must enclose the same total area and cross only once.
- Interpret the change in the shaded area beyond Eₐ when the temperature is raised, and link it to rate.
- Identify errors in a given sketch of two distribution curves and state how each should be corrected.
Metallic bonding
Metallic bonding involves attraction between delocalised electrons and positive ions arranged in a lattice.
- Explain why magnesium has a higher melting point than sodium in terms of ionic charge and delocalised electrons.
- Describe metallic bonding using the three elements of the mark scheme definition.
- Account for the malleability and electrical conductivity of a metal from its structure.
Modification of alkanes by cracking
Catalytic cracking takes place at a slight pressure, high temperature and in the presence of a zeolite catalyst and is used mainly to produce motor fuels and aromatic hydrocarbons (mechanism not required).
- Explain why the catalytic route is chosen to make motor fuel, naming the branched alkanes, cycloalkanes and aromatics it produces.
- Write a balanced equation for the catalytic cracking of a named alkane into a fuel-range alkane and a cycloalkane.
- Explain why the catalytic route is cheaper to operate than the thermal route.
Modification of alkanes by cracking
Cracking involves breaking C–C bonds in alkanes.
- Write a balanced equation for the cracking of a named long chain alkane.
- Explain the economic reason for cracking in terms of the supply of and demand for different fractions.
- Identify which bonds break during cracking and why an unsaturated product is usually formed.
Modification of alkanes by cracking
Thermal cracking takes place at high pressure and high temperature and produces a high percentage of alkenes (mechanism not required).
- State the conditions used for thermal cracking with appropriate units.
- Write a balanced equation for a thermal cracking reaction that produces an alkene.
- Compare thermal with catalytic cracking in terms of conditions and product mix.
Nature of covalent and dative covalent bonds
A co-ordinate (dative covalent) bond contains a shared pair of electrons with both electrons supplied by one atom.
- Identify the lone-pair donor and the electron-deficient acceptor in the formation of a named dative covalent bond.
- Explain why all four bonds in the ammonium ion are identical despite one forming differently.
- Draw the product of ammonia reacting with boron trifluoride, showing the origin of the shared pair.
Nature of covalent and dative covalent bonds
A single covalent bond contains a shared pair of electrons.
- Draw a dot-and-cross diagram for a simple covalent molecule showing shared pairs and lone pairs separately.
- Explain in terms of nuclei and the shared pair why two atoms are held together in a covalent bond.
- Recognise that covalent bonds are directional, leading to simple molecules or giant covalent lattices.
Nature of covalent and dative covalent bonds
Multiple bonds contain multiple pairs of electrons.
- State how many electrons are shared in a double bond and in a triple bond.
- Explain why a carbon-carbon triple bond is shorter and stronger than a carbon-carbon single bond.
- Predict the shape of carbon dioxide by treating each double bond as a single region of electron density.
Nature of covalent and dative covalent bonds
Students should be able to represent: a covalent bond using a line a co-ordinate bond using an arrow.
- Draw the ammonium ion showing three ordinary bonds as lines and the fourth as an arrow in the correct direction.
- Distinguish, in a diagram, between a dative bond arrow and a curly arrow used in a mechanism.
- Show the charge on a polyatomic ion correctly using a square bracket.
Nomenclature
IUPAC rules for nomenclature.
- Name a branched alkane, alkene, alcohol, aldehyde, ketone or carboxylic acid using IUPAC rules.
- Number a carbon chain so that the principal functional group receives the lowest possible locant.
- Deduce a structure from an IUPAC name and check it by renaming the structure you have drawn.
Nomenclature
Organic compounds can be represented by: empirical formula molecular formula general formula structural formula displayed formula skeletal formula.
- Convert between the empirical, molecular, structural, displayed and skeletal representations of a given compound.
- Draw a displayed formula showing every bond and a skeletal formula showing every non-carbon atom.
- Explain why a molecular formula can be ambiguous where a structural formula is not.
Nomenclature
The characteristics of a homologous series, a series of compounds containing the same functional group.
- List the characteristics of a homologous series and illustrate each with the alkanes or the alcohols.
- Write the general formula of a given homologous series and use it to deduce a member's molecular formula.
- Explain the rise in boiling point along a homologous series in terms of London forces.
Nuclear magnetic resonance spectroscopy (A-level only)
Appreciation that scientists have developed a range of analytical techniques which together enable the structures of new compounds to be confirmed.
- State what mass spectrometry, infrared spectroscopy, ¹³C NMR and ¹H NMR each contribute to a structure determination.
- Deduce a structure from combined infrared, ¹³C NMR and ¹H NMR data, quoting a value for every deduction.
- Test a proposed structure against each piece of data given and reject it when one piece disagrees.
- Given a ¹H NMR spectrum with no integration values and no resolved splitting, state which two deductions can no longer be made from it.
Nuclear magnetic resonance spectroscopy (A-level only)
Chemical shift depends on the molecular environment.
- Explain shielding and deshielding in terms of the electron density around a nucleus.
- Predict which of two protons in a molecule appears at the higher chemical shift, and justify the choice.
- Assign a peak at a stated shift to an environment using the Data Booklet.
- Count the proton environments in a given structure and so predict the number of peaks.
Nuclear magnetic resonance spectroscopy (A-level only)
Students should be able to: explain why TMS is a suitable substance to use as a standard 1 13 use H NMR and C NMR spectra and chemical shift data from the Chemistry Data Booklet to suggest possible structures or part structures for molecules use integration data from H NMR spectra to determine the relative numbers of equivalent protons in the molecule use the n+1 rule to deduce the spin–spin splitting patterns of adjacent, non-equivalent protons, limited to doublet, triplet and quartet formation in aliphatic compounds.
- Explain three properties that make tetramethylsilane a suitable standard.
- Deduce the number of protons on a neighbouring carbon from a doublet, triplet or quartet.
- Combine an integration ratio such as 2:2:2:3:3 with the molecular formula to give the actual number of protons in each environment.
- Assemble shift, integration and splitting data into a part structure, then into a full structure that fits the molecular formula.
Nuclear magnetic resonance spectroscopy (A-level only)
The use of the δ scale for recording chemical shift.
- State the reference compound for the delta scale and the shift assigned to it.
- Explain why chemical shift is quoted in parts per million rather than in hertz.
- Use a shift from a spectrum together with the Data Booklet to name the environment responsible.
- Predict whether two spectrometers of different field strength give the same delta value for a peak, and justify the answer.
Nucleophilic properties (A-level only)
Amines are nucleophiles.
- Define a nucleophile and identify the feature of an amine that lets it act as one.
- Explain why the same lone pair accounts for both the basic and the nucleophilic behaviour of an amine.
- Predict which of two amines is the better nucleophile and justify the choice.
Nucleophilic properties (A-level only)
Students should be able to outline the mechanisms of: these nucleophilic substitution reactions the nucleophilic addition–elimination reactions of ammonia and primary amines with acyl chlorides.
- Draw the nucleophilic substitution mechanism of ammonia or an amine with a named halogenoalkane.
- Draw the addition-elimination mechanism of ammonia or a primary amine with a named acyl chloride.
- Explain why the second mole of ammonia or amine is needed and show the step in which it acts.
Nucleophilic properties (A-level only)
The nucleophilic addition–elimination reactions of ammonia and primary amines with acyl chlorides and acid anhydrides.
- Draw and name the amide formed when a named acyl chloride reacts with ammonia or with a primary amine.
- Write a balanced equation that accounts for both moles of the nitrogen compound consumed.
- Predict how the products change when an acid anhydride replaces the acyl chloride.
Nucleophilic properties (A-level only)
The nucleophilic substitution reactions of ammonia and amines with halogenoalkanes to form primary, secondary, tertiary amines and quaternary ammonium salts.
- Write equations for the successive reactions of ammonia with a named halogenoalkane, through to the quaternary salt.
- Explain why a mixture of amines is obtained and how the conditions shift which product dominates.
- Identify the quaternary ammonium salt in a set of structures and say why it cannot be deprotonated.
Nucleophilic properties (A-level only)
The use of quaternary ammonium salts as cationic surfactants.
- Draw a quaternary ammonium ion that would work as a cationic surfactant and label its head and tail.
- Explain how such an ion binds to hair or fabric and what effect the coating has.
- Contrast a cationic surfactant with soap in terms of the charge on the head group.
Nucleophilic substitution
Halogenoalkanes contain polar bonds.
- Label the partial charges on a carbon–halogen bond and justify the direction from electronegativity values.
- Deduce which atom in a given halogenoalkane a nucleophile will attack, and say why.
- Distinguish the polarity argument from the bond enthalpy argument when explaining relative rates of hydrolysis.
Nucleophilic substitution
Halogenoalkanes undergo substitution reactions with the – – nucleophiles OH , CN and NH Students should be able to: outline the nucleophilic substitution mechanisms of these reactions explain why the carbon–halogen bond enthalpy influences the rate of reaction.
- Draw the full mechanism, arrows included, for the reaction of 1-bromopropane with hydroxide, cyanide and ammonia.
- Identify the reagent and conditions needed to convert a named halogenoalkane into a nitrile, an alcohol or an amine.
- Rank the hydrolysis rates of a chloro-, bromo- and iodoalkane using bond enthalpy data and justify the order.
Optical isomerism (A-level only)
A mixture of equal amounts of enantiomers is called a racemic mixture (racemate).
- Decide from the stated proportions of two enantiomers in a sample whether it is a racemate and whether it would rotate plane polarised light.
- Explain why a racemic mixture does not rotate plane polarised light.
- Distinguish between a chiral molecule and an optically active sample using a racemate as the example.
Optical isomerism (A-level only)
An asymmetric carbon atom is chiral and gives rise to optical isomers (enantiomers), which exist as non super- imposable mirror images and differ in their effect on plane polarised light.
- Draw a pair of enantiomers in three dimensions using wedge and dashed bonds either side of a mirror line.
- Sort properties such as melting point, solubility, reaction with an achiral reagent, rotation of plane polarised light and the action of an enzyme into those identical for two enantiomers and those that differ.
- Explain why an enzyme can act on one enantiomer of a substrate but not on the other.
Optical isomerism (A-level only)
Optical isomerism is a form of stereoisomerism and occurs as a result of chirality in molecules, limited to molecules with a single chiral centre.
- Define stereoisomerism and state how optical isomerism differs from E-Z isomerism.
- Identify the chiral centre in a given structure by listing the four groups attached to a carbon.
- Deduce how many optical isomers a molecule with one chiral centre has, and justify the number.
Optical isomerism (A-level only)
Students should be able to: draw the structural formulas and displayed formulas of enantiomers understand how racemic mixtures (racemates) are formed and why they are optically inactive.
- Draw the structural and displayed formulae of both enantiomers of a compound with one chiral centre.
- Explain how a planar intermediate leads to equal amounts of the two enantiomers.
- Explain why the resulting racemate shows no rotation of plane polarised light.
Organic synthesis (A-level only)
Students should be able to: explain why chemists aim to design processes that do not require a solvent and that use non- hazardous starting materials explain why chemists aim to design production methods with fewer steps that have a high percentage atom economy use reactions in this specification to devise a synthesis, with up to four steps, for an organic compound.
- Calculate the percentage atom economy of a reaction from its equation.
- Explain why addition reactions have a higher atom economy than substitution or elimination reactions.
- Justify a solvent-free process by naming two specific advantages over one that uses a solvent.
- Devise a route of up to four steps to a named target, giving reagents, conditions and intermediates.
Organic synthesis (A-level only)
The synthesis of an organic compound can involve several steps.
- Identify the reagents and conditions for each step of a given two-step or three-step route.
- Draw the intermediate formed at each stage of a stated synthesis.
- Calculate the overall percentage yield of a route from the yield of each separate step.
- Explain why a step that produces a mixture of products lowers the yield of the desired compound.
Oxidation of alcohols
Acidified potassium dichromate(VI) is a suitable oxidising agent.
- Name the oxidising agent fully and state the colour change seen when it oxidises an alcohol.
- Write the half equation showing dichromate(VI) reduced to chromium(III) and explain why acid is needed.
- Use [O] to write a balanced equation for the oxidation of a named alcohol.
Oxidation of alcohols
Alcohols are classified as primary, secondary and tertiary.
- Classify a given alcohol as primary, secondary or tertiary by counting alkyl groups on the hydroxyl carbon.
- Draw one primary, one secondary and one tertiary alcohol with the molecular formula C₄H₁₀O.
- Predict from the classification what, if anything, the alcohol is oxidised to by acidified dichromate.
Oxidation of alcohols
Primary alcohols can be oxidised to aldehydes which can be further oxidised to carboxylic acids.
- Write equations using [O] for both oxidation steps of a named primary alcohol.
- Explain how the apparatus chosen determines whether an aldehyde or a carboxylic acid is isolated.
- Describe the colour change observed when acidified potassium dichromate(VI) oxidises a primary alcohol.
Oxidation of alcohols
Secondary alcohols can be oxidised to ketones.
- Write the equation using [O] for the oxidation of a named secondary alcohol and name the product.
- Explain why the ketone formed cannot be oxidised further by acidified potassium dichromate(VI).
- Contrast the products obtained from oxidising a primary and a secondary alcohol under reflux.
Oxidation of alcohols
Students should be able to: write equations for these oxidation reactions (equations showing [O] as oxidant are acceptable) explain how the method used to oxidise a primary alcohol determines whether an aldehyde or carboxylic acid is obtained use chemical tests to distinguish between aldehydes and ketones including Fehling’s solution and Tollens’ reagent.
- Decide, when writing an oxidation equation with [O], whether water is a product: it is when an alcohol is oxidised, but not when an aldehyde becomes a carboxylic acid.
- Find the fault in a given method for distinguishing an aldehyde from a ketone, such as Fehling's solution used without warming, and state the observation it would wrongly produce.
- Describe a test-tube procedure, with observations, that distinguishes propanal from propanone.
Oxidation of alcohols
Tertiary alcohols are not easily oxidised.
- Explain why a tertiary alcohol is not readily oxidised by acidified potassium dichromate(VI).
- Predict and describe the observation when a tertiary alcohol is warmed with the reagent.
- Design a simple test-tube procedure that distinguishes a tertiary alcohol from a secondary one.
Oxidation, reduction and redox equations
Oxidation is the process of electron loss and oxidising agents are electron acceptors.
- Explain why the oxidising agent is the species that is itself reduced, tracing the electrons from the species oxidised to the agent in a given equation.
- Identify the species oxidised and the oxidising agent in a given equation, quoting oxidation states as evidence.
- Explain, with reference to electrons, why a named reaction is classified as redox.
- Deduce the role of concentrated sulfuric acid in its reaction with a solid halide from the products formed.
Oxidation, reduction and redox equations
Reduction is the process of electron gain and reducing agents are electron donors.
- Deduce the number of electrons gained per mole of the species reduced from its change in oxidation state, for example two when sulfuric acid is reduced to sulfur dioxide and eight when it is reduced to hydrogen sulfide.
- Identify the reducing agent in a given equation and justify the choice using oxidation states.
- Explain why iodide is a stronger reducing agent than chloride, using the products of their reactions with concentrated sulfuric acid.
- Deduce from oxidation states that a given reaction is a disproportionation.
Oxidation, reduction and redox equations
Students should be able to: work out the oxidation state of an element in a compound or ion from the formula write half-equations identifying the oxidation and reduction processes in redox reactions combine half-equations to give an overall redox equation.
- Construct a balanced half equation for a species in acidic solution using H₂O, H⁺ and electrons in the correct order.
- Identify which half equation represents oxidation and which represents reduction from the side the electrons appear on.
- Combine two half equations into an overall redox equation by scaling so that the electrons cancel.
- Verify a redox equation by checking that atoms and total charge balance on both sides.
Oxidation, reduction and redox equations
The rules for assigning oxidation states.
- Deduce the oxidation state of an element in a compound or polyatomic ion by applying the rules in order.
- Identify the exceptions for hydrogen and oxygen and apply them to a metal hydride and to a peroxide.
- Calculate the oxidation state of the central atom in an oxoanion from the charge on the ion.
- Explain how a change in oxidation state between two species shows that oxidation or reduction has occurred.
Ozone depletion
Appreciate that results of research by different groups in the scientific community provided evidence for legislation to ban the use of CFCs as solvents and refrigerants.
- Explain why agreement between independent research groups carries more weight than a single result.
- Describe the kinds of evidence that supported restricting chlorofluorocarbons, naming at least one measurement.
- Discuss the role of peer review and publication in turning a chemical hypothesis into international regulation.
Ozone depletion
Chemists have now developed alternative chlorine-free compounds.
- Explain why a replacement refrigerant containing no C–Cl bond cannot deplete ozone.
- Name one chlorine-free refrigerant and state the property that makes it a workable substitute.
- Evaluate a proposed replacement by weighing its ozone safety against flammability or greenhouse effects.
Ozone depletion
Chlorine atoms are formed in the upper atmosphere when ultraviolet radiation causes C–Cl bonds in chlorofluorocarbons (CFCs) to break.
- Write an equation for the ultraviolet fission of a named chlorofluorocarbon, showing the radicals produced.
- Justify, using bond enthalpy data, why the carbon–chlorine bond breaks rather than the carbon–fluorine bond.
- Explain why a compound chosen for being unreactive ends up causing damage in the stratosphere.
Ozone depletion
Chlorine atoms catalyse the decomposition of ozone and contribute to the hole in the ozone layer.
- Write the two propagation steps of the chlorine-catalysed decomposition of ozone and combine them into the overall equation.
- Distinguish the catalyst from the intermediate in the cycle and justify each label.
- Explain why a very small concentration of chlorine radicals destroys a large amount of ozone.
Ozone depletion
Ozone, formed naturally in the upper atmosphere, is beneficial because it absorbs ultraviolet radiation.
- Name a specific consequence of a thinner ozone layer, such as skin cancer, cataracts or damage to crops, and link it to the radiation the layer would otherwise absorb.
- Write equations showing ozone being formed from, and broken back to, oxygen in the upper atmosphere.
- Distinguish the consequences of ozone depletion from the consequences of the enhanced greenhouse effect.
Ozone depletion
Students should be able to: use equations, such as the following, to explain how chlorine atoms catalyse decomposition of ozone:
- Combine two propagation equations into a correct overall equation by cancelling species that appear on both sides.
- Deduce from a pair of equations which species is the catalyst and which is the intermediate.
- Apply the same reasoning to an unfamiliar radical catalyst such as nitrogen monoxide acting on ozone.
pH curves, titrations and indicators (A-level only)
Required practical 9 Investigate how pH changes when a weak acid reacts with a strong base and when a strong acid reacts with a weak base.
- Describe how to obtain a pH curve for ethanoic acid titrated with sodium hydroxide, including calibration of the meter.
- Explain why the volume increments are reduced close to the equivalence point.
- Use a plotted curve to determine the equivalence volume and the pKa of the weak acid.
pH curves, titrations and indicators (A-level only)
Students should be able to: perform calculations for these titrations based on experimental results.
- Calculate the concentration of hydrochloric acid neutralised by 22.40 cm³ of 0.100 mol dm⁻³ barium hydroxide, using the 2 : 1 ratio.
- Convert a titre and a concentration into the mass of solid present in an original sample.
- Estimate the percentage uncertainty in a titre of 22.40 cm³ when each burette reading is uncertain by 0.05 cm³.
pH curves, titrations and indicators (A-level only)
Students should be able to: sketch and explain the shapes of typical pH curves use pH curves to select an appropriate indicator.
- Sketch and explain the shape of the pH curve for the titration of a weak acid with a strong base.
- Select a suitable indicator for the titration of ethanoic acid with sodium hydroxide and justify the choice from the curve.
- Explain, in terms of the steepness of the pH curve, why an indicator cannot be used for butanoic acid titrated with ethylamine.
pH curves, titrations and indicators (A-level only)
Titrations of acids with bases.
- Describe how to make 250 cm³ of a standard solution from a weighed solid, including the washings and final mixing.
- Explain why the funnel is removed from the burette before the titration, referring to the effect on the titre.
- Calculate a mean titre from a set of readings, selecting only the concordant ones.
pH curves, titrations and indicators (A-level only)
Typical pH curves for acid–base titrations in all combinations of weak and strong monoprotic acids and bases.
- Sketch the pH curve for 25.0 cm³ of 0.100 mol dm⁻³ ethanoic acid titrated with 0.100 mol dm⁻³ sodium hydroxide, marking equivalence above pH 7.
- Explain why the equivalence point for ammonia titrated with hydrochloric acid lies below pH 7.
- Explain why the curve for butanoic acid titrated with ethylamine has no steep section.
Physical properties of Period 3 elements
Students should be able to: explain the trends in atomic radius and first ionisation energy explain the melting point of the elements in terms of their structure and bonding.
- Explain why atomic radius decreases across Period 3 in terms of nuclear charge and shielding.
- Write a three-factor explanation for why the first ionisation energy of magnesium exceeds that of sodium.
- Explain why aluminium melts at a higher temperature than sodium, referring to ionic charge, ionic size and delocalised electrons.
Physical properties of Period 3 elements
The trends in atomic radius, first ionisation energy and melting point of the elements Na–Ar The reasons for these trends in terms of the structure of and bonding in the elements.
- Account for the fall in first ionisation energy between magnesium and aluminium in terms of sub-shells.
- Rank sodium, silicon and sulfur by melting point and justify each position from its structure and bonding.
- Predict which element from sodium to argon has the lowest melting point and explain the answer using van der Waals forces between monatomic argon atoms.
Preparation (A-level only)
Aromatic amines, prepared by the reduction of nitro compounds, are used in the manufacture of dyes.
- Write a balanced equation using [H] for the reduction of a named nitro compound to an aromatic amine.
- Explain why excess sodium hydroxide is required at the end of the preparation.
- Outline the two-stage conversion of benzene into phenylamine, giving reagents and conditions for each stage.
Preparation (A-level only)
Primary aliphatic amines can be prepared by the reaction of ammonia with halogenoalkanes and by the reduction of nitriles.
- Write equations for both routes to a named primary aliphatic amine, using [H] for the reduction.
- Explain why the ammonia route produces a mixture and how the conditions are chosen to favour the primary amine.
- Recognise that the two-step synthesis via a nitrile lengthens the carbon chain, while the reduction step itself does not.
Properties of Period 3 elements and their oxides (A-level only)
The reactions of Na and Mg with water.
- Write balanced equations for sodium with cold water and for magnesium with both cold water and steam, naming the different products.
- Predict the approximate pH of each resulting solution and justify the difference using the solubilities of NaOH and Mg(OH)₂.
- Describe what is seen in each reaction, distinguishing a white solid formed in steam from a precipitate formed in solution.
Proteins (A-level only)
Amino acids can be located on a chromatogram using developing agents such as ninhydrin or ultraviolet light and identified by their R values.
- Name a suitable developing agent for amino acids and state why one is required.
- Calculate an Rf value from measurements taken off a chromatogram.
- Explain why an unknown must be run alongside a known standard on the same plate before it can be identified.
Proteins (A-level only)
Amino acids can be separated and identified by thin-layer chromatography.
- Describe the correct setup of a TLC tank, including a pencil baseline above the solvent and a lid to saturate the atmosphere.
- Explain how amino acids are visualised using a chemical locating agent such as ninhydrin.
- Calculate Rf values from a chromatogram and use them to identify amino acids by comparison with known standards.
Proteins (A-level only)
Hydrolysis of the peptide link produces the constituent amino acids.
- Draw the amino acids released when a given di- or tripeptide is hydrolysed, in the form the stated conditions give.
- State the reagent, concentration and conditions used to hydrolyse a protein completely.
- Deduce the number of water molecules consumed when a peptide of a stated length is hydrolysed.
Proteins (A-level only)
Proteins are sequences of amino acids joined by peptide links.
- Draw the structure of a dipeptide or tripeptide formed from named amino acids in a stated order.
- Complete a peptide structure that is only partly drawn, putting the hydrogen on the nitrogen of every peptide link and the right group - acid or amine - at the end of the chain.
- Calculate how many molecules of water are released when a chain of a given length forms.
Proteins (A-level only)
Students should be able to: draw the structure of a peptide formed from up to three amino acids draw the structure of the amino acids formed by hydrolysis of a peptide identify primary, secondary and tertiary structures in diagrams explain how these structures are maintained by hydrogen bonding and S–S bonds calculate Rf values from a chromatogram.
- Deduce how many spots a chromatogram of a fully hydrolysed peptide should show when two of its amino acids are the same, and say what each spot represents.
- Identify where the hydrogen bonds of an α-helix or β-pleated sheet run - between the N-H and C=O of the backbone, not between R groups - and use that to tell secondary from tertiary structure in a diagram.
- Calculate an Rf value from a chromatogram and use it with a table of standards to identify an amino acid.
Proteins (A-level only)
The importance of hydrogen bonding and sulfur–sulfur bonds in proteins.
- Identify from a pair of R groups which interaction, hydrogen bonding or a disulfide bridge, can form between them.
- Explain why a disulfide bridge is a stronger interaction than a hydrogen bond.
- Draw a hydrogen bond between two named R groups, showing the lone pair and the partial charges.
Proteins (A-level only)
The primary, secondary (α-helix and β–pleated sheets) and tertiary structure of proteins.
- Define primary, secondary and tertiary structure and state the bond or force that maintains each.
- Identify the α-helix, the β-pleated sheet and the tertiary fold in a labelled protein diagram.
- Explain why breaking hydrogen bonds changes a protein's shape without changing its primary structure.
Rate equations (A-level only)
The orders m and n are restricted to the values 0, 1, and 2.
- Deduce the order with respect to a reactant from the effect of a concentration change on the rate.
- Write a rate equation from stated orders, omitting any species of zero order.
- Calculate the value of the rate constant from a rate and a set of concentrations and derive its units.
- Explain why the orders cannot be read from the balanced equation for the reaction.
Rate equations (A-level only)
These equations and the gas constant, R, will be given when required.
- Use the given logarithmic form of the Arrhenius equation to determine an activation energy from a gradient.
- Plot ln k against 1/T from tabulated data and draw an appropriate line of best fit.
- Convert an activation energy obtained with R in J K⁻¹ mol⁻¹ into kJ mol⁻¹.
- Predict the effect on k of raising the temperature or of lowering the activation energy, using the form of the equation.
Reaction mechanisms
Free-radical mechanisms: the unpaired electron in a radical is represented by a dot the use of curly arrows is not required for radical mechanisms.
- Draw a radical with the dot placed on the atom that carries the unpaired electron.
- Complete a propagation step by deducing the missing radical, so that the atoms balance and one unpaired electron is shown on each side.
- Distinguish a radical from an ion by showing an unpaired electron rather than a charge.
Reaction mechanisms
Other mechanisms: the formation of a covalent bond is shown by a curly arrow that starts from a lone electron pair or from another covalent bond the breaking of a covalent bond is shown by a curly arrow starting from the bond.
- Draw the bond-forming arrow for a named step, taking it from the lone pair on the oxygen of OH⁻ to the δ⁺ carbon of a haloalkane.
- Draw the bond-breaking arrow for that same step, from the C–Br bond to the bromine, so that Br⁻ leaves with both electrons.
- Identify and correct curly arrows that start from an atom or from a charge.
Reaction mechanisms
Reactions of organic compounds can be explained using mechanisms.
- Name the mechanism of a given organic reaction using both the attacking species and the reaction type, or the single word 'elimination'.
- Distinguish homolytic from heterolytic fission and state the type of intermediate each produces.
- Identify the nucleophile, electrophile or radical in a given reaction step.
Reaction mechanisms
Students should be able to: outline mechanisms by drawing the structures of the species involved and curly arrows to represent the movement of electron pairs.
- Outline a named mechanism by drawing reactant structures, curly arrows, the intermediate and the product.
- Select the correct carbocation intermediate for a stated major product and justify the choice by stability.
- Draw lone pairs in the positions where the mechanism requires an arrow to originate.
Reaction mechanisms
Students should be able to: write balanced equations for the steps in a free- radical mechanism.
- Write balanced initiation, propagation and termination equations for the chlorination of an alkane.
- Classify a given radical equation as initiation, propagation or termination by counting radicals on each side.
- Combine two propagation steps to produce the balanced overall equation for the reaction.
Reactions of ions in aqueous solution (A-level only)
Required practical 11 Carry out simple test-tube reactions to identify transition metal ions in aqueous solution.
- Identify Cu²⁺, Fe²⁺, Fe³⁺, Al³⁺ and Cr³⁺ from their reactions with sodium hydroxide and with ammonia.
- Record separate observations for dropwise addition and for addition of excess reagent.
- Name reagents completely, giving sodium hydroxide rather than hydroxide ion in every answer.
Relative atomic mass and relative molecular mass
Relative atomic mass and relative molecular mass in terms of C.
- Define relative atomic mass as a two-part fraction, naming the reference isotope and the one-twelfth factor.
- Calculate the relative molecular mass of a compound by summing the relative atomic masses in its formula.
- Explain why relative atomic masses have no units and why most are not whole numbers.
Relative atomic mass and relative molecular mass
The term relative formula mass will be used for ionic compounds.
- Calculate the relative formula mass of an ionic compound whose formula contains a bracketed compound ion.
- Explain why the term relative molecular mass is inappropriate for sodium chloride.
- Include water of crystallisation correctly when calculating the relative formula mass of a hydrated salt.
Shapes of complex ions (A-level only)
Cisplatin is the cis isomer.
- Draw the structure of cisplatin and distinguish it from the trans isomer.
- Explain how cisplatin prevents cell division through ligand substitution and binding to guanine in DNA.
- Evaluate the use of cisplatin by weighing its action on cancer cells against damage to healthy dividing cells.
Shapes of complex ions (A-level only)
Square planar complexes are also formed and can display cis–trans isomerism.
- Draw the cis and trans isomers of a square planar complex such as Pt(NH₃)₂Cl₂, including any overall charge.
- Name the type of isomerism shown and justify why it exists in square planar but not in tetrahedral complexes.
- State the bond angle between two ligands in the cis arrangement and in the trans arrangement.
Shapes of complex ions (A-level only)
Transition metal ions commonly form tetrahedral – complexes with larger ligands (eg Cl ).
- Predict the shape and co-ordination number of a complex from the size of its ligand.
- Draw [CuCl₄]²⁻ as a tetrahedral complex and label the 109.5° bond angle.
- Write the equation for the conversion of [Cu(H₂O)₆]²⁺ into [CuCl₄]²⁻ and state the colour change seen.
Shapes of simple molecules and ions
Bonding pairs and lone (non-bonding) pairs of electrons as charge clouds that repel each other.
- Count the bonding pairs and lone pairs around a central atom from a dot-and-cross diagram.
- Explain why a lone pair occupies more space around the central atom than a bonding pair.
- Justify why lone pairs must be included when predicting a molecular shape.
Shapes of simple molecules and ions
Lone pair–lone pair repulsion is greater than lone pair– bond pair repulsion, which is greater than bond pair–bond pair repulsion.
- State the order of the three types of electron pair repulsion and use it to explain a reduced bond angle.
- Predict the bond angle in a four-pair species from the number of lone pairs it carries.
- Explain why the bond angle in water is smaller than that in ammonia.
Shapes of simple molecules and ions
Pairs of electrons in the outer shell of atoms arrange themselves as far apart as possible to minimise repulsion.
- Predict the arrangement adopted by two, three, four, five or six electron pairs around a central atom.
- State the bond angle associated with each of those arrangements.
- Explain why the named shape of a molecule can differ from the arrangement of its electron pairs.
Shapes of simple molecules and ions
Students should be able to: explain the shapes of, and bond angles in, simple molecules and ions with up to six electron pairs (including lone pairs of electrons) surrounding the central atom.
- Calculate the number of bonding and lone pairs around the central atom of a molecular ion, adjusting for its charge.
- Draw and name the shape of a species with up to six electron pairs, including any lone pairs that influence it.
- Explain a predicted shape and bond angle using the relative strengths of electron pair repulsions.
Shapes of simple molecules and ions
The effect of electron pair repulsion on bond angles.
- Predict the bond angle in a molecule or ion from its total electron pairs and its lone pair count.
- Quote both bond angles present in a trigonal bipyramidal species and say which positions they connect.
- Apply the 2.5° angle reduction rule correctly to tetrahedral base geometries without overgeneralising to other shapes.
Structure, bonding and reactivity
Alkenes are unsaturated hydrocarbons.
- Define unsaturated in terms of the carbon–carbon double bond and the ability to undergo addition.
- Describe what is seen when bromine water is shaken with an alkene and with an alkane.
- Deduce from a molecular formula whether a hydrocarbon could contain one carbon–carbon double bond.
Structure, bonding and reactivity
Bonding in alkenes involves a double covalent bond, a centre of high electron density.
- Describe how the π bond in an alkene arises from sideways overlap of p orbitals.
- Explain why alkenes are attacked by electrophiles rather than by nucleophiles.
- Deduce whether a given alkene shows E-Z isomerism, referring to restricted rotation about the double bond.
Substitution reactions (A-level only)
Bidentate and multidentate ligands replace monodentate ligands from complexes. This is called the chelate effect.
- Write balanced equations for the displacement of monodentate ligands by bidentate or multidentate ligands.
- Show that the co-ordination number is unchanged while the number of free particles increases.
- Name the chelate effect and identify it as the reason such reactions go almost to completion.
Substitution reactions (A-level only)
Carbon monoxide is toxic because it replaces oxygen co- ordinately bonded to Fe(II) in haemoglobin.
- Explain the toxicity of carbon monoxide as a ligand substitution at the Fe(II) site of haemoglobin.
- Compare the strength and reversibility of the bonds to carbon monoxide and to oxygen, and deduce the consequence for oxygen transport.
- Identify the donor atom in carbon monoxide and draw the co-ordinate bond that it forms.
Substitution reactions (A-level only)
Haem is an iron(II) complex with a multidentate ligand.
- State the oxidation state of the iron in haem and the co-ordination number of the complex.
- Account for all six co-ordination sites in haem, naming the ligand that occupies each one.
- Distinguish haem from haemoglobin in terms of the globin protein and the number of haem groups present.
Substitution reactions (A-level only)
Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin, enabling oxygen to be transported in the blood.
- Explain oxygen transport as reversible co-ordinate bonding of O₂ at the sixth site on Fe(II).
- Predict whether oxygen will bind or be released from the partial pressure of oxygen in a given tissue.
- State what occupies the sixth co-ordination site when oxygen is not bound.
Substitution reactions (A-level only)
Students should be able to: explain the chelate effect, in terms of the balance between the entropy and enthalpy change in these reactions.
- Explain why ΔH is close to zero for a chelate substitution by comparing the bonds broken and the bonds made.
- Deduce the sign of ΔS from the number of particles on each side of a given equation.
- Use ΔG = ΔH − TΔS to justify why chelate substitutions are feasible and describe them as entropy driven.
The ideal gas equation
Students should be able to: use the equation in calculations.
- Calculate the volume occupied by a known mass of a gas at a stated temperature and pressure, giving the answer in cubic centimetres.
- Determine the relative molecular mass of a gas from its mass, volume, temperature and pressure.
- Identify which conversions are needed before substitution by inspecting the units given in the question.
The ideal gas equation
The ideal gas equation pV = nRT with the variables in SI units.
- Convert pressure, volume and temperature data into the SI units required before substituting into the gas equation.
- Rearrange pV = nRT to make any one of its variables the subject.
- Explain how the units of the gas constant determine the units in which the other quantities must be expressed.
The ionic product of water, Kw (A-level only)
Water is slightly dissociated.
- Write the expression for Kw and state its units.
- Calculate the pH of 0.100 mol dm⁻³ sodium hydroxide at 298 K, obtaining 13.00.
- Explain why water at 323 K has a pH of about 6.6 and is still neutral.
The mole and the Avogadro constant
The Avogadro constant as the number of particles in a mole.
- Calculate the number of ions of a named type in a given mass of an ionic compound.
- Convert the mass of a single atom in grams into a molar mass using the Avogadro constant.
- Explain why one mole of calcium chloride does not contain the same number of positive and negative ions.
The mole and the Avogadro constant
The mole as applied to electrons, atoms, molecules, ions, formulas and equations.
- Calculate the amount in moles of a substance from its mass, or from the concentration and volume of its solution.
- Use the coefficients of a balanced equation to convert an amount of one species into an amount of another.
- Determine the amount of electrons transferred per mole of reactant from a given half-equation.
Trends in properties
Students should be able to explain why: silver nitrate solution is used to identify halide ions the silver nitrate solution is acidified ammonia solution is added.
- Explain why dilute nitric acid is added before the silver nitrate, naming one ion that it removes.
- Justify the choice of nitric acid over hydrochloric and sulfuric acid for acidifying the sample.
- Describe how to distinguish silver chloride from silver bromide when the precipitate colours look similar.
Trends in properties
Students should be able to: explain the trend in electronegativity explain the trend in the boiling point of the elements in terms of their structure and bonding.
- Define electronegativity and use the definition to compare fluorine with bromine.
- Explain the boiling point of bromine in terms of its structure, the force between its molecules and the number of electrons per molecule.
- Predict, with a reason, whether astatine would be a solid, a liquid or a gas at room temperature.
Trends in properties
The trend in oxidising ability of the halogens down the group, including displacement reactions of halide ions in aqueous solution.
- Write the ionic equation and state the colour change when chlorine is bubbled through aqueous sodium iodide.
- Predict, with a reason, whether bromine reacts with aqueous sodium chloride.
- Explain why iodine is a weaker oxidising agent than chlorine, using atomic radius and shielding.
Trends in properties
The trend in reducing ability of the halide ions, including the reactions of solid sodium halides with concentrated sulfuric acid.
- Write the equation for solid sodium chloride with concentrated sulfuric acid and state the observation.
- Deduce the change in oxidation state of sulfur when sodium iodide reduces concentrated sulfuric acid to hydrogen sulfide.
- Explain why sodium bromide gives sulfur dioxide with concentrated sulfuric acid while sodium chloride does not.
Trends in properties
The trend in solubility of the silver halides in ammonia.
- Describe what happens to silver chloride, silver bromide and silver iodide when dilute and then concentrated ammonia are added.
- Explain how the ammonia test distinguishes a white precipitate from a cream one.
- Write the equation for silver chloride dissolving in ammonia to form a soluble complex ion.
Trends in properties
The trends in electronegativity and boiling point of the halogens.
- Explain why chlorine is a gas and iodine a solid at room temperature, in terms of van der Waals forces.
- Explain why the electronegativity of iodine is lower than that of chlorine, referring to atomic radius and shielding.
- State the colour and physical state of chlorine, bromine and iodine at 298 K.
Trends in properties
The use of acidified silver nitrate solution to identify and distinguish between halide ions.
- Describe the colour of the precipitate formed when acidified silver nitrate is added to separate solutions of sodium chloride, bromide and iodide.
- Write the ionic equation, with state symbols, for the formation of silver bromide.
- Explain why dilute nitric acid rather than hydrochloric or sulfuric acid is used to acidify the solution.
Uses of chlorine and chlorate(I)
Appreciate that society assesses the advantages and disadvantages when deciding if chemicals should be added to water supplies.
- Give one benefit and one risk of adding chlorine to drinking water and reach a justified conclusion.
- Explain why the decision to fluoridate a water supply is not one that chemists can take alone.
- Explain why the risk from chlorination depends on the concentration used rather than on the presence of chlorine alone.
Uses of chlorine and chlorate(I)
Appreciate that the benefits to health of water treatment by chlorine outweigh its toxic effects.
- Evaluate the use of chlorine in drinking water, giving a named benefit, a named risk and a conclusion.
- Explain why 'chlorine is toxic' is not on its own a sufficient argument against chlorinating water.
- Use the idea of dose to explain why a substance once used as a chemical weapon is added to drinking water.
Uses of chlorine and chlorate(I)
The reaction of chlorine with cold, dilute, aqueous NaOH and uses of the solution formed.
- Write the balanced equation for chlorine with cold dilute sodium hydroxide and state the conditions required.
- Deduce the two oxidation states of chlorine in the products and use them to show that the reaction is disproportionation.
- Name the compound responsible for the bleaching action of the solution formed and state one household use.
Uses of chlorine and chlorate(I)
The reaction of chlorine with water to form chloride ions and chlorate(I) ions.
- Write the equation for chlorine dissolving in water and name both products.
- Deduce the oxidation state of chlorine in Cl₂, in Cl⁻ and in ClO⁻, and use them to justify the term disproportionation.
- Explain, in terms of electron transfer, why the reaction of chlorine with water is a redox reaction.
Uses of chlorine and chlorate(I)
The reaction of chlorine with water to form chloride ions and oxygen.
- Write the balanced equation for chlorine reacting with water in sunlight.
- Deduce which element is oxidised and which is reduced in 2Cl₂ + 2H₂O → 4H⁺ + 4Cl⁻ + O₂.
- Explain why chlorine water is stored in a dark bottle.
Uses of chlorine and chlorate(I)
The use of chlorine in water treatment.
- State why chlorine is added to drinking water and name one disease that its use prevents.
- Explain why the concentration of chlorine used in water treatment must be carefully controlled.
- Identify the species formed when chlorine dissolves in water that is responsible for killing bacteria.
Variable oxidation states (A-level only)
The redox potential for a transition metal ion changing from a higher to a lower oxidation state is influenced by pH and by the ligand.
- Explain how changing pH alters the electrode potential of a half-equation that includes H⁺.
- Predict the effect of adding ammonia on the ease of oxidation of a cobalt(II) or iron(II) complex.
- Justify why a reaction may occur under given conditions even though standard potentials suggest otherwise.
Variable oxidation states (A-level only)
Transition elements show variable oxidation states.
- Deduce the oxidation state of a transition metal from the formula of any of its compounds or ions.
- Explain variable oxidation states in terms of the similar energies of the 4s and 3d sub-levels.
- Identify which element is oxidised and which is reduced in a given transition metal redox equation.
Variable oxidation states (A-level only)
Vanadium species in oxidation states IV, III and II are formed by the reduction of vanadate(V) ions by zinc in acidic solution.
- Write the half-equation for each reduction step from vanadium(V) to vanadium(II), balancing with H⁺ and H₂O.
- State the colour of each vanadium species and predict the sequence seen as zinc reacts with acidified vanadate(V).
- Combine the vanadium and zinc half-equations into a balanced overall equation.
Weak acids and bases, Ka, for weak acids (A-level only)
Ka is the dissociation constant for a weak acid.
- Write the Ka expression for butanoic acid and state its units.
- Calculate the pH of 0.100 mol dm⁻³ butanoic acid from Ka = 1.51 × 10⁻⁵ mol dm⁻³.
- Convert a pKa of 4.82 into a Ka value and use pKa values to rank three acids by strength.
Weak acids and bases, Ka, for weak acids (A-level only)
Weak acids and weak bases dissociate only slightly in aqueous solution.
- Write the dissociation equation of propanoic acid with an equilibrium sign and identify its conjugate base.
- Sketch mass-loss curves for equal concentrations of chloroethanoic acid and ethanoic acid reacting with excess sodium hydrogencarbonate.
- Explain why 0.100 mol dm⁻³ ethanoic acid has a pH near 2.9 while hydrochloric acid of the same concentration has pH 1.00.
Regulatory & Verification Details
- Qualification Number (QN)
- 601/5731/8
- Regulated Status
- Available to learners
- Awarding Body
- AQA
- Qualification Type
- GCE A Level
- Operational Start
- Sept 2015