Chemistry-9701 A level

Atomic structure

Understand that atoms are mostly empty space surrounding a very small, dense nucleus that contains protons and neutrons; electrons are found in shells in the empty space around the nucleus
Identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses
Understand the terms atomic and proton number; mass and nucleon number
Describe the distribution of mass and charge within an atom
Describe the behaviour of beams of protons, neutrons and electrons moving at the same velocity in an electric field
Determine the numbers of protons, neutrons and electrons present in both atoms and ions given atomic or proton number, mass or nucleon number and charge
State and explain qualitatively the variations in atomic radius and ionic radius across a period and down a group
Define the term isotope in terms of numbers of protons and neutrons
Understand the notation x yA for isotopes, where x is the mass or nucleon number and y is the atomic or proton number
State that and explain why isotopes of the same element have the same chemical properties
State that and explain why isotopes of the same element have different physical properties, limited to mass and density
Understand the terms: • shells, sub-shells and orbitals • principal quantum number (n) • ground state, limited to electronic configuration
Describe the number of orbitals making up s, p and d sub-shells, and the number of electrons that can fill s, p and d sub-shells
Describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p sub-shells
Describe the electronic configurations to include the number of electrons in each shell, sub-shell and orbital
Explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion
Determine the electronic configuration of atoms and ions given the atomic or proton number and charge, using either of the following conventions: e.g. for Fe: 1s2 2s2 2p6 3s2 3p6 3d6 4s2 (full electronic configuration) or [Ar] 3d6 4s2
Understand and use the electrons in boxes notation e.g. for Fe: [Ar]
Describe and sketch the shapes of s and p orbitals
Describe a free radical as a species with one or more unpaired electrons
Define and use the term first ionisation energy, IE
Construct equations to represent first, second and subsequent ionisation energies
Identify and explain the trends in ionisation energies across a period and down a group of the Periodic Table
Identify and explain the variation in successive ionisation energies of an element
Understand that ionisation energies are due to the attraction between the nucleus and the outer electron
Explain the factors influencing the ionisation energies of elements in terms of nuclear charge, atomic/ionic radius, shielding by inner shells and sub-shells and spin-pair repulsion
Deduce the electronic configurations of elements using successive ionisation energy data
Deduce the position of an element in the Periodic Table using successive ionisation energy data

Atoms, molecules and stoichiometry

Define the unified atomic mass unit as one twelfth of the mass of a carbon-12 atom
Define relative atomic mass, Ar , relative isotopic mass, relative molecular mass, Mr , and relative formula mass in terms of the unified atomic mass unit
Define and use the term mole in terms of the Avogadro constant
Write formulas of ionic compounds from ionic charges and oxidation numbers (shown by a Roman numeral), including:
(a) the prediction of ionic charge from the position of an element in the Periodic Table (b) recall of the names and formulas for the following ions: NO3 – , CO3 2–, SO4 2–, OH– , NH4 + , Zn2+, Ag+ , HCO3 – , PO4 3–
(a) write and construct equations (which should be balanced), including ionic equations (which should not include spectator ions) (b) use appropriate state symbols in equations
Define and use the terms empirical and molecular formula
Understand and use the terms anhydrous, hydrated and water of crystallisation
Calculate empirical and molecular formulas, using given data
Perform calculations including use of the mole concept, involving: (a) reacting masses (from formulas and equations) including percentage yield calculations (b) volumes of gases (e.g. in the burning of hydrocarbons)
(c) volumes and concentrations of solutions (d) limiting reagent and excess reagent
Answers should reflect the number of significant figures given or asked for in the question. When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified
(e) deduce stoichiometric relationships from calculations such as those in 2.4.1(a)–(d)

Chemical bonding

Define electronegativity as the power of an atom to attract electrons to itself
Explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells
State and explain the trends in electronegativity across a period and down a group of the Periodic Table
Use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds (the presence of covalent character in some ionic compounds will not be assessed) (Pauling electronegativity values will be given where necessary)
Define ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions)
Describe ionic bonding including the examples of sodium chloride, magnesium oxide and calcium fluoride
Define metallic bonding as the electrostatic attraction between positive metal ions and delocalised electrons
Define covalent bonding as electrostatic attraction between the nuclei of two atoms and a shared pair of electrons (a) describe covalent bonding in molecules including: • hydrogen, H2 • oxygen, O2 • nitrogen, N2
• chlorine, Cl 2 • hydrogen chloride, HCl • carbon dioxide, CO2 • ammonia, NH3 • methane, CH4 • ethane, C2H6 • ethene, C2H4
(b) understand that elements in period 3 can expand their octet including in the compounds sulfur dioxide, SO2, phosphorus pentachloride, PCl 5 , and sulfur hexafluoride, SF6
(c) describe coordinate (dative covalent) bonding, including in the reaction between ammonia and hydrogen chloride gases to form the ammonium ion, NH4 + , and in the Al 2Cl 6 molecule
(a) describe covalent bonds in terms of orbital overlap giving σ and π bonds: • σ bonds are formed by direct overlap of orbitals between the bonding atoms • π bonds are formed by the sideways overlap of adjacent p orbitals above and below the σ bond
(b) describe how the σ and π bonds form in molecules including H₂, C₂H₆, C₂H₄, HCN and N₂ (c) use the concept of hybridisation to describe sp, sp² and sp³ orbitals
(a) define the terms: • bond energy as the energy required to break one mole of a particular covalent bond in the gaseous state • bond length as the internuclear distance of two covalently bonded atoms
(b) use bond energy values and the concept of bond length to compare the reactivity of covalent molecules
State and explain the shapes of, and bond angles in, molecules by using VSEPR theory, including as simple examples: • BF3 (trigonal planar, 120°) • CO2 (linear, 180°) • CH4 (tetrahedral, 109.5°)
• NH3 (pyramidal, 107°) • H2O (non-linear, 104.5°) • SF6 (octahedral, 90°) • PF5 (trigonal bipyramidal, 120° and 90°)
Predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.5.1
Describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples
Use the concept of hydrogen bonding to explain the anomalous properties of H₂O (ice and water): • its relatively high melting and boiling points • its relatively high surface tension • the density of the solid ice compared with the liquid water
Use the concept of electronegativity to explain bond polarity and dipole moments of molecules
(a) describe van der Waals’ forces as the intermolecular forces between molecular entities other than those due to bond formation, and use the term van der Waals’ forces as a generic term to describe all intermolecular forces
Describe the types of van der Waals’ forces: • instantaneous dipole–induced dipole (id-id) forces, also called London dispersion forces • permanent dipole–permanent dipole (pd-pd) forces, including hydrogen bonding
Describe hydrogen bonding and understand that hydrogen bonding is a special case of permanent dipole–permanent dipole forces between molecules where hydrogen is bonded to a highly electronegative atom
State that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces
Use dot-and-cross diagrams to illustrate ionic, covalent and coordinate bonding including the representation of any compounds stated in 3.4 and 3.5
(dot-and-cross diagrams may include species with atoms which have an expanded octet or species with an odd number of electrons)

States of matter

Explain the origin of pressure in a gas in terms of collisions between gas molecules and the wall of the container
Understand that ideal gases have zero particle volume and no intermolecular forces of attraction
State and use the ideal gas equation pV = nRT in calculations, including in the determination of Mr
Describe, in simple terms, the lattice structure of a crystalline solid which is: (a) giant ionic, including sodium chloride and magnesium oxide (b) simple molecular, including iodine, buckminsterfullerene C60 and ice
(c) giant molecular, including silicon(IV) oxide, graphite and diamond (d) giant metallic, including copper
Describe, interpret and predict the effect of different types of structure and bonding on the physical properties of substances, including melting point, boiling point, electrical conductivity and solubility
Deduce the type of structure and bonding present in a substance from given information

Chemical energetics

Understand that chemical reactions are accompanied by enthalpy changes and these changes can be exothermic (ΔH is negative) or endothermic (ΔH is positive
Construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy
Define and use the terms: (a) standard conditions (this syllabus assumes that these are 298K and 101kPa) shown by ⦵ . (b) enthalpy change with particular reference to: reaction, ΔHr , formation, ΔHf , combustion, ΔHc , neutralisation, ΔHneut
Understand that energy transfers occur during chemical reactions because of the breaking and making of chemical bonds
Use bond energies (ΔH positive, i.e. bond breaking) to calculate enthalpy change of reaction, ΔH
Understand that some bond energies are exact and some bond energies are averages
Calculate enthalpy changes from appropriate experimental results, including the use of the relationships q = mcΔT and ΔH = –mcΔT/ n
Apply Hess’s law to construct simple energy cycles
Carry out calculations using cycles and relevant energy terms, including: (a) determining enthalpy changes that cannot be found by direct experiment (b) use of bond energy data

Electrochemistry

Equilibria

(a) understand what is meant by a reversible reaction (b) understand what is meant by dynamic equilibrium in terms of the rate of forward and reverse reactions being equal and the concentration of reactants and products remaining constant
(c) understand the need for a closed system in order to establish dynamic equilibrium
Define Le Chatelier’s principle as: if a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change
Use Le Chatelier’s principle to deduce qualitatively (from appropriate information) the effects of changes in temperature, concentration, pressure or presence of a catalyst on a system at equilibrium
Deduce expressions for equilibrium constants in terms of concentrations, Kc
Use the terms mole fraction and partial pressure
Deduce expressions for equilibrium constants in terms of partial pressures, Kp (use of the relationship between Kp and Kc is not required)
Use the Kc and Kp expressions to carry out calculations (such calculations will not require the solving of quadratic equations)
Calculate the quantities present at equilibrium, given appropriate data
State whether changes in temperature, concentration or pressure or the presence of a catalyst affect the value of the equilibrium constant for a reaction
Describe and explain the conditions used in the Haber process and the Contact process, as examples of the importance of an understanding of dynamic equilibrium in the chemical industry and the application of Le Chatelier’s principle
State the names and formulas of the common acids, limited to hydrochloric acid, HCl, sulfuric acid, H2SO4, nitric acid, HNO3, ethanoic acid, CH3COOH
State the names and formulas of the common alkalis, limited to sodium hydroxide, NaOH, potassium hydroxide, KOH, ammonia, NH3
Describe the Brønsted–Lowry theory of acids and bases
Describe strong acids and strong bases as fully dissociated in aqueous solution and weak acids and weak bases as partially dissociated in aqueous solution
Appreciate that water has pH of 7, acid solutions pH of below 7 and alkaline solutions pH of above 7
Explain qualitatively the differences in behaviour between strong and weak acids including the reaction with a reactive metal and difference in pH values by use of a pH meter, universal indicator or conductivity
Understand that neutralisation reactions occur when H+ (aq) and OH– (aq) form H2O(l)
Understand that salts are formed in neutralisation reactions
Sketch the pH titration curves of titrations using combinations of strong and weak acids with strong and weak alkalis
Select suitable indicators for acid-alkali titrations, given appropriate data (pKa values will not be used)

Reaction kinetics

Explain and use the term rate of reaction, frequency of collisions, effective collisions and non-effective collisions
Explain qualitatively, in terms of frequency of effective collisions, the effect of concentration and pressure changes on the rate of a reaction
Use experimental data to calculate the rate of a reaction
Define activation energy, EA, as the minimum energy required for a collision to be effective
Sketch and use the Boltzmann distribution to explain the significance of activation energy
Explain qualitatively, in terms both of the Boltzmann distribution and of frequency of effective collisions, the effect of temperature change on the rate of a reaction
Explain and use the terms catalyst and catalysis: (a) explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower activation energy
(b) explain this catalytic effect in terms of the Boltzmann distribution (c) construct and interpret a reaction pathway diagram, for a reaction in the presence and absence of an effective catalyst

The Periodic Table: chemical periodicity

Describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements
Explain the variation in melting point and electrical conductivity in terms of the structure and bonding of the elements
Describe, and write equations for, the reactions of the elements with oxygen (to give Na2O, MgO, Al 2O3, P4O10, SO2), chlorine (to give NaCl, MgCl 2, AlCl 3, SiCl 4, PCl 5) and water (Na and Mg only)
State and explain the variation in the oxidation number of the oxides (Na2O, MgO, Al 2O3, P4O10, SO2 and SO3 only) and chlorides (NaCl, MgCl 2, AlCl 3, SiCl 4, PCl 5 only) in terms of their outer shell (valence shell) electrons
Describe, and write equations for, the reactions, if any, of the oxides Na2O, MgO, Al 2O3, SiO2, P4O10, SO2 and SO3 with water including the likely pHs of the solutions obtained
Describe, explain, and write equations for, the acid/base behaviour of the oxides Na2O, MgO, Al 2O3, P4O10, SO2 and SO3 and the hydroxides NaOH, Mg(OH)2 and Al(OH)3 including, where relevant, amphoteric behaviour in reactions with acids and bases NaOH
Describe, explain, and write equations for, the reactions of the chlorides NaCl, MgCl 2, Al Cl 3, SiCl 4, PCl 5 with water including the likely pHs of the solutions obtained
Explain the variations and trends in 9.2.2, 9.2.3, 9.2.4 and 9.2.5 in terms of bonding and electronegativity
Suggest the types of chemical bonding present in the chlorides and oxides from observations of their chemical and physical properties
Predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity
Deduce the nature, possible position in the Periodic Table and identity of unknown elements from given information about physical and chemical properties

Group 2

Group 17

Describe the colours and the trend in volatility of chlorine, bromine and iodine
Describe and explain the trend in the bond strength of the halogen molecules
Interpret the volatility of the elements in terms of instantaneous dipole–induced dipole forces
Describe the relative reactivity of the elements as oxidising agents
Describe the reactions of the elements with hydrogen and explain their relative reactivity in these reactions
Describe the relative thermal stabilities of the hydrogen halides and explain these in terms of bond strengths
Describe the relative reactivity of halide ions as reducing agents
Describe and explain the reactions of halide ions with: (a) aqueous silver ions followed by aqueous ammonia (the formation and formula of the [Ag(NH3) 2] + complex is not required) (b) concentrated sulfuric acid, to include balanced chemical equations
Describe and interpret, in terms of changes in oxidation number, the reaction of chlorine with cold and with hot aqueous sodium hydroxide and recognise these as disproportionation reactions
Explain, including by use of an equation, the use of chlorine in water purification to include the production of the active species HOCl and ClO– which kill bacteria

Nitrogen and sulfur

An introduction to AS Level organic chemistry

Define the term hydrocarbon as a compound made up of C and H atoms only
Understand that alkanes are simple hydrocarbons with no functional group
Understand that the compounds in the table on pages 29 and 30 contain a functional group which dictates their physical and chemical properties
Interpret and use the general, structural, displayed and skeletal formulas of the classes of compound stated in the table on pages 29 and 30
Understand and use systematic nomenclature of simple aliphatic organic molecules with functional groups detailed in the table on pages 29 and 30, up to six carbon atoms (six plus six for esters, straight chains only for esters and nitriles)
Deduce the molecular and/or empirical formula of a compound, given its structural, displayed or skeletal formula
Interpret and use the following terminology associated with types of organic compounds and reactions: (a) homologous series (b) saturated and unsaturated (c) homolytic and heterolytic fission (d) free radical, initiation, propagation, termination
(e) nucleophile, electrophile, nucleophilic, electrophilic (f) addition, substitution, elimination, hydrolysis, condensation (g) oxidation and reduction
Understand and use the following terminology associated with types of organic mechanisms: (a) free-radical substitution (b) electrophilic addition (c) nucleophilic substitution (d) nucleophilic addition
Describe and explain the shape of, and bond angles in, molecules containing sp, sp2 and sp3 hybridised atoms
Describe the arrangement of σ and π bonds in molecules containing sp, sp2 and sp3 hybridised atoms
Understand and use the term planar when describing the arrangement of atoms in organic molecules, for example ethene
Describe structural isomerism and its division into chain, positional and functional group isomerism
Describe stereoisomerism and its division into geometrical (cis/trans) and optical isomerism (use of E/Z nomenclature is acceptable but is not required)
Describe geometrical (cis/trans) isomerism in alkenes, and explain its origin in terms of restricted rotation due to the presence of π bonds
Explain what is meant by a chiral centre and that such a centre gives rise to two optical isomers (enantiomers)
Identify chiral centres and geometrical (cis/trans) isomerism in a molecule of given structural formula including cyclic compounds
Deduce the possible isomers for an organic molecule of known molecular formula

Hydrocarbons

Recall the reactions (reagents and conditions) by which alkanes can be produced: (a) addition of hydrogen to an alkene in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat (b) cracking of a longer chain alkane, heat with Al 2O3
Describe: (a) the complete and incomplete combustion of alkanes (b) the free-radical substitution of alkanes by Cl 2 or Br2 in the presence of ultraviolet light, as exemplified by the reactions of ethane
Describe the mechanism of free-radical substitution with reference to the initiation, propagation and termination steps
Suggest how cracking can be used to obtain more useful alkanes and alkenes of lower Mr from heavier crude oil fractions
Understand the general unreactivity of alkanes, including towards polar reagents in terms of the strength of the C–H bonds and their relative lack of polarity
Recognise the environmental consequences of carbon monoxide, oxides of nitrogen and unburnt hydrocarbons arising from the combustion of alkanes in the internal combustion engine and of their catalytic removal
Recall the reactions (including reagents and conditions) by which alkenes can be produced: (a) elimination of HX from a halogenoalkane by ethanolic NaOH and heat
(b) dehydration of an alcohol, by using a heated catalyst (e.g. Al 2O3) or a concentrated acid (e.g. concentrated H2SO4) (c) cracking of a longer chain alkane
Describe the following reactions of alkenes: (a) the electrophilic addition of (i) hydrogen in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat (ii) steam, H2O(g) and H3PO4 catalyst
(iii) a hydrogen halide, HX(g), at room temperature (iv) a halogen, X2
The oxidation by cold dilute acidified KMnO4 to form the diol
The oxidation by hot concentrated acidified KMnO4 leading to the rupture of the carbon–carbon double bond and the identities of the subsequent products to determine the position of alkene linkages in larger molecules
Addition polymerisation exemplified by the reactions of ethene and propene
Describe the use of aqueous bromine to show the presence of a C=C bond
Describe the mechanism of electrophilic addition in alkenes, using bromine/ethene and hydrogen bromide/propene as examples
Describe and explain the inductive effects of alkyl groups on the stability of primary, secondary and tertiary cations formed during electrophilic addition (this should be used to explain Markovnikov addition)

Halogen compounds

Recall the reactions (reagents and conditions) by which halogenoalkanes can be produced: (a) the free-radical substitution of alkanes by Cl 2 or Br2 in the presence of ultraviolet light, as exemplified by the reactions of ethane
(b) electrophilic addition of an alkene with a halogen, X2, or hydrogen halide, HX(g), at room temperature
(c) substitution of an alcohol, e.g. by reaction with HX(g); or with KCl and concentrated H2SO4 or concentrated H3PO4; or with PCl 3 and heat; or with PCl 5; or with SOCl 2
Classify halogenoalkanes into primary, secondary and tertiary
Describe the following nucleophilic substitution reactions: (a) the reaction with NaOH(aq) and heat to produce an alcohol (b) the reaction with KCN in ethanol and heat to produce a nitrile
(c) the reaction with NH3 in ethanol heated under pressure to produce an amine (d) the reaction with aqueous silver nitrate in ethanol as a method of identifying the halogen present as exemplified by bromoethane
Describe the elimination reaction with NaOH in ethanol and heat to produce an alkene as exemplified by bromoethane
Describe the SN1 and SN2 mechanisms of nucleophilic substitution in halogenoalkanes including the inductive effects of alkyl groups
Recall that primary halogenoalkanes tend to react via the SN2 mechanism; tertiary halogenoalkanes via the SN1 mechanism; and secondary halogenoalkanes by a mixture of the two, depending on structure
Describe and explain the different reactivities of halogenoalkanes (with particular reference to the relative strengths of the C–X bonds as exemplified by the reactions of halogenoalkanes with aqueous silver nitrates)

Hydroxy compounds

Recall the reactions (reagents and conditions) by which alcohols can be produced: (a) electrophilic addition of steam to an alkene, H2O(g) and H3PO4 catalyst (b) reaction of alkenes with cold dilute acidified potassium manganate(VII) to form a dio
(c) substitution of a halogenoalkane using NaOH(aq) and heat (d) reduction of an aldehyde or ketone using NaBH4 or LiAlH4
(e) reduction of a carboxylic acid using LiAlH4 (f) hydrolysis of an ester using dilute acid or dilute alkali and heat
Describe: (a) the reaction with oxygen (combustion) (b) substitution to form halogenoalkanes, e.g. by reaction with HX(g); or with KCl and concentrated H2SO4 or concentrated H3PO4; or with PCl 3 and heat; or with PCl 5; or with SOCl 2
(c) the reaction with Na(s) (d) oxidation with acidified K2Cr2O7 or acidified KMnO4 to: (i) carbonyl compounds by distillation (ii) carboxylic acids by refluxing
(e) dehydration to an alkene, by using a heated catalyst, e.g. Al 2O3 or a concentrated acid (f) formation of esters by reaction with carboxylic acids and concentrated H2SO4 as catalyst as exemplified by ethano
(a) classify alcohols as primary, secondary and tertiary alcohols, to include examples with more than one alcohol group (b) state characteristic distinguishing reactions, e.g. mild oxidation with acidified K2Cr2O7 , colour change from orange to green
Deduce the presence of a CH3CH(OH)– group in an alcohol, CH3CH(OH)–R, from its reaction with alkaline I2(aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO2 –
Explain the acidity of alcohols compared with water

Carbonyl compounds

Recall the reactions (reagents and conditions) by which aldehydes and ketones can be produced: (a) the oxidation of primary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to produce aldehydes
(b) the oxidation of secondary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to produce ketones
Describe: (a) the reduction of aldehydes and ketones using NaBH4 or LiAlH4 to produce alcohols (b) the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat to produce hydroxynitriles as exemplified by ethanal and propanone
Describe the mechanism of the nucleophilic addition reactions of hydrogen cyanide with aldehydes and ketones in 17.1.2(b)
Describe the use of 2,4-dinitrophenylhydrazine (2,4-DNPH reagent) to detect the presence of carbonyl compounds
Deduce the nature (aldehyde or ketone) of an unknown carbonyl compound from the results of simple tests (Fehling’s and Tollens’ reagents; ease of oxidation)
Deduce the presence of a CH3CO– group in an aldehyde or ketone, CH3CO–R, from its reaction with alkaline I2(aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO2

Carboxylic acids and derivatives

Recall the reactions by which carboxylic acids can be produced: (a) oxidation of primary alcohols and aldehydes with acidified K2Cr2O7 or acidified KMnO4 and refluxing
(b) hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification (c) hydrolysis of esters with dilute acid or dilute alkali and heat followed by acidification
Describe: (a) the redox reaction with reactive metals to produce a salt and H2(g) (b) the neutralisation reaction with alkalis to produce a salt and H2O(l )
(c) the acid–base reaction with carbonates to produce a salt and H2O(l) and CO2(g) (d) esterification with alcohols with concentrated H2SO4 as catalyst (e) reduction by LiAlH4 to form a primary alcohol
Recall the reaction (reagents and conditions) by which esters can be produced: (a) the condensation reaction between an alcohol and a carboxylic acid with concentrated H₂SO₄ as catalyst
Describe the hydrolysis of esters by dilute acid and by dilute alkali and heat

Nitrogen compounds

Polymerisation

Organic synthesis

Analytical techniques

Analyse an infrared spectrum of a simple molecule to identify functional groups (see the Data section for the functional groups required)
Analyse mass spectra in terms of m/e values and isotopic abundances (knowledge of the working of the mass spectrometer is not required)
Calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum
Deduce the molecular mass of an organic molecule from the molecular ion peak in a mass spectrum
Suggest the identity of molecules formed by simple fragmentation in a given mass spectrum
Deduce the number of carbon atoms, n, in a compound using the [M + 1]+ peak and the formula n = 100 × abundance of [M + 1]+ ion 1.1 × abundance of M+ ion
Deduce the presence of bromine and chlorine atoms in a compound using the [M + 2]+ peak

Chemical energetics

Define and use the terms: (a) enthalpy change of atomisation, ΔHat (b) lattice energy, ΔHlatt (the change from gas phase ions to solid lattice)
(a) define and use the term first electron affinity, EA (b) explain the factors affecting the electron affinities of elements (c) describe and explain the trends in the electron affinities of the Group 16 and Group 17 elements
Construct and use Born–Haber cycles for ionic solids (limited to +1 and +2 cations, –1 and –2 anions)
Carry out calculations involving Born–Haber cycles
Explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy
Define and use the term enthalpy change with reference to hydration, ΔHhyd, and solution, ΔHso
Construct and use an energy cycle involving enthalpy change of solution, lattice energy and enthalpy change of hydration
Carry out calculations involving the energy cycles in 23.2.2
Explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of an enthalpy change of hydration
Define the term entropy, S, as the number of possible arrangements of the particles and their energy in a given system
Predict and explain the sign of the entropy changes that occur: (a) during a change in state, e.g. melting, boiling and dissolving (and their reverse
(b) during a temperature change (c) during a reaction in which there is a change in the number of gaseous molecules
Calculate the entropy change for a reaction, ΔS, given the standard entropies, S⦵ , of the reactants and products, ΔS⦵ = ΣS⦵ (products) – ΣS⦵ (reactants) (use of ΔS⦵ = ΔSsurr + ΔSsys is not required)
State and use the Gibbs equation ΔG⦵ = ΔH⦵ – TΔS⦵
Perform calculations using the equation ΔG⦵ = ΔH⦵ – TΔS⦵
State whether a reaction or process will be feasible by using the sign of ΔG
Predict the effect of temperature change on the feasibility of a reaction, given standard enthalpy and entropy changes

Electrochemistry

State and apply the relationship F = Le between the Faraday constant, F, the Avogadro constant, L, and the charge on the electron, e
Calculate: (a) the quantity of charge passed during electrolysis, using Q = It (b) the mass and/or volume of substance liberated during electrolysis
Predict the identities of substances liberated during electrolysis from the state of electrolyte (molten or aqueous), position in the redox series (electrode potential) and concentration
Describe the determination of a value of the Avogadro constant by an electrolytic method
Define the terms: (a) standard electrode (reduction) potential (b) standard cell potentiaL
Describe the standard hydrogen electrode
Describe methods used to measure the standard electrode potentials of: (a) metals or non-metals in contact with their ions in aqueous solution (b) ions of the same element in different oxidation states
Calculate a standard cell potential by combining two standard electrode potentials
Use standard cell potentials to: (a) deduce the polarity of each electrode and hence explain/deduce the direction of electron flow in the external circuit of a simple cell (b) predict the feasibility of a reaction
Deduce from E ⦵ values the relative reactivity of elements, compounds and ions as oxidising agents or as reducing agents
Construct redox equations using the relevant half-equations
Predict qualitatively how the value of an electrode potential, E, varies with the concentrations of the aqueous ions
Sse the Nernst equation, e.g. E = E ⦵ + (0.059/z) log [oxidised species]/ [reduced species] ,
to predict quantitatively how the value of an electrode potential varies with the concentrations of the aqueous ions; examples include Cu2+(aq) + 2e– ⇌ Cu(s), Fe3+(aq) + e– ⇌ Fe2+(aq)
Understand and use the equation ΔG⦵ = –nE ⦵ cell F

Equilibria

Understand and use the terms conjugate acid and conjugate base
Define conjugate acid–base pairs, identifying such pairs in reactions
Define mathematically the terms pH, Ka, pKa and Kw and use them in calculations (Kb and the equation Kw = Ka × Kb will not be tested)
Calculate [H+ (aq)] and pH values for: (a) strong acids (b) strong alkalis (c) weak acids
(a) define a buffer solution (b) explain how a buffer solution can be made
(c) explain how buffer solutions control pH; use chemical equations in these explanations (d) describe and explain the uses of buffer solutions, including the role of HCO3 – in controlling pH in blood
Calculate the pH of buffer solutions, given appropriate data
Understand and use the term solubility product, Ksp
Write an expression for Ksp
Calculate Ksp from concentrations and vice versa
(a) understand and use the common ion effect to explain the different solubility of a compound in a solution containing a common ion (b) perform calculations using Ksp values and concentration of a common ion
State what is meant by the term partition coefficient, Kpc
Calculate and use a partition coefficient for a system in which the solute is in the same physical state in the two solvents
Understand the factors affecting the numerical value of a partition coefficient in terms of the polarities of the solute and the solvents used

Reaction kinetics

Explain and use the terms rate equation, order of reaction, overall order of reaction, rate constant, half-life, rate-determining step and intermediate
(a) understand and use rate equations of the form rate = k [A]m[B]n (for which m and n are 0, 1 or 2)
(c) interpret experimental data in graphical form, including concentration–time and rate–concentration graphs (d) calculate an initial rate using concentration data(e) construct a rate equation
(a) show understanding that the half-life of a first-order reaction is independent of concentration (b) use the half-life of a first-order reaction in calculations
Calculate the numerical value of a rate constant, for example by: (a) using the initial rates and the rate equation (b) using the half-life, t ½ , and the equation k = 0.693/t ½
For a multi-step reaction: (a) suggest a reaction mechanism that is consistent with the rate equation and the equation for the overall reaction (b) predict the order that would result from a given reaction mechanism and rate-determining step
(c) deduce a rate equation using a given reaction mechanism and rate-determining step for a given reaction (d) identify an intermediate or catalyst from a given reaction mechanism
Describe qualitatively the effect of temperature change on the rate constant and hence the rate of a reaction
Explain that catalysts can be homogeneous or heterogeneous
Describe the mode of action of a heterogeneous catalyst to include adsorption of reactants, bond weakening and desorption of products, for example:
(a) iron in the Haber process (b) palladium, platinum and rhodium in the catalytic removal of oxides of nitrogen from the exhaust gases of car engines
Describe the mode of action of a homogeneous catalyst by being used in one step and reformed in a later step, for example: (a) atmospheric oxides of nitrogen in the oxidation of atmospheric sulfur dioxide (b) Fe2+ or Fe3+ in the I – /S2O8 2– reaction

Group 2

Chemistry of transition elements

Define a transition element as a d-block element which forms one or more stable ions with incomplete d orbitals
Sketch the shape of a 3dxy orbital and 3dz² orbital
Understand that transition elements have the following properties: (a) they have variable oxidation states (b) they behave as catalysts (c) they form complex ions (d) they form coloured compounds
Explain why transition elements have variable oxidation states in terms of the similarity in energy of the 3d and the 4s sub-shells
Explain why transition elements behave as catalysts in terms of having more than one stable oxidation state, and vacant d orbitals that are energetically accessible and can form dative bonds with ligands
Explain why transition elements form complex ions in terms of vacant d orbitals that are energetically accessible
Describe and explain the reactions of transition elements with ligands to form complexes, including the complexes of copper(II) and cobalt(II) ions with water and ammonia molecules and hydroxide and chloride ions
Define the term ligand as a species that contains a lone pair of electrons that forms a dative covalent bond to a central metal atom/ion
Understand and use the terms: (a) monodentate ligand including as examples H2O, NH3, Cl – and CN– (b) bidentate ligand including as examples 1,2-diaminoethane, en, H2NCH2CH2NH2 and the ethanedioate ion, C2O4 2–
(c) polydentate ligand including as an example EDTA4–
Define the term complex as a molecule or ion formed by a central metal atom/ion surrounded by one or more ligands
Describe the geometry (shape and bond angles) of transition element complexes which are linear, square planar, tetrahedral or octahedral
(a) state what is meant by coordination number (b) predict the formula and charge of a complex ion, given the metal ion, its charge or oxidation state, the ligand and its coordination number or geometry
Explain qualitatively that ligand exchange can occur, including the complexes of copper(II) ions and cobalt(II) ions with water and ammonia molecules and hydroxide and chloride ions
Predict, using E ⦵ values, the feasibility of redox reactions involving transition elements and their ions
Describe the reactions of, and perform calculations involving: (a) MnO4 – /C2O4 2– in acid solution given suitable data (b) MnO4 – /Fe2+ in acid solution given suitable data (c) Cu2+ / I – given suitable data
Perform calculations involving other redox systems given suitable data
Define and use the terms degenerate and non-degenerate d orbitals
Describe the splitting of degenerate d orbitals into two non-degenerate sets of d orbitals of higher energy, and use of ΔE in: (a) octahedral complexes, two higher and three lower d orbitals (b) tetrahedral complexes, three higher and two lower d orbitals
Explain why transition elements form coloured compounds in terms of the frequency of light absorbed as an electron is promoted between two non-degenerate d orbitals
Describe, in qualitative terms, the effects of different ligands on ΔE, frequency of light absorbed, and hence the complementary colour that is observed
Use the complexes of copper(II) ions and cobalt(II) ions with water and ammonia molecules and hydroxide and chloride ions as examples of ligand exchange affecting the colour observed
Describe the types of stereoisomerism shown by complexes, including those associated with bidentate ligands:
(a) geometrical (cis/trans) isomerism, e.g. square planar such as [Pt(NH₃)₂Cl₂] and octahedral such as [Co(NH3) 4(H2O)2] 2+ and [Ni(H2NCH2CH2NH2) 2(H2O)2] 2+ (b) optical isomerism, e.g. [Ni(H2NCH2CH2NH2) 3] 2+ and [Ni(H2NCH2CH2NH2) 2(H2O)2] 2+
Deduce the overall polarity of complexes such as those described in 28.4.1(a) and 28.4.1(b)
Define the stability constant, Kstab, of a complex as the equilibrium constant for the formation of the complex ion in a solvent (from its constituent ions or molecules)
Write an expression for a Kstab of a complex ([H₂O] should not be included)
Use Kstab expressions to perform calculations
Describe and explain ligand exchanges in terms of Kstab values and understand that a large Kstab is due to the formation of a stable complex ion

An introduction to A Level organic chemistry

Understand that the compounds in the table on page 47 contain a functional group which dictates their physical and chemical properties
Interpret and use the general, structural, displayed and skeletal formulas of the classes of compound stated in the table on page 47
Understand and use systematic nomenclature of simple aliphatic organic molecules (including cyclic compounds containing a single ring of up to six carbon atoms) with functional groups detailed in the table on page 47, up to six carbon atoms
Understand and use systematic nomenclature of simple aromatic molecules with one benzene ring and one or more simple substituents, for example 3-nitrobenzoic acid or 2,4,6-tribromophenol
Understand that enantiomers have identical physical and chemical properties apart from their ability to rotate plane polarised light and their potential biological activity
Understand and use the terms optically active and racemic mixture
Describe the effect on plane polarised light of the two optical isomers of a single substance
explain the relevance of chirality to the synthetic preparation of drug molecules including: (a) the potential different biological activity of the two enantiomers (b) the need to separate a racemic mixture into two pure enantiomers
(c) the use of chiral catalysts to produce a single pure optical isomer

Hydrocarbons

Describe the chemistry of arenes as exemplified by the following reactions of benzene and methylbenzene: (a) substitution reactions with Cl₂ and with Br₂ in the presence of a catalyst, AlCl₃ or AlBr₃, to form halogenoarenes (aryl halides)
(b) nitration with a mixture of concentrated HNO₃ and concentrated H₂SO₄ at a temperature between 25°C and 60°C (c) Friedel–Crafts alkylation by CH₃Cl and AlCl₃ and heat
(d) Friedel–Crafts acylation by CH₃COCl and AlCl₃ and heat (e) complete oxidation of the side-chain using hot alkaline KMnO₄ and then dilute acid to give a benzoic acid
Describe the mechanism of electrophilic substitution in arenes: (a) as exemplified by the formation of nitrobenzene and bromobenzene
(b) with regards to the effect of delocalisation (aromatic stabilisation) of electrons in arenes to explain the predomination of substitution over addition
Predict whether halogenation will occur in the side-chain or in the aromatic ring in arenes depending on reaction conditions
Describe that in the electrophilic substitution of arenes, different substituents direct to different ring positions (limited to the directing effects of –NH₂, –OH, –R, –NO₂, –COOH and –COR)

Halogens Compounds

Hydroxy Compounds

Describe the reaction with acyl chlorides to form esters using ethyl ethanoate
Recall the reactions (reagents and conditions) by which phenol can be produced: (a) reaction of phenylamine with HNO₂ or NaNO₂ and dilute acid below 10°C to produce the diazonium salt; further warming of the diazonium salt with H₂O to give phenol
Recall the chemistry of phenol, as exemplified by the following reactions: (a) with bases, for example NaOH(aq) to produce sodium phenoxide (b) with Na(s) to produce sodium phenoxide and H2(g) (c) in NaOH(aq) with diazonium salts, to give azo compound
(d) nitration of the aromatic ring with dilute HNO₃(aq) at room temperature to give a mixture of 2-nitrophenol and 4-nitrophenol (e) bromination of the aromatic ring with Br₂(aq) to form 2,4,6-tribromophenol
Explain the acidity of phenol
Describe and explain the relative acidities of water, phenol and ethanol
Explain why the reagents and conditions for the nitration and bromination of phenol are different from those for benzene
Recall that the hydroxyl group of a phenol directs to the 2-, 4- and 6-positions
Apply knowledge of the reactions of phenol to those of other phenolic compounds, e.g. naphthol

Carboxylic acids and derivatives

Recall the reaction by which benzoic acid can be produced: (a) reaction of an alkylbenzene with hot alkaline KMnO₄ and then dilute acid, exemplified by methylbenzene
Describe the reaction of carboxylic acids with PCl₃ and heat, PCl₅ or SOCl₂ to form acyl chlorides
Recognise that some carboxylic acids can be further oxidised: (a) the oxidation of methanoic acid, HCOOH, with Fehling’s reagent or Tollens’ reagent or acidified KMnO₄ or acidified K₂Cr₂O₇ to carbon dioxide and water
(b) the oxidation of ethanedioic acid, HOOCCOOH, with warm acidified KMnO₄ to carbon dioxide
Describe and explain the relative acidities of carboxylic acids, phenols and alcohols
Describe and explain the relative acidities of chlorine-substituted carboxylic acids
Recall the reaction by which esters can be produced: (a) reaction of alcohols with acyl chlorides using the formation of ethyl ethanoate and phenyl benzoate as examples
Recall the reactions (reagents and conditions) by which acyl chlorides can be produced: (a) reaction of carboxylic acids with PCl₃ and heat, PCl₅ or SOCl₂
Describe the following reactions of acyl chlorides: (a) hydrolysis on addition of water at room temperature to give the carboxylic acid and HCl (b) reaction with an alcohol at room temperature to produce an ester and HCl
(c) reaction with phenol at room temperature to produce an ester and HCl (d) reaction with ammonia at room temperature to produce an amide and HCl (e) reaction with a primary or secondary amine at room temperature to produce an amide and HCl
Describe the addition–elimination mechanism of acyl chlorides in reactions in 33.3.2(a)–(e)
Explain the relative ease of hydrolysis of acyl chlorides, alkyl chlorides and halogenoarenes (aryl chlorides)

Nitrogen compounds

Recall the reactions (reagents and conditions) by which primary and secondary amines are produced: (a) reaction of halogenoalkanes with NH₃ in ethanol heated under pressure (b) reaction of halogenoalkanes with primary amines in ethanol
(c) the reduction of amides with LiAlH₄ (d) the reduction of nitriles with LiAlH₄ or H₂/Ni
Describe the condensation reaction of ammonia or an amine with an acyl chloride at room temperature to give an amide
Describe and explain the basicity of aqueous solutions of amines
Describe the preparation of phenylamine via the nitration of benzene to form nitrobenzene followed by reduction with hot Sn/concentrated HCl followed by NaOH(aq)
Describe: (a) the reaction of phenylamine with Br₂(aq) at room temperature (b) the reaction of phenylamine with HNO₂ or NaNO₂ and dilute acid below 10°C to produce the diazonium salt; further warming of the diazonium salt with H₂O to give phenol
Describe and explain the relative basicities of aqueous ammonia, ethylamine and phenylamine
Recall the following about azo compounds: (a) describe the coupling of benzenediazonium chloride with phenol in NaOH(aq) to form an azo compound
(b) identify the azo group (c) state that azo compounds are often used as dyes (d) that other azo dyes can be formed via a similar route
Recall the reactions (reagents and conditions) by which amides are produced: (a) the reaction between ammonia and an acyl chloride at room temperature (b) the reaction between a primary amine and an acyl chloride at room temperature
Describe the reactions of amides: (a) hydrolysis with aqueous alkali or aqueous acid (b) the reduction of the CO group in amides with LiAlH₄ to form an amine
State and explain why amides are much weaker bases than amines
Describe the acid/base properties of amino acids and the formation of zwitterions, to include the isoelectric point
Describe the formation of amide (peptide) bonds between amino acids to give di- and tripeptides
Interpret and predict the results of electrophoresis on mixtures of amino acids and dipeptides at varying pHs (the assembling of the apparatus will not be tested)

Polymerisation

Describe the formation of polyesters: (a) the reaction between a diol and a dicarboxylic acid or dioyl chloride (b) the reaction of a hydroxycarboxylic acid
Describe the formation of polyamides: (a) the reaction between a diamine and a dicarboxylic acid or dioyl chloride (b) the reaction of an aminocarboxylic acid (c) the reaction between amino acids
Deduce the repeat unit of a condensation polymer obtained from a given monomer or pair of monomers
Identify the monomer(s) present in a given section of a condensation polymer molecule
Predict the type of polymerisation reaction for a given monomer or pair of monomers
Deduce the type of polymerisation reaction which produces a given section of a polymer molecule
Recognise that poly(alkenes) are chemically inert and can therefore be difficult to biodegrade
Recognise that some polymers can be degraded by the action of light
Recognise that polyesters and polyamides are biodegradable by acidic and alkaline hydrolysis

Organic synthesis

Analytical techniques

Describe and understand the terms (a) stationary phase, for example aluminium oxide (on a solid support) (b) mobile phase; a polar or non-polar solvent (c) Rf value (d) solvent front and baseline
Interpret Rf values
Explain the differences in Rf values in terms of interaction with the stationary phase and of relative solubility in the mobile phase
Describe and understand the terms (a) stationary phase; a high boiling point non-polar liquid (on a solid support) (b) mobile phase; an unreactive gas (c) retention time
Interpret gas /liquid chromatograms in terms of the percentage composition of a mixture
Explain retention times in terms of interaction with the stationary phase
Analyse and interpret a carbon-13 NMR spectrum of a simple molecule to deduce: (a) the different environments of the carbon atoms present (b) the possible structures for the molecule
Predict or explain the number of peaks in a carbon-13 NMR spectrum for a given molecule
Analyse and interpret a proton (1 H) NMR spectrum of a simple molecule to deduce: (a) the different environments of proton present using chemical shift values (b) the relative numbers of each type of proton present from relative peak areas
(c) the number of equivalent protons on the carbon atom adjacent to the one to which the given proton is attached from the splitting pattern, using the n + 1 rule (d) the possible structures for the molecule
Predict the chemical shifts and splitting patterns of the protons in a given molecule
Describe the use of tetramethylsilane, TMS, as the standard for chemical shift measurements
State the need for deuterated solvents, e.g. CDCl₃, when obtaining a proton NMR spectrum
Describe the identification of O–H and N–H protons by proton exchange using D₂O

Course Content

Atomic structure

Understand that atoms are mostly empty space surrounding a very small, dense nucleus that contains protons and neutrons; electrons are found in shells in the empty space around the nucleus

Identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses

Understand the terms atomic and proton number; mass and nucleon number

Describe the distribution of mass and charge within an atom

Describe the behaviour of beams of protons, neutrons and electrons moving at the same velocity in an electric field

Determine the numbers of protons, neutrons and electrons present in both atoms and ions given atomic or proton number, mass or nucleon number and charge

State and explain qualitatively the variations in atomic radius and ionic radius across a period and down a group

Define the term isotope in terms of numbers of protons and neutrons

Understand the notation x yA for isotopes, where x is the mass or nucleon number and y is the atomic or proton number

State that and explain why isotopes of the same element have the same chemical properties

State that and explain why isotopes of the same element have different physical properties, limited to mass and density

Understand the terms: • shells, sub-shells and orbitals • principal quantum number (n) • ground state, limited to electronic configuration

Describe the number of orbitals making up s, p and d sub-shells, and the number of electrons that can fill s, p and d sub-shells

Describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p sub-shells

Describe the electronic configurations to include the number of electrons in each shell, sub-shell and orbital

Explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion

Determine the electronic configuration of atoms and ions given the atomic or proton number and charge, using either of the following conventions: e.g. for Fe: 1s2 2s2 2p6 3s2 3p6 3d6 4s2 (full electronic configuration) or [Ar] 3d6 4s2

Understand and use the electrons in boxes notation e.g. for Fe: [Ar]

Describe and sketch the shapes of s and p orbitals

Describe a free radical as a species with one or more unpaired electrons

Define and use the term first ionisation energy, IE

Construct equations to represent first, second and subsequent ionisation energies

Identify and explain the trends in ionisation energies across a period and down a group of the Periodic Table

Identify and explain the variation in successive ionisation energies of an element

Understand that ionisation energies are due to the attraction between the nucleus and the outer electron

Explain the factors influencing the ionisation energies of elements in terms of nuclear charge, atomic/ionic radius, shielding by inner shells and sub-shells and spin-pair repulsion

Deduce the electronic configurations of elements using successive ionisation energy data

Deduce the position of an element in the Periodic Table using successive ionisation energy data

Atoms, molecules and stoichiometry

Define the unified atomic mass unit as one twelfth of the mass of a carbon-12 atom

Define relative atomic mass, Ar , relative isotopic mass, relative molecular mass, Mr , and relative formula mass in terms of the unified atomic mass unit

Define and use the term mole in terms of the Avogadro constant

Write formulas of ionic compounds from ionic charges and oxidation numbers (shown by a Roman numeral), including:

(a) the prediction of ionic charge from the position of an element in the Periodic Table (b) recall of the names and formulas for the following ions: NO3 – , CO3 2–, SO4 2–, OH– , NH4 + , Zn2+, Ag+ , HCO3 – , PO4 3–

(a) write and construct equations (which should be balanced), including ionic equations (which should not include spectator ions) (b) use appropriate state symbols in equations

Define and use the terms empirical and molecular formula

Understand and use the terms anhydrous, hydrated and water of crystallisation

Calculate empirical and molecular formulas, using given data

Perform calculations including use of the mole concept, involving: (a) reacting masses (from formulas and equations) including percentage yield calculations (b) volumes of gases (e.g. in the burning of hydrocarbons)

(c) volumes and concentrations of solutions (d) limiting reagent and excess reagent

Answers should reflect the number of significant figures given or asked for in the question. When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified

(e) deduce stoichiometric relationships from calculations such as those in 2.4.1(a)–(d)

Chemical bonding

Define electronegativity as the power of an atom to attract electrons to itself

Explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells

State and explain the trends in electronegativity across a period and down a group of the Periodic Table

Use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds (the presence of covalent character in some ionic compounds will not be assessed) (Pauling electronegativity values will be given where necessary)

Define ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions)

Describe ionic bonding including the examples of sodium chloride, magnesium oxide and calcium fluoride

Define metallic bonding as the electrostatic attraction between positive metal ions and delocalised electrons

Define covalent bonding as electrostatic attraction between the nuclei of two atoms and a shared pair of electrons (a) describe covalent bonding in molecules including: • hydrogen, H2 • oxygen, O2 • nitrogen, N2

• chlorine, Cl 2 • hydrogen chloride, HCl • carbon dioxide, CO2 • ammonia, NH3 • methane, CH4 • ethane, C2H6 • ethene, C2H4

(b) understand that elements in period 3 can expand their octet including in the compounds sulfur dioxide, SO2, phosphorus pentachloride, PCl 5 , and sulfur hexafluoride, SF6

(c) describe coordinate (dative covalent) bonding, including in the reaction between ammonia and hydrogen chloride gases to form the ammonium ion, NH4 + , and in the Al 2Cl 6 molecule

(a) describe covalent bonds in terms of orbital overlap giving σ and π bonds: • σ bonds are formed by direct overlap of orbitals between the bonding atoms • π bonds are formed by the sideways overlap of adjacent p orbitals above and below the σ bond

(b) describe how the σ and π bonds form in molecules including H₂, C₂H₆, C₂H₄, HCN and N₂ (c) use the concept of hybridisation to describe sp, sp² and sp³ orbitals

(a) define the terms: • bond energy as the energy required to break one mole of a particular covalent bond in the gaseous state • bond length as the internuclear distance of two covalently bonded atoms

(b) use bond energy values and the concept of bond length to compare the reactivity of covalent molecules

State and explain the shapes of, and bond angles in, molecules by using VSEPR theory, including as simple examples: • BF3 (trigonal planar, 120°) • CO2 (linear, 180°) • CH4 (tetrahedral, 109.5°)

• NH3 (pyramidal, 107°) • H2O (non-linear, 104.5°) • SF6 (octahedral, 90°) • PF5 (trigonal bipyramidal, 120° and 90°)

Predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.5.1

Describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples

Use the concept of hydrogen bonding to explain the anomalous properties of H₂O (ice and water): • its relatively high melting and boiling points • its relatively high surface tension • the density of the solid ice compared with the liquid water

Use the concept of electronegativity to explain bond polarity and dipole moments of molecules

(a) describe van der Waals’ forces as the intermolecular forces between molecular entities other than those due to bond formation, and use the term van der Waals’ forces as a generic term to describe all intermolecular forces

Describe the types of van der Waals’ forces: • instantaneous dipole–induced dipole (id-id) forces, also called London dispersion forces • permanent dipole–permanent dipole (pd-pd) forces, including hydrogen bonding

Describe hydrogen bonding and understand that hydrogen bonding is a special case of permanent dipole–permanent dipole forces between molecules where hydrogen is bonded to a highly electronegative atom

State that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces

Use dot-and-cross diagrams to illustrate ionic, covalent and coordinate bonding including the representation of any compounds stated in 3.4 and 3.5

(dot-and-cross diagrams may include species with atoms which have an expanded octet or species with an odd number of electrons)

States of matter

Explain the origin of pressure in a gas in terms of collisions between gas molecules and the wall of the container

Understand that ideal gases have zero particle volume and no intermolecular forces of attraction

State and use the ideal gas equation pV = nRT in calculations, including in the determination of Mr

Describe, in simple terms, the lattice structure of a crystalline solid which is: (a) giant ionic, including sodium chloride and magnesium oxide (b) simple molecular, including iodine, buckminsterfullerene C60 and ice

(c) giant molecular, including silicon(IV) oxide, graphite and diamond (d) giant metallic, including copper

Describe, interpret and predict the effect of different types of structure and bonding on the physical properties of substances, including melting point, boiling point, electrical conductivity and solubility