11 AS/A Chemistry Annualplan Vinay

Note: The following plan is for a period of 2 years covering the scope of both A and AS levels. The scope of each topic is at both A and AS levels. In the first year, only the scope of AS levels will be addressed. The following year, in the corresponding month, the scope of A levels will be covered. Some topics are only for A levels which will be mentioned in the following plan

July

Atoms, molecules and stoichiometry

  • Relative masses of atoms and molecules
  • The mole, the Avogadro constant
  • The determination of relative atomic masses, Ar , and relative molecular masses, Mr , from mass spectra
  • The calculation of empirical and molecular formulae
  • Reacting masses and volumes (of solutions and gases)

Learning Outcomes

(a) define and use the terms relative atomic, isotopic, molecular and formula masses, based on the 12C scale
(b) define and use the term molein terms of the Avogadro constant
(c) analyse mass spectra in terms of isotopic abundances and molecular fragments[knowledge of the working of the mass spectrometer is not required]
(d) calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum
(e) define and use the terms empirical and molecular formulae
(f) calculate empirical and molecular formulae, using combustion data or composition by mass
(g) write and/or construct balanced equations
(h) perform calculations, including use of the mole concept, involving:
(i) reacting masses (from formulae and equations)
(ii) volumes of gases (e.g. in the burning of hydrocarbons)
(iii) volumes and concentrations of solutions
When performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question.
When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified
(i) deduce stoichiometric relationships from calculations such as those in (h)

Atomic Structure

  • The nucleus of the atom: neutrons and protons, isotopes, proton and nucleon numbers
  • Electrons: electronic energy levels, ionisation energies, atomic orbitals, extranuclear structure

Learning Outcomes

(a) identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses
(b) deduce the behaviour of beams of protons, neutrons and electrons in electric fields
(c) describe the distribution of mass and charges within an atom
(d) deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge)
(e) (i) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number
(ii) distinguish between isotopes on the basis of different numbers of neutrons present
(iii) recognise and use the symbolism xAyfor isotopes, where x is a the nucleon number and y is the proton number
(f) describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals.
(g) describe the shapes of s and p orbitals and d orbitals
(h) state the electronic configuration of atoms and ions given the proton number (and charge), using the convention 1s2, 2s2, 2p6 etc
(i) (i) explain and use the terms ionisation energy and electron affinity
(ii) explain the factors influencing the ionisation energies of elements
(iii) explain the trends in ionisation energies across a Period and down a Group of the Periodic Table
(j) deduce the electronic configurations of elements from successive ionisation energy data
(k) interpret successive ionisation energy data of an element in terms of the position of that element within the Periodic Table

August

Chemical Bonding

  • Ionic (electrovalent) bonding
  • Covalent bonding and co-ordinate (dative covalent) bonding
    • The shapes of simple molecules
    • Bond energies, bond lengths and bond polarities
  • Intermolecular forces, including hydrogen bonding
  • Metallic bonding
  • Bonding and physical properties

Learning Outcomes

(a) describe ionic (electrovalent) bonding, as in sodium chloride and magnesium oxide, including the use of ‘dot-and-cross’ diagrams
(b) describe, including the use of ‘dot-and-cross’ diagrams,
(i) covalent bonding, as in hydrogen, oxygen, chlorine, hydrogen chloride, carbon dioxide, methane, ethene
(ii) co-ordinate (dative covalent) bonding, as in the formation of the ammonium ion and in the Al2Cl6molecule
(c) explain the shapes of, and bond angles in, molecules by using the qualitative model of electron-pair repulsion (including lone pairs), using as simple examples: BF3(trigonal), CO2(linear), CH4(tetrahedral), NH3(pyramidal), H2O (non-linear), SF6 (octahedral), PF5 (trigonal bipyramid))
(d) describe covalent bonding in terms of orbital overlap, giving σ and πbonds, including the concept of hybridisation to form sp, sp2 and sp3 orbitals
(e) explain the shape of, and bond angles in, the ethane, ethene and benzene molecules in terms of σ and π bonds
(f) predict the shapes of, and bond angles in, molecules analogous to those specified in (c)and (e)
(g) describe hydrogen bonding, using ammonia and water as simple examples of molecules containing N-H and O-H groups
(h) understand, in simple terms, the concept of electronegativity and apply it to explain the properties of molecules such as bond polarity, the dipole moments of molecules, the behaviour of oxides with water and the acidities of chlorine-substituted ethanoic acids
(i) explain the terms bond energy, bond lengthand bond polarityand use them to compare the reactivities of covalent bonds
(j) describe intermolecular forces (van der Waals’ forces), based on permanent and induced dipoles, as in CHCl3 , Br2 and the liquid noble gases
(k) describe metallic bonding in terms of a lattice of positive ions surrounded by mobile electrons
(l) describe, interpret and/or predict the effect of different types of bonding (ionic bonding, covalent bonding, hydrogen bonding, other intermolecular interactions, metallic bonding) on the physical properties of substances
(m) deduce the type of bonding present from given information
(n) show understanding of chemical reactions in terms of energy transfers associated with the breaking and making of chemical bonds

September

States of Matter

  • The gaseous state:
    • Ideal gas behaviour and deviations from it
    • pV =nRT and its use in determining a value for Mr
  • The liquid state
    • The kinetic concept of the liquid state and simple kinetic-molecular descriptions of changes of state
  • The solid state
    • Lattice structures

Learning Outcomes

(a) state the basic assumptions of the kinetic theory as applied to an ideal gas
(b) explain qualitatively in terms of intermolecular forces and molecular size:
(i) the conditions necessary for a gas to approach ideal behaviour
(ii) the limitations of ideality at very high pressures and very low temperatures
(c) state and use the general gas equation pV= nRTin calculations, including the determination of Mr
(d) describe, using a kinetic-molecular model: the liquid state, melting, vaporisation, vapour pressure
(e) describe, in simple terms, the lattice structure of a crystalline solid which is:
(i) ionic, as in sodium chloride, magnesium oxide
(ii) simple molecular, as in iodine
(iii) giant molecular, as in silicon(IV) oxide and the graphite and diamond allotropes of carbon
(iv) hydrogen-bonded, as in ice
(v) metallic, as in copper
[the concept of the ‘unit cell’ is not required]
(f) explain the strength, high melting point and electrical insulating properties of ceramics in terms of their giant molecular structure
(g) relate the uses of ceramics, based on magnesium oxide, aluminium oxide and silicon(IV) oxide, to their properties (suitable examples include furnace linings, electrical insulators, glass, crockery)
(h) discuss the finite nature of materials as a resource and the importance of recycling processes
(i) outline the importance of hydrogen bonding to the physical properties of substances, including ice and water (for example, boiling and melting points, viscosity and surface tension)
(j) suggest from quoted physical data the type of structure and bonding present in a substance

Chemical Energetics

  • Enthalpy changes: ∆Hof formation, combustion, hydration, solution, neutralisation and atomisation; bond energy; lattice energy; electron affinity
  • Hess’ Law, including BornHaber cycles

Learning Outcomes

(a) explain that some chemical reactions are accompanied by energy changes, principally in the form of heat energy; the energy changes can be exothermic (∆H, negative) or endothermic
(b) explain and use the terms:
(i) enthalpy change of reactionand standard conditions, with particular reference to: formation, combustion, hydration, solution, neutralisation, atomisation
(ii) bond energy(∆Hpositive, i.e. bond breaking)
(iii) lattice energy (∆H negative, i.e. gaseous ions to solid lattice)
(c) calculate enthalpy changes from appropriate experimental results, including the use of the relationship enthalpy change, ∆H= –mc∆T
(d) explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy
(e) apply Hess’ Law to construct simple energy cycles, and carry out calculations involving such cycles and relevant energy terms, with particular reference to:
(i) determining enthalpy changes thatcannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion
(ii) average bond energies
(iii) the formation of a simple ionic solid and of its aqueous solution
(iv) Born-Haber cycles (including ionisation energy and electron affinity)
(f) construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy

October

Electrochemistry

  • Redox processes: electron transfer and changes in oxidation number (oxidation state)
  • Electrode potentials
    • Standard electrode (redox) potentials ; the redox series
    • Standard cell potentials, and their uses
    • Batteries and fuel cells
  • Electrolysis
    • Factors affecting the amount of substance liberated during electrolysis
    • The Faraday constant: the Avogadro constant: their relationship
    • Industrial uses of electrolysis

Learning Outcomes

(a)
(i) calculate oxidation numbers of elements in compounds and ions
(ii) describe and explain redox processes in terms of electron transfer and/or changes in oxidation number (oxidation state)
(iii) use changes in oxidation numbers to help balance chemical equations
(b) explain, including the electrode reactions, the industrial processes of:
(i) the electrolysis of brine, using a diaphragm cell
(ii) the extraction of aluminium from molten aluminium oxide/cryolite
(iii) the electrolytic purification of copper
(c) define the terms:
(i) standard electrode (redox) potential
(ii) standard cell potential
(d) describe the standard hydrogen electrode
(e) describe methods used to measure the standard electrode potentials of:
(i) metals or non-metals in contact with their ions in aqueous solution
(ii) ions of the same element in different oxidation states
(f) calculate a standard cell potential by combining two standard electrode potentials
(g) use standard cell potentials to:
(i) explain/deduce the direction of electron flow from a simple cell
(ii) predict the feasibility of a reaction
(h) construct redox equations using the relevant half-equations
(i) predict qualitatively how the value of an electrode potential varies with the concentration of the aqueous ion
(j) state the possible advantages of developing other types of cell, e.g. the H2/O2 fuel cell and improved batteries (as in electric vehicles) in terms of smaller size, lower mass and higher voltage
(k) state the relationship, F = Le, between the Faraday constant, the Avogadro constant and the charge on the electron
(l) predict the identity of the substance liberated during electrolysis from the state of electrolyte (molten or aqueous), position in the redox series (electrode potential) and concentration
(m) calculate:
(i) the quantity of charge passed during electrolysis
(ii) the mass and/or volume of substance liberated during electrolysis, including those in the electrolysis of H2SO4(aq), Na2SO4(aq)
(n) describe the determination of a value of the Avogadro constant by an electrolytic method

Chemical Equilibria

  • Chemical equilibria:reversible reactions; dynamic equilibrium
    • Factors affecting chemical equilibria
    • Equilibrium constants
    • The Haber process; the Contact process
  • Ionic equilibria
    • Brønsted-Lowry theory of acids and bases
    • Acid dissociation constants, Ka and the use of pKa
    • The ionic product of water, Kw
    • pH: choice of pH indicators
    • Buffer solutions
    • Solubility product; the common ion effect

Learning Outcomes

(a) explain, in terms of rates of the forward and reverse reactions, what is meant by a reversible reactionand dynamic equilibrium
(b) state Le Chatelier’s Principle and apply it to deduce qualitatively (from appropriate information) the effects of changes in temperature, concentration or pressure, on a system at equilibrium
(c) state whether changes in concentration, pressure or temperature or the presence of a catalyst affect the value of the equilibrium constant for a reaction
(d) deduce expressions for equilibrium constants in terms of concentrations, Kc, and partial pressures, Kp
(e) calculate the values of equilibrium constants in terms of concentrations or partial pressures from appropriate data
(f) calculate the quantities present at equilibrium, given appropriate data
(g) describe and explain the conditions used in the Haber process and the Contact process, as examples of the importance of an understanding of chemical equilibrium in the chemical industry
(h) show understanding of, and use, the Brønsted-Lowry theory of acids and bases, including the use of the acid-I, base-II concept
(i) explain qualitatively the differences in behaviour between strong and weak acids and bases and the pH values of their aqueous solutions in terms of the extent of dissociation
(j) explain the terms pH, Ka, pKa, Kwand use them in calculations
(k) calculate [H+(aq)] and pH values for strong and weak acids and strong bases
(l) explain the choice of suitable indicators for acid-base titrations, given appropriate data
(m) describe the changes in pH during acid-base titrations and explain these changes in terms of the strengths of the acids and bases
(n)
(i) explain how buffer solutions control pH
(ii) describe and explain their uses, including the role of HCO3in controlling pH in blood
(o) calculate the pH of buffer solutions, given appropriate data
(p) show understanding of, and use, the concept of solubility product, Ksp
(q) calculate Kspfrom concentrations and vice versa
(r) show understanding of the common ion effect

November

Chemical Kinetics

  • Simple rate equations; orders of reaction; rate constants
  • Effect of temperature on rate constants; the concept of activation energy
  • Homogeneous and heterogeneous catalysis

Learning Outcomes

(a) explain and use the terms: rate of reaction, activation energy, catalysis, rate equation, order of reaction, rate constant, halflife of a reaction, rate-determining step
(b) explain qualitatively, in terms of collisions, the effect of concentration changes on the rate of a reaction
(c) show understanding, including reference to the Boltzmann distribution, of what is meant by the term activation energy
(d) explain qualitatively, in terms both of the Boltzmann distribution and of collision frequency, the effect of temperature change on the rate of a reaction
(e)
(i) explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower activation energy
(ii) interpret this catalytic effect in terms of the Boltzmann distribution
(f) describe enzymes as biological catalysts (proteins) which may have specific activity
(g) construct and use rate equations of the form rate = k[A]m[B]n
(limited to simple cases of single step reactions and of multistep processes with a rate-determining step, for which m and n are 0, 1 or 2), including:
(i) deducing the order of a reaction from concentration-time graphs, by the initial rates method and half-life methods
(ii) deducing, for zero- and first-order reactions, the order of reaction from concentration-time graphs
(iii) verifying that a suggested reaction mechanism is consistent with the observed kinetics
(iv) predicting the order that would result from a given reaction mechanism (and vice versa)
(v) calculating an initial rate using concentration data
(h)
(i) show understanding that the half-life of a first-order reaction is independent of concentration
(ii) use the half-life of a first-order reaction in calculations
(i) calculate a rate constant, for example by using the initial rates or half-life method
(j) devise a suitable experimental technique for studying the rate of a reaction, from given information
(k) outline the different modes of action of homogeneous and heterogeneous catalysis, including:
(i) the Haber process
(ii) the catalytic removal of oxides of nitrogen in the exhaust gases from car engines
(iii) the catalytic role of atmospheric oxides of nitrogen in the oxidation of atmospheric sulfur dioxide
(iv) catalytic role of Fe3+ in the I/S2O82- reaction

Chemical Periodicity

  • Periodicity of physical properties of the elements: variation with proton number across the third period (sodium to argon) of:
    • atomic radius and ionic radius
    • melting point
    • electrical conductivity
    • ionisation energy
  • Periodicity of chemical properties of the elements in the third period
    • Reaction of the elements with oxygen, chlorine and water
    • Variation in oxidation number of the oxides and of the chlorides
    • Reactions of these oxides and chlorides with water
    • Acid/base behaviour of these oxides and the corresponding hydroxides

Learninig Outcomes

(a) describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements
(b) explain qualitatively the variation in atomic radius and ionic radius
(c) interpret the variation in melting point and in electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements
(d) explain the variation in first ionisation energy
(e) describe the reactions, if any, of the elements with oxygen (to give Na2O, MgO, Al2O3, P4O10, SO2, SO3), chlorine (to give NaCl,MgCl2, Al2Cl6, SiCl4,PCl5) and water (Na and Mg only)
(f) state and explain the variation in oxidation number of the oxides and chlorides in terms of their valance shell electrons
(g) describe the reactions of the oxides with water
(h) describe and explain the acid/base behaviour of oxides and hydroxides including, where relevant, amphoteric behaviour in reaction with sodium hydroxide (only) and acids
(i) describe and explain the reactions of the chlorides with water
(j) interpret the variations and trends in (f), (g), (h), and (i) in terms of bonding and electronegativity
(k) suggest the types of chemical bonding present in chlorides and oxides from observations of their chemical and physical properties
In addition, candidates should be able to:
(l) predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity
(m) deduce the nature, possible position in the Periodic Table, and identity of unknown elements from given information about physical and chemical properties

December

Group 2 and 7 elements (This is only for AS levels and will be covered in the first year)

  • Similarities and trends in the properties of the Group II metals magnesium to barium and their compounds
  • Some uses of Group II compounds
  • Characteristic physical properties of Group VII
  • The relative reactivity of the elements as oxidising agents
  • Some reactions of the halide ions
  • The manufacture of chlorine
  • The reactions of chlorine with aqueous sodium hydroxide
  • The important uses of the halogens and of halogen compounds

Learning Outcomes

(a) describe the reactions of the elements with oxygen, water and dilute acids
(b) describe the behaviour of the oxides, hydroxides and carbonates with water and with dilute acids
(c) describe the thermal decomposition of the nitrates and carbonates
(d) interpret, and make predictions from, the trends in physical and chemical properties of the elements and their compounds
(e) explain the use of magnesium oxide as a refractory lining material
(f) describe and explain the use of lime in agriculture
(g) interpret and explain qualitatively the trend in the thermal stability of the nitrates and carbonates in terms of the charge density of the cation and the polarisability of the large anion
(h) interpret and explain qualitatively the variation in solubility of the sulfates in terms of relative magnitudes of the enthalpy change of hydration and the corresponding lattice energy
The following is for Group VII elements
(a) describe the colours of, and the trend in volatility of chlorine, bromine and iodine
(b) interpret the volatility of the elements in terms of van der Waals’ forces
(c) describe and deduce from E values the relative reactivity of the elements as oxidising agents
(d) describe and explain the reactions of the elements with hydrogen
(e)
(i) describe and explain the relative thermal stabilities of the hydrides
(ii) interpret these relative stabilities in terms of bond energies
(f) describe and explain the reactions of halide ions with
(i) aqueous silver ions followed by aqueous ammonia
(ii) concentrated sulfuric acid
(g) outline a method for the manufacture of chlorine from brine by a diaphragm cell
(h) describe and interpret in terms of changes of oxidation number the reaction of chlorine with cold, and with hot, aqueous sodium hydroxide
(i) explain the use of chlorine in water purification
(j) state the industrial importance and environmental significance of the halogens and their compounds (e.g. for bleaches, PVC, halogenated hydrocarbons as solvents, refrigerants and in aerosols)

Group 6 and Transition Elements(This is only for A levels and will be dealt in the second year)

  • The variation in melting points and electrical conductivities of the elements
  • The bonding, molecular shape, volatility and hydrolysis of the tetrachlorides
  • The bonding, acid/base nature and thermal stability of the oxides of oxidation states II and IV
  • The relative stability of higher and lower oxidation states for the elements in their oxides and aqueous cations
  • General physical and characteristic chemical properties of the first set of transition elements, titanium to copper
  • Colour of complexes

Learning Outcomes

(a) outline the variation in melting point and in electrical conductivity of the elements and interpret them in terms of structure and bonding
(b) describe and explain the bonding in, molecular shape and volatility of the tetrachlorides
(c) describe and explain the reactions of the tetrachlorides with water in terms of structure and bonding
(d) describe and explain the bonding, acid-base nature and thermal stability of the oxides of oxidation states IIand IV
The following is for transition elements
(a) explain what is meant by a transition element, in terms of d-block elements forming one or more stable ions with incomplete d orbitals
(b) state the electronic configuration of a first row transition element and of its ions
(c) contrast, qualitatively, the melting points and densities of the transition elements with those of calcium as a typical s-block element
(d) describe the tendency of transition elements to have variable oxidation states
(e) predict from a given electronic configuration, the likely oxidation states of a transition element
(f) describe and explain the use of Fe3+/Fe2+, MnO4/Mn2+ and Cr2O72−/Cr3+ as examples of redox systems
(g) predict, using E values, the likelihood of redox reactions
(h) explain the reactions of transition elements with ligands to form complexes, including the complexes of copper(II) ions with water, hydroxide, ammonia and chloride ions
(i)
(i) define the term ligand as a species that contains a lone pair of electrons that forms a dative bond to a central metal atom/ion.
(ii) define the term complexas a molecule or ion formed by a central metal atom/ion surrounded by one or more ligands
(iii) describe transition metal complexes as linear, octahedral, tetrahedral or square planar
(j) explain qualitatively that ligand exchange may occur, including the complexes of copper(II) ions with water, hydroxide, ammonia and chloride ions
(k) describe the shape and symmetry of the d orbitals, and the splitting of degenerate d orbitals into two energy levels in octahedral complexes using the complexes of copper(II) ions with water and ammonia as examples
(l) explain the origin of colour in transition element complexes resulting from the absorption of light energy as an electron moves between two non-degenerate d orbitals
(m) describe, in qualitative terms, the effects of different ligands on absorption, and hence colour, using the complexes of copper(II) ions with water, hydroxide, ammonia and chloride
ions as examples
(n) apply the above ideas of ligands and complexes to other metals, given information
(e) describe and explain the relative stability of higher and lower oxidation states of the elements in their oxides and aqueous cations including, where relevant, E values

January

Nitrogen and Sulphur (AS Level)

  • Nitrogen
    • Its unreactivity
    • Ammonia, the ammonium ion, nitric acid and fertilisers
    • The environmental impact of nitrogen oxides and nitrates
  • Sulfur
    • The formation of atmospheric sulfur dioxide, its role in acid rain formation, the use of sulfur dioxide in food preservation
    • Sulfuric acid

Learning Outcomes

(a) explain the lack of reactivity of nitrogen
(b) describe and explain:
(i) the basicity of ammonia
(ii) the structure of the ammonium ion and its formation by an acid-base reaction
(iii) the displacement of ammonia from its salts
(c) describe the Haber process for the manufacture of ammonia from its elements, giving essential operating conditions, and interpret these conditions (qualitatively) in terms of the principles of kinetics and equilibria
(d) state the industrial importance of ammonia and nitrogen compounds derived from ammonia
(e) state and explain the environmental consequences of the uncontrolled use of nitrate fertilisers
(f) state and explain the natural and man-made occurrences of oxides of nitrogen and their catalytic removal from the exhaust gases of internal combustion engines
(g) explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulfur dioxide
(h) describe the formation of atmospheric sulfur dioxide from the combustion of sulfur contaminated carbonaceous fuels
(i) state the role of sulfur dioxide in the formation of acid rain and describe the main environmental consequences of acid rain
(j) state the main details of the Contact process for sulfuric acid production
(k) describe the use of sulfur dioxide in food preservation

Introduction to Organic Chemistry

  • Empirical, molecular, structural, displayed and skeletal formulae
  • Functional groups and the naming of organic compounds
  • Characteristic organic reactions
  • Shapes of organic molecules; σand πbonds
  • Isomerism: structural and stereoisomerism

Learning Outccomes

(a) interpret, and use the general, structural, displayed and skeletal formulae of the following classes of compound:
(i)alkanes, alkenes and arenes
(ii)halogenoalkanes and halogenoarenes
(iii) alcohols (including primary, secondary and tertiary) and phenols
(iv) aldehydes and ketones
(v) carboxylic acids, esters and acyl chlorides
(vi) amines (primary only), nitriles, amides and amino acids
(b) interpret, and use the following terminology associated with organic reactions:
(i) functional group
(ii) homolytic and heterolytic fission
(iii) free radical, initiation, propagation, termination
(iv)nucleophile, electrophile
(v) addition, substitution, elimination, hydrolysis
(vi) oxidation and reduction
(c)
(i) describe the shapes of the ethane, ethene and benzene molecules
(ii) predict the shapes of other related molecules
(d) explain the shapes of the ethane, ethene and benzenemolecules in terms of σ and π carbon-carbon bonds
(e) describe structural isomerism, and its division into chain, positional and functional group isomerism
(f) describe stereoisomerism, and its division into geometrical (cis-trans) and optical isomerism
(g) describe cis-trans isomerism in alkenes, and explain its origin in terms of restricted rotation due to the presence of πbonds
(h) explain what is meant by a chiral centreand that such a centre gives rise to optical isomerism
(i) identify chiral centres and/or cis-trans isomerism in a molecule of given structural formula
(j) deduce the possible isomers for an organic molecule of known molecular formula
(k) deduce the molecular formula of a compound, given its structural, displayed or skeletal formula

February

Hydrocarbons

  • Alkanes (exemplified by ethane)
    • Free-radical reactions
    • Crude oil and ‘cracking’
  • Alkenes (exemplified by ethene)
    • Addition and oxidation reactions
    • Industrial importance
  • Arenes (exemplified by benzene and methylbenzene)
    • Influence of delocalised π electrons on structure and properties
    • Substitution reactions with electrophiles
    • Oxidation of side-chain
  • Hydrocarbons as fuels

Learning Outcomes

(a) show awareness of the general unreactivity of alkanes, including towards polar reagents
(b) describe the chemistry of alkanes as exemplified by the following reactions of ethane:
(i) combustion
(ii) substitution by chlorine and by bromine
(c) describe the mechanism of free-radical substitution at methyl groups with particular reference to the initiation, propagation and termination reactions
(d) describe the chemistry of alkenes as exemplified, where relevant, by the following reactions of ethene and propene (including the Markovnikov addition of asymmetric electrophiles to propene):
(i) addition of hydrogen, steam, hydrogen halides and halogens
(ii) oxidation by cold, dilute, acidified manganate(VII) ions to form the diol
(iii) oxidation by hot, concentrated, acidified manganate(VII) ions leading to the rupture of the carbon-to-carbon double bond in order to determine the position of alkene linkages in larger molecules
(iv) polymerisation
(e) describe the mechanism of electrophilic addition in alkenes, using bromine/ethene and hydrogen bromide/propene as examples
(f) explain the use of crude oil as a source of both aliphatic and aromatic hydrocarbons
(g) suggest how ‘cracking’ can be used to obtain more useful alkanes and alkenes of lower Mr from larger hydrocarbon molecules
(h) describe and explain how the combustion reactions of alkanes lead to their use as fuels in industry, in the home and in transport
(i) recognise the environmental consequences of:
(i) carbon monoxide, oxides of nitrogen and unburnt hydrocarbons arising from the internal combustion engine and of their catalytic removal
(ii) gases that contribute to the enhanced greenhouse effect
(j) describe the chemistry of arenes as exemplified by the following reactions of benzene and methylbenzene:
(i) substitution reactions with chlorine and with bromine
(ii) nitration
(iii) complete oxidation of the side-chain to give a benzoic acid
(iv) hydrogenation of the benzene ring to form a cyclohexane ring
(k)
(i) describe the mechanism of electrophilic substitution in arenes, as exemplified by the formation of nitrobenzene and bromobenzene
(ii) suggest the mechanism of other electrophilic substitution reactions, given data
(iii) describe the effect of the delocalisation of electrons in arenes in such reactions
(l) predict whether halogenation will occur in the side-chain or aromatic nucleus in arenes depending on reaction conditions
(m) apply the knowledge of positions of substitution in the electrophilic substitution of arenes

Halogen Derivates

  • Halogenoalkanes and halogenoarenes
    • Nucleophilic substitution
    • Hydrolysis
    • Formation of nitriles, primary amines
    • Elimination
  • Relative strength of the C-Hal bond

Learning Outcomes

(a) recall the chemistry of halogenoalkanes as exemplified by
(i) the following nucleophilic substitution reactions of bromoethane: hydrolysis, formation of nitriles, formation of primary amines by reaction with ammonia
(ii) the elimination of hydrogen bromide from 2-bromopropane
(b) describe the mechanism of nucleophilic substitution (by both SN1 and SN2 mechanisms) in halogenoalkanes
(c) interpret the different reactivities of halogenoalkanes and chlorobenzene(with particular reference to hydrolysis and to the relative strengths of the C-Hal bonds)
(d) explain the uses of fluoroalkanes and fluorohalogenoalkanes in terms of their relative chemical inertness
(e) recognise the concern about the effect of chlorofluoroalkanes on the ozone layer

Hydroxy Compounds

  • Alcohols (exemplified by ethanol)
    • Formation of halogenoalkanes
    • Reaction with sodium; oxidation; dehydration; esterification; acylation
    • The tri-iodomethane test
  • Phenol
    • Its acidity; reaction with sodium
    • Nitration of, and bromination of, the aromatic ring

Learning Outcomes

a) recall the chemistry of alcohols, exemplified by ethanol:
(i) combustion
(ii) substitution to give halogenoalkanes
(iii) reaction with sodium
(iv) oxidation to carbonyl compounds and carboxylic acids
(v) dehydration to alkenes
(vi) formation of esters by esterification with carboxylic acids and acylation with acyl chlorides
(b)
(i) classify hydroxy compounds into primary, secondary and tertiary alcohols
(ii) suggest characteristic distinguishing reactions, e.g. mild oxidation
(c) deduce the presence of a CH3CH(OH)– group in an alcohol from its reaction with alkaline aqueous iodine to form tri-iodomethane
(d) recall the chemistry of phenol, as exemplified by the following reactions:
(i) with bases
(ii) with sodium
(iii)with diazonium salts
(iv) nitration of, and bromination of, the aromatic ring
(e) describe and explain the relative acidities of water, phenol and ethanol

March

Carbonyl compounds

  • Aldehydes (exemplified by ethanal)
    • Oxidation to carboxylic acids
    • Reaction with hydrogen cyanide
    • Characteristic tests for aldehydes
  • Ketones (exemplified by propanone and phenylethanone)
    • Reaction with hydrogen cyanide
    • Characteristic tests for ketones

Learning Outcomes

(a) describe
(i) the formation of aldehydes and ketones from primary and secondary alcohols respectively using Cr2O72–/H+
(ii) the reduction of aldehydes and ketones e.g. using NaBH4 or LiAlH4
(iii) the reaction of aldehydes and ketones with HCN and NaCN
(b) describe the mechanism of the nucleophilic addition reactions of hydrogen cyanide with aldehydes and ketones
(c) describe the use of 2,4-dinitrophenylhydrazine (2,4-DNPH) reagent to detect the presence of carbonyl compounds
(d) deduce the nature (aldehyde or ketone) of an unknown carbonyl compound from the results of simple tests (i.e. Fehling’s and Tollens’ reagents; ease of oxidation)
(e) describe the reaction of CH3CO– compounds with alkaline aqueous iodine to give tri-iodomethane

Carboxylic acids and Derivates

  • Carboxylic acids (exemplified by ethanoic acid and benzoic acid)
    • Formation from primary alcohols and nitriles
    • Salt, ester and acyl chloride formation
  • Acyl chlorides (exemplified by ethanoyl chloride)
    • Ease of hydrolysis compared with alkyl and aryl chlorides
    • Reaction with alcohols, phenols and primary amines
  • Esters (exemplified by ethyl ethanoate and phenyl benzoate)
    • Formation from carboxylic acids and from acyl chlorides
    • Hydrolysis (under acidic and under basic conditions)
    • Uses of esters

Learning Outcomes

(a)describe the formation of carboxylic acids from alcohols, aldehydes and nitriles
(b) describe the reactions of carboxylic acids in the formation of
(i) salts, by the use of reactive metals, alkalis or carbonates
(ii)esters
(iii)acyl chlorides
(c) explain the acidity of carboxylic acids and of chlorinesubstituted ethanoic acids in terms of their structures
(d) describe the hydrolysis of acyl chlorides
(e) describe the reactions of acyl chlorides with alcohols, phenols and primary amines
(f) explain the relative ease of hydrolysis of acyl chlorides, alkyl chlorides and aryl chlorides
(g) describe the formation of esters from carboxylic acids or acyl chlorides, using ethyl ethanoate and phenyl benzoateas examples
(h) describe the acid and base hydrolysis of esters
(i) describe the formation of polyesters
(j) state the major commercial uses of esters e.g. solvents, perfumes, flavourings

Nitrogen Compunds

  • Primary amines (exemplified by ethylamine and phenylamine)
    • Formation
    • Salt formation
    • Other reactions of phenylamine
  • Amides (exemplified by ethanamide)
  • Formation from acyl chlorides
  • Hydrolysis
  • Reduction
  • Amino acids (exemplified by aminoethanoic acid)
    • Acid and base properties
    • Zwitterion formation
  • Proteins
    • Structure, based on the peptide linkage
    • Hydrolysis of proteins

Learning Outcomes

(a) describe the formation of alkyl amines such as ethylamine (by the reaction of ammonia with halogenoalkanes; the reduction of amides with LiAlH4; the reduction of nitriles with LiAlH4 or H2/Ni) and of phenylamine (by the reduction of nitrobenzene with tin/concentrated HCl)
(b) describe and explain the basicity of amines
(c) explain the relative basicities of ammonia, ethylamine and phenylamine in terms of their structures
(d) describe the reaction of phenylamine with:
(i) aqueous bromine
(ii) nitrous acid to give the diazonium salt and phenol
(e) describe the coupling of benzenediazonium chloride and phenol and the use of similar reactions in the formation of dyestuff
(f) describe the formation of amides from the reaction between RNH2 and R’COCl
(g) recognise that amides are neutral
(h)
(i) describe amide hydrolysis on treatment with aqueous alkali or acid
(ii) describe the reduction of amides with LiAlH4
(i) describe the acid/base properties of amino acids and the formation of zwitterions
(j) describe the formation of peptide bonds between amino acids and, hence, explain protein formation
(k) describe the hydrolysis of proteins
(l) describe the formation of polyamides

Polymerization

  • Addition polymerisation
  • Condensation polymerisation

Learning Outcomes

(a) describe the characteristics of addition polymerisation as exemplified by poly(ethene) and PVC
(b) recognise the difficulty of the disposal of poly(alkene)s, i.e. nonbiodegradability and harmful combustion products
(c) describe the characteristics of condensation polymerisation
(i) in polyesters as exemplified by Terylene
(ii) in polyamides as exemplified by peptides, proteins, nylon 6 and nylon 6,6
(d) predict the type of polymerisation reaction for a given monomer or pair of monomers
(e) deduce the repeat unit of a polymer obtained from a given monomer or pair of monomers
(f) deduce the type of polymerisation reaction which produces a given section of a polymer molecule
(g) identify the monomer(s) present in a given section of a polymer molecule

April

THe Chemistry of Life

  • Protein chemistry
  • Genetic information
  • Energy
  • Metals in biological systems

Learning Outcomes

(a) recall that proteins are condensation polymers formed from amino acid monomers and recognise and describe the generalised structure of amino acids
(b) explain the importance of amino acid sequence (primary structure) in determining the properties of proteins
(c) distinguish between the primary, secondary (α-helix and β-sheet) and tertiary structures of proteins and explain the stabilisation of secondary (through hydrogen bonding between C=O and N-H bonds of peptide groups) and tertiary (through interactions between R-groups) structure
(d) describe and explain the characteristics of enzyme catalysis, including
(i) specificity (using a simple lock and key model) and the idea of competitive inhibition
(ii) structural integrity in relation to denaturation and noncompetitive inhibition
(e) given information, use core chemistry to explain how small molecules interact with proteins and how they can modify the structure and function of biological systems (for example, as enzyme inhibitors or cofactors, disrupting protein-protein
interactions, blocking ion channels)
(f) describe the double helical structure of DNA in terms of a sugar-phosphate backbone and attached bases
(g) explain the significance of hydrogen-bonding in the pairing of bases in DNA in relation to the replication of genetic information
(h) explain in outline how DNA encodes for the amino acid sequence of proteins with reference to mRNA, tRNA and the ribosome in translation and transcription
(i) explain the chemistry of DNA mutation from provided data
(j) discuss the genetic basis of disease (for example, sickle cell anaemia) in terms of altered base sequence, causing alterations in protein structure and function
(k) explain how modification to protein/enzyme primary structure can result in new structure and/or function
(l) outline, in terms of the hydrolysis of ATP to ADP + Pi , the provision of energy for the cell
(m) understand why some metals are essential to life and, be able to explain the chemistry involved (for
example, iron in haemoglobin, sodium and potassium in transmission of nerve impulses , zinc as an enzyme cofactor
(n) recognise that some metals are toxic and discuss, in chemical terms, the problems associated with heavy metals in the environment entering the food chain, for example mercury

Analytical Chemistry

  • Methods of detection and analysis
  • Applications in chemistry and society

Learning Outcomes

(a) describe simply the process of electrophoresis and the effect of pH, using peptides and amino acids as examples
(b) explain, in simple terms, the technique of DNA fingerprinting and its applications in forensic science, archaeology and medicine
(c) describe the importance to modern medicine, and the challenges, of separating and characterising the proteins in cells
(d) outline in simple terms the principles of nuclear magnetic resonance in 1H and be able to interpret simple NMR spectra, using chemical shift values, splitting patterns and the effect of adding D2O to a sample
(e) show awareness of the use of NMR and X-ray crystallography in determining the structure of macromolecules and in understanding their function
(f) state what is meant by partition coefficient and calculate a partition coefficient for a system in which the solute is in the same molecular state in the two solvents
(g) understand qualitatively paper, high performance liquid, thin layer and gas/liquid chromatography in terms of adsorption and/or partition and be able to interpret data from these techniques
(h) explain the concept of mass spectroscopy, deduce the number of carbon atoms in a compound using the M+1 peak and the presence of bromine and chlorine atoms using the
M+2 peak and suggest the identity of molecules formed by simple fragmentation in a given mass spectrum
(i) draw conclusions given appropriate information and data from environmental monitoring (for example, PCBs in the atmosphere, isotopic ratios in ice cores)

Design and Materials

  • Medicinal chemistry and drug delivery
  • Properties of polymers
  • Nanotechnology
  • Environment and energy

Learning Outcomes

(a) discuss the challenges of drug design and explain in simple terms how molecules may be identified and developed to overcome these problems
(b) discuss the challenges of drug delivery and explain in simple terms how materials may be developed to overcome these problems
(c) discuss the properties and structure of polymers based on their methods of formation
(d) discuss how the presence of side-chains and intermolecular forces affect the properties of polymeric materials (for example, spider silk)
(e) show awareness of nanotechnology and, given information and data, be able to discuss the chemistry involved
(f) discuss how a knowledge of chemistry can be used to overcome environmental problems (for example, ground water contamination, oil spillage, CFCs)
(g) discuss how a knowledge of chemistry can be used to extend the life of existing resources, to identify alternative resources and to improve the efficiency of energy production and use

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