AQA Oxford A LEVEL Chemistry Syllabus

	Notes	Questions	Answers	Part	AS	A2
ATOMIC STRUCTURE	Open Here	Open Here	Open Here	PHYSICAL CHEMISTRY	✔
AMOUNT OF SUBSTANCE				PHYSICAL CHEMISTRY	✔
BONDING				PHYSICAL CHEMISTRY	✔
ENERGETICS				PHYSICAL CHEMISTRY	✔
KINETICS				PHYSICAL CHEMISTRY	✔
CHEMICAL EQUILIBRIA				PHYSICAL CHEMISTRY	✔
REDOX EQUATIONS				PHYSICAL CHEMISTRY	✔
THERMODYNAMICS				PHYSICAL CHEMISTRY		✔
RATE EQUATIONS				PHYSICAL CHEMISTRY		✔
EQUILIBRIUM CONSTANT Kp				PHYSICAL CHEMISTRY		✔
ELECTRODE POTENTIALS AND ELECTROCHEMICAL CELLS				PHYSICAL CHEMISTRY		✔
ACIDS AND BASES	Download Notes	Download Questions	Download answers	PHYSICAL CHEMISTRY		✔

3.1 PHYSICAL CHEMISTRY
3.1.1 ATOMIC STRUCTURE (INTERNATIONAL AS)

The chemical properties of elements depend on their atomic structure and in particular on the arrangement of electrons around the nucleus. The arrangement of electrons in orbitals is linked to the way elements are organised in the Periodic Table. Chemists can measure the mass of atoms and molecules to a high degree of accuracy in a mass spectrometer. The principles of operation of a modern mass spectrometer are studied.
3.1.1.1 Fundamental particles
Appreciate that knowledge and understanding of atomic structure has evolved over time. Protons, neutrons and electrons: relative charge and relative mass.
An atom consists of a nucleus containing protons and neutrons surrounded by electrons.3.1.1.2 Mass number and isotopes
Mass number (A) and atomic (proton) number (Z).
Students should be able to:

• determine the number of fundamental particles in atoms and ions using mass number, atomic number and charge
• explain the existence of isotopes.
  The principles of a simple time of flight (TOF) mass spectrometer, limited to ionisation, acceleration to give all
  ions constant kinetic energy, ion drift, ion detection, data analysis.
  The mass spectrometer gives accurate information about relative isotopic mass and also about the relative abundance of isotopes.
  Mass spectrometry can be used to identify elements.
  Mass spectrometry can be used to determine relative molecular mass. Students should be able to:

• interpret simple mass spectra of elements
• calculate relative atomic mass from isotopic abundance, limited to mononuclear ions.

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3.1.1.3 Electron configuration
Electron configurations of atoms and ions up to Z = 36 in terms of shells and sub-shells (orbitals) s, p and d. Ionisation energies.
Students should be able to:

• define first ionisation energy
• write equations for first and successive ionisation energies
• explain how first and successive ionisation energies in Period 3 (Na–Ar) and in Group 2 (Be–Ba) give
  evidence for electron configuration in sub-shells and in shells.
  3.1.2 AMOUNT OF SUBSTANCE (INTERNATIONAL AS)
  When chemists measure an amount of a substance, they use an amount in moles. The mole is a useful quantity because one mole of a substance always contains the same number of entities of the substance. An amount in moles can be measured out by mass in grams, by volume in dm3 of a solution of known concentration and by volume in dm3 of a gas.

3.1.2.1 Relative atomic mass and relative molecular mass
Relative atomic mass and relative molecular mass in terms of 12C. The term relative formula mass will be used for ionic compounds. Students should be able to:

• define relative atomic mass (Ar)
• define relative molecular mass (Mr).

3.1.2.2 The mole and the Avogadro constant
The Avogadro constant as the number of particles in a mole.
The mole as applied to electrons, atoms, molecules, ions, formulas and equations.
The concentration of a substance in solution, measured in mol dm–3.
Students should be able to carry out calculations:

• using the Avogadro constant
• using mass of substance, Mr , and amount in moles
• using concentration, volume and amount of substance in a solution.
  Students will not be expected to recall the value of the Avogadro constant.

3.1.2.3 The ideal gas equation
The ideal gas equation pV = nRT with the variables in SI units.
Students should be able to:• use the equation in calculations.
Students will not be expected to recall the value of the gas constant, R.

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3.1.2.4 Empirical and molecular formula
Empirical formula is the simplest whole number ratio of atoms of each element in a compound. Molecular formula is the actual number of atoms of each element in a compound.
The relationship between empirical formula and molecular formula.
Students should be able to:

• calculate empirical formula from data giving composition by mass or percentage by mass
• calculate molecular formula from the empirical formula and relative molecular mass.

3.1.2.5 Balanced equations and associated calculations
Equations (full and ionic). Percentage atom economy is:
molecular mass of desired product × 100sum of molecular masses of all reactants
Economic, ethical and environmental advantages for society and for industry of developing chemical processes with a high atom economy.
Students should be able to:

• write balanced equations for reactions studied
• balance equations for unfamiliar reactions when reactants and products are specified.
  Students should be able to use balanced equations to calculate:

• masses
• volumes of gases
• percentage yields
• percentage atom economies
• concentrations and volumes for reactions in solutions.
  Required practical 1:
  Make up a volumetric solution and carry out a simple acid–base titration.

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3.1.3 BONDING (INTERNATIONAL AS)
The physical and chemical properties of compounds depend on the ways in which the compounds are held together by chemical bonds and intermolecular forces. Theories of bonding explain how atoms or ions are held together in these structures. Materials scientists use knowledge of structure and bonding to engineer new materials with desirable properties. These new materials may offer new applications in a range of different modern technologies.
3.1.3.1 Ionic bonding
Ionic bonding involves electrostatic attraction between oppositely charged ions in a lattice. The formulas of compound ions, eg sulfate, hydroxide, nitrate, carbonate and ammonium. Students should be able to:

• predict the charge on a simple ion using the position of the element in the Periodic Table
• construct formulas for ionic compounds.

3.1.3.2 Nature of covalent and dative covalent bonds
A single covalent bond contains a shared pair of electrons. Multiple bonds contain multiple pairs of electrons.
A co-ordinate (dative covalent) bond contains a shared pair of electrons with both electrons supplied by one atom.
Students should be able to represent:

• a covalent bond using a line
• a co-ordinate bond using an arrow.

3.1.3.3 Metallic bonding
Metallic bonding involves attraction between delocalised electrons and positive ions arranged in a lattice.
3.1.3.4 Bonding and physical properties
The four types of crystal structure:

• ionic
• metallic
• macromolecular (giant covalent)
• molecular.
  The structures of the following crystals as examples of these four types of crystal structure:

• diamond
• graphite
• ice
• iodine
• magnesium
• sodium chloride.

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Students should be able to:

• relate the melting point and conductivity of materials to the type of structure and the bonding present
• explain the energy changes associated with changes of state
• draw diagrams to represent these structures involving specified numbers of particles.

3.1.3.5 Shapes of simple molecules and ions
Bonding pairs and lone (non-bonding) pairs of electrons as charge clouds that repel each other.
Pairs of electrons in the outer shell of atoms arrange themselves as far apart as possible to minimise repulsion.
Lone pair–lone pair repulsion is greater than lone pair–bond pair repulsion, which is greater than bond pair– bond pair repulsion.
The effect of electron pair repulsion on bond angles.
Students should be able to:
• name and explain the shapes of, and bond angles in, simple molecules and ions with up to six electron pairs (including lone pairs of electrons) surrounding the central atom.
3.1.3.6 Bond polarity
Electronegativity as the power of an atom to attract the pair of electrons in a covalent bond.
The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical. This produces a polar covalent bond, and may cause a molecule to have a permanent dipole.
Students should be able to:

• use partial charges to show that a bond is polar
• explain why some molecules with polar bonds do not have a permanent dipole.

3.1.3.7 Forces between molecules
Forces between molecules:

• permanent dipole–dipole forces
• induced dipole–dipole (van der Waals, dispersion, London) forces
• hydrogen bonding.
  The melting and boiling points of molecular substances are influenced by the strength of these intermolecular forces.
  The importance of hydrogen bonding in the low density of ice and the anomalous boiling points of compounds.
  Students should be able to:

• explain the existence of these forces between familiar and unfamiliar molecules
• explain how melting and boiling points are influenced by these intermolecular forces.

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3.1.4 ENERGETICS (INTERNATIONAL AS)
The enthalpy change in a chemical reaction can be measured accurately. It is important to know this value for chemical reactions that are used as a source of heat energy in applications such as domestic boilers and internal combustion engines.
3.1.4.1 Enthalpy change
Reactions can be endothermic or exothermic.
Enthalpy change (∆H) is the heat energy change measured under conditions of constant pressure.
Standard enthalpy changes refer to standard conditions, ie 100 kPa and a stated temperature (eg ∆H298q).
Students should be able to:

• define standard enthalpy of combustion (∆cHq)
• define standard enthalpy of formation (∆fHq).
  3.1.4.2 Calorimetry
  The heat change, q, in a reaction is given by the equation
  q = mc∆T
  where m is the mass of the substance that has a temperature change ∆T and a specific heat capacity c.
  Students should be able to:

• use this equation to calculate the molar enthalpy change for a reaction
• use this equation in related calculations.
  Students will not be expected to recall the value of the specific heat capacity, c, of a substance.
  Required practical 2:
  Measure an enthalpy change.

3.1.4.3 Applications of Hess’s law
Hess’s law.
Students should be able to:
• use Hess’s law to perform calculations, including calculation of enthalpy changes for reactions from enthalpies of combustion or from enthalpies of formation.
3.1.4.4 Bond enthalpies
Mean bond enthalpy. Students should be able to:

• define the term mean bond enthalpy
• use mean bond enthalpies to calculate an approximate value of ∆H for reactions in the gaseous phase
• explain why values from mean bond enthalpy calculations differ from those determined using Hess’s law.

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3.1.5 KINETICS (INTERNATIONAL AS)
The study of kinetics enables chemists to determine how a change in conditions affects the speed of a chemical reaction. Whilst the reactivity of chemicals is a significant factor in how fast chemical reactions proceed, there are variables that can be manipulated in order to speed them up or slow them down.
3.1.5.1 Collision theory
Reactions can only occur when collisions take place between particles having sufficient energy. This energy is called the activation energy.
Students should be able to:

• define the term activation energy
• explain why most collisions do not lead to a reaction.

3.1.5.2 Maxwell–Boltzmann distribution
Maxwell–Boltzmann distribution of molecular energies in gases.
Students should be able to:
• draw and interpret distribution curves for different temperatures.
3.1.5.3 Effect of temperature on reaction rate
Meaning of the term rate of reaction.
The qualitative effect of temperature changes on the rate of reaction.
Students should be able to:
• use the Maxwell–Boltzmann distribution to explain why a small temperature increase can lead to a large increase in rate.
3.1.5.4 Effect of concentration and pressure
The qualitative effect of changes in concentration on collision frequency.
The qualitative effect of a change in the pressure of a gas on collision frequency.
Students should be able to:
• explain how a change in concentration or a change in pressure influences the rate of a reaction.
3.1.5.5 Catalysts
A catalyst is a substance that increases the rate of a chemical reaction without being changed in chemical composition or amount.
Catalysts work by providing an alternative reaction route of lower activation energy.
Students should be able to:
• use a Maxwell–Boltzmann distribution to help explain how a catalyst increases the rate of a reaction involving a gas.

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3.1.6 CHEMICAL EQUILIBRIA, LE CHATELIER’S PRINCIPLE AND Kc(INTERNATIONAL AS)
In contrast with kinetics, which is a study of how quickly reactions occur, a study of equilibria indicates how far reactions will go. Le Chatelier’s principle can be used to predict the effects of changes in temperature, pressure and concentration on the yield of a reversible reaction. This has important consequences for many industrial processes. The further study of the equilibrium constant, Kc, considers how the mathematical expression for the equilibrium constant enables us to calculate how an equilibrium yield will be influenced by the concentration of reactants and products.
3.1.6.1 Chemical equilibria and Le Chatelier’s principle
Many chemical reactions are reversible.
In a reversible reaction at equilibrium:

• forward and reverse reactions proceed at equal rates
• the concentrations of reactants and products remain constant.
  Le Chatelier’s principle.
  Le Chatelier’s principle can be used to predict the effects of changes in temperature, pressure and concentration on the position of equilibrium in homogeneous reactions.
  A catalyst does not affect the position of equilibrium.
  Students should be able to:

• use Le Chatelier’s principle to predict qualitatively the effect of changes in temperature, pressure and concentration on the position of equilibrium
• explain why, for a reversible reaction used in an industrial process, a compromise temperature and pressure may be used.

3.1.6.2 Equilibrium constant Kc for homogeneous systemsThe equilibrium constant Kc is deduced from the equation for a reversible reaction.
The concentration, in mol dm–3, of a species X involved in the expression for Kc is represented by [X] The value of the equilibrium constant is not affected either by changes in concentration or addition of a
catalyst.
Students should be able to:

• construct an expression for Kc for a homogeneous system in equilibrium
• calculate a value for Kc from the equilibrium concentrations for a homogeneous system at constant temperature
• perform calculations involving Kc
• predict the qualitative effects of changes of temperature on the value of Kc.

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3.1.7 OXIDATION, REDUCTION AND REDOX EQUATIONS (INTERNATIONAL AS)
Redox reactions involve a transfer of electrons from the reducing agent to the oxidising agent. The change in the oxidation state of an element in a compound or ion is used to identify the element that has been oxidised or reduced in a given reaction. Separate half-equations are written for the oxidation or reduction processes. These half-equations can then be combined to give an overall equation for any redox reaction.
3.1.7.1 Oxidation, reduction and redox equations
Oxidation is the process of electron loss and oxidising agents are electron acceptors. Reduction is the process of electron gain and reducing agents are electron donors. The rules for assigning oxidation states.
Students should be able to:

• work out the oxidation state of an element in a compound or ion from the formula
• write half-equations identifying the oxidation and reduction processes in redox reactions
• combine half-equations to give an overall redox equation.
  3.1.8 THERMODYNAMICS (INTERNATIONAL A2)
  The further study of thermodynamics builds on the Energetics section and is important in understanding the stability of compounds and why chemical reactions occur. Enthalpy change is linked with entropy change enabling the free-energy change to be calculated.

3.1.8.1 Born–Haber cycles
Lattice enthalpy can be defined as either enthalpy of lattice dissociation or enthalpy of lattice formation. Born–Haber cycles are used to calculate lattice enthalpies using the following data:

• enthalpy of formation
• ionisation energy
• enthalpy of atomisation
• bond enthalpy
• electron affinity. Students should be able to:

• define each of the above terms and lattice enthalpy
• construct Born–Haber cycles to calculate lattice enthalpies using these enthalpy changes
• construct Born–Haber cycles to calculate one of the other enthalpy changes
• compare lattice enthalpies from Born–Haber cycles with those from calculations based on a perfect ionic model to provide evidence for covalent character in ionic compounds.
  Cycles can be used to calculate enthalpies of solution for ionic compounds from lattice enthalpies and enthalpies of hydration.

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Students should be able to:

• define the term enthalpy of hydration
• perform calculations of an enthalpy change using these cycles.

3.1.8.2 Gibbs free-energy change, ∆G, and entropy change, ∆S∆H, whilst important, is not sufficient to explain feasible change.
The concept of increasing disorder (entropy change, ∆S).∆S accounts for the above deficiency, illustrated by physical changes and chemical changes.
The balance between entropy and enthalpy determines the feasibility of a reaction given by the relationship:∆G = ∆H – T∆S (derivation not required).
For a reaction to be feasible, the value of ∆G must be zero or negative.
Students should be able to:

• calculate entropy changes from absolute entropy values
• use the relationship ∆G = ∆H – T∆S to determine how ∆G varies with temperature
• use the relationship ∆G = ∆H – T∆S to determine the temperature at which a reaction becomes feasible.
  3.1.9 RATE EQUATIONS (INTERNATIONAL A2)
  In rate equations, the mathematical relationship between rate of reaction and concentration gives information about the mechanism of a reaction that may occur in several steps.

3.1.9.1 Rate equations
The rate of a chemical reaction is related to the concentration of reactants by a rate equation of the form:Rate = k[A]m [B]nwhere m and n are the orders of reaction with respect to reactants A and B and k is the rate constant. The orders m and n are restricted to the values 0, 1, and 2.
The rate constant k varies with temperature as shown by the equation:k = Ae–Ea/RT
where A is a constant, known as the Arrhenius constant, Ea is the activation energy and T is the temperature in K.
Students should be able to:

• define the terms order of reaction and rate constant
• perform calculations using the rate equation
• explain the qualitative effect of changes in temperature on the rate constant k
• perform calculations using the equation k = Ae–Ea/RT
• understand that the equation k = Ae–Ea/RT can be rearranged into the form ln k = –Ea/RT + ln A and know how to use this rearranged equation with experimental data to plot a straight line graph with slope –Ea/R.
  These equations and the gas constant, R, will be given when required.

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3.1.9.2 Determination of rate equation
The rate equation is an experimentally determined relationship.
The orders with respect to reactants can provide information about the mechanism of a reaction.
Students should be able to:

• use concentration–time graphs to deduce the rate of a reaction
• use initial concentration–time data to deduce the initial rate of a reaction
• use rate–concentration data or graphs to deduce the order (0, 1 or 2) with respect to a reactant
• derive the rate equation for a reaction from the orders with respect to each of the reactants
• use the orders with respect to reactants to provide information about the rate determining/limiting step of a reaction.
  Required practical 6:
  Measure the rate of reaction by an initial rate method and a continuous monitoring method.
  3.1.10 EQUILIBRIUM CONSTANT Kp FOR HOMOGENEOUS SYSTEMS (INTERNATIONAL A2)
  The further study of equilibria considers how the mathematical expression for the equilibrium constant Kpenables us to calculate how an equilibrium yield will be influenced by the partial pressures of reactants and products. This has important consequences for many industrial processes.
  3.1.10.1 Equilibrium constant Kp for homogeneous systemsThe equilibrium constant Kp is deduced from the equation for a reversible reaction occurring in the gas phase.
  Kp is the equilibrium constant calculated from partial pressures for a system at constant temperature. Students should be able to:

• derive partial pressure from mole fraction and total pressure
• construct an expression for Kp for a homogeneous system in equilibrium
• perform calculations involving Kp
• predict the qualitative effects of changes in temperature and pressure on the position of equilibrium
• predict the qualitative effects of changes in temperature on the value of Kp
• understand that, whilst a catalyst can affect the rate of attainment of an equilibrium, it does not affect the value of the equilibrium constant.

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3.1.11 ELECTRODE POTENTIALS AND ELECTROCHEMICAL CELLS (INTERNATIONAL A2)
Redox reactions take place in electrochemical cells where electrons are transferred from the reducing agent to the oxidising agent indirectly via an external circuit. A potential difference is created that can drive an electric current to do work. Electrochemical cells have very important commercial applications as a portable supply of electricity to power electronic devices such as mobile phones, tablets and laptops. On a larger scale, they can provide energy to power a vehicle.
3.1.11.1 Electrode potentials and cells
IUPAC convention for writing half-equations for electrode reactions.
The conventional representation of cells.
Cells are used to measure electrode potentials by reference to the standard hydrogen electrode.
The importance of the conditions when measuring the electrode potential, E (Nernst equation not required).
Standard electrode potential, Eq, refers to conditions of 298 K, 100 kPa and 1.00 mol dm−3 solution of ions.
Standard electrode potentials can be listed as an electrochemical series.
Students should be able to:

• use Eq values to predict the direction of simple redox reactions
• calculate the EMF of a cell
• write and apply the conventional representation of a cell.
  Required practical 7:
  Measure the EMF of an electrochemical cell.
  3.1.11.2 Commercial applications of electrochemical cells
  Electrochemical cells can be used as a commercial source of electrical energy. The simplified electrode reactions in a lithium cell:
  Positive electrode: Li+ + CoO2 + e– g Li+[CoO2]–Negative electrode: LigLi+ + e–
  Cells can be non-rechargeable (irreversible), rechargeable or fuel cells.
  Fuel cells are used to generate an electric current and do not need to be electrically recharged. The electrode reactions in an alkaline hydrogen–oxygen fuel cell.
  The benefits and risks to society associated with using these cells.
  Students should be able to:

• use given electrode data to deduce the reactions occurring in non-rechargeable and rechargeable cells
• deduce the EMF of a cell
• explain how the electrode reactions can be used to generate an electric current.

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3.1.12 ACIDS AND BASES (INTERNATIONAL A2)
Acids and bases are important in domestic, environmental and industrial contexts. Acidity in aqueous solutions is caused by hydrogen ions and a logarithmic scale, pH, has been devised to measure acidity. Buffer solutions, which can be made from partially neutralised weak acids, resist changes in pH and find many important industrial and biological applications.
3.1.12.1 Brønsted–Lowry acid–base equilibria in aqueous solution
An acid is a proton donor.
A base is a proton acceptor.
Acid–base equilibria involve the transfer of protons.
3.1.12.2 Definition and determination of pH
The concentration of hydrogen ions in aqueous solution covers a very wide range. Therefore, a logarithmic scale, the pH scale, is used as a measure of hydrogen ion concentration.
pH = –log10[H+]
Students should be able to:

• convert concentration of hydrogen ions into pH and vice versa
• calculate the pH of a solution of a strong acid from its concentration.
  3.1.12.3 The ionic product of water, KwWater is slightly dissociated.
  Kw is derived from the equilibrium constant for this dissociation.Kw = [H+][OH–]
  The value of Kw varies with temperature.
  Students should be able to:

• use Kw to calculate the pH of a strong base from its concentration.
3.1.12.4 Weak acids and bases; Ka for weak acidsWeak acids and weak bases dissociate only slightly in aqueous solution.
Ka is the dissociation constant for a weak acid. pKa = –log10 KaStudents should be able to:

• construct an expression for Ka
• perform calculations relating the pH of a weak acid to the concentration of the acid and the dissociation constant, Ka
• convert Ka into pKa and vice versa.

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3.1.12.5 pH curves, titrations and indicators
Titrations of acids with bases.
Students should be able to:• perform calculations for these titrations based on experimental results.
Typical pH curves for acid–base titrations in all combinations of weak and strong monoprotic acids and bases. Students should be able to:

• sketch and explain the shapes of typical pH curves
• use pH curves to select an appropriate indicator.
  Required practical 8:
  Investigate how pH changes when a weak acid reacts with a strong base and when a strong acid reacts with a weak base.
  3.1.12.6 Buffer action
  A buffer solution maintains an approximately constant pH, despite dilution or addition of small amounts of acid or base.
  Acidic buffer solutions contain a weak acid and the salt of that weak acid. Basic buffer solutions contain a weak base and the salt of that weak base. Applications of buffer solutions.
  Students should be able to:

• explain qualitatively the action of acidic and basic buffers
• calculate the pH of acidic buffer solutions.

3.2 INORGANIC CHEMISTRY
3.2.1 PERIODICITY (INTERNATIONAL AS)
The Periodic Table provides chemists with a structured organisation of the known chemical elements from which they can make sense of their physical and chemical properties. The historical development of the Periodic Table and models of atomic structure provide good examples of how scientific ideas and explanations develop over time.
3.2.1.1 Classification
An element is classified as s, p, d or f block according to its position in the Periodic Table, which is determined by its proton number.
3.2.1.2 Physical properties of Period 3 elements
The trends in atomic radius, first ionisation energy and melting point of the elements Na–Ar The reasons for these trends in terms of the structure of and bonding in the elements. Students should be able to:

• •

explain the trends in atomic radius and first ionisation energy
explain the melting point of the elements in terms of their structure and bonding.

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3.2.2 GROUP 2, THE ALKALINE EARTH METALS (INTERNATIONAL AS)
The elements in Group 2 are called the alkaline earth metals. The trends in the solubilities of the hydroxides and the sulfates of these elements are linked to their use. Barium sulfate, magnesium hydroxide and magnesium sulfate have applications in medicines whilst calcium hydroxide is used in agriculture to change soil pH, which is essential for good crop production and maintaining the food supply.
3.2.2.1 Group 2, the alkaline earth metals
The trends in atomic radius, first ionisation energy and melting point of the elements Mg–Ba Students should be able to:

• explain the trends in atomic radius and first ionisation energy
• explain the melting point of the elements in terms of their structure and bonding.
  The reactions of the elements Mg–Ba with water.
  The use of magnesium in the extraction of titanium from TiCl4The relative solubilities of the hydroxides of the elements Mg–Ba in water. Mg(OH)2 is sparingly soluble.
  The use of Mg(OH)2 in medicine and of Ca(OH)2 in agriculture.
  The use of CaO or CaCO3 to remove SO2 from flue gases.
  The relative solubilities of the sulfates of the elements Mg–Ba in water. BaSO4 is insoluble.
  The use of acidified BaCl2 solution to test for sulfate ions.
  The use of BaSO4 in medicine.
  Students should be able to:

• explain why BaCl2 solution is used to test for sulfate ions and why it is acidified.
3.2.3 GROUP 7(17), THE HALOGENS (INTERNATIONAL AS)
The halogens in Group 7 are very reactive non-metals. Trends in their physical properties are examined and explained. Fluorine is too dangerous to be used in a school laboratory but the reactions of chlorine are studied. Challenges in studying the properties of elements in this group include explaining the trends in ability of the halogens to behave as oxidising agents and the halide ions to behave as reducing agents.
3.2.3.1 Trends in properties
The trends in electronegativity and boiling point of the halogens.
Students should be able to:

• explain the trend in electronegativity
• explain the trend in the boiling point of the elements in terms of their structure and bonding.
  The trend in oxidising ability of the halogens down the group, including displacement reactions of halide ions in aqueous solution.

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The trend in reducing ability of the halide ions, including the reactions of solid sodium halides with concentrated sulfuric acid.
The use of acidified silver nitrate solution to identify and distinguish between halide ions. The trend in solubility of the silver halides in ammonia.
Students should be able to explain why:

• silver nitrate solution is used to identify halide ions
• the silver nitrate solution is acidified
• ammonia solution is added.

3.2.3.2 Uses of chlorine and chlorate(I)
The reaction of chlorine with water to form chloride ions and chlorate(I) ions. The reaction of chlorine with water to form chloride ions and oxygen.
Appreciate that society assesses the advantages and disadvantages when deciding if chemicals should be added to water supplies.
The use of chlorine in water treatment.
Appreciate the benefits to health of water treatment by chlorine outweigh its toxic effects.
The reaction of chlorine with cold, dilute, aqueous NaOH and uses of the solution formed.
Required practical 3:
Carry out simple test-tube reactions to identify:

• cations – Group 2, NH4+
• anions – Group 7 (halide ions), OH–, CO32–, SO42–
  3.2.4 PROPERTIES OF PERIOD 3 ELEMENTS AND THEIR OXIDES AND CHLORIDES (INTERNATIONAL A2)
  The reactions of the Period 3 elements with oxygen and chlorine are considered. The pH of the solutions formed when the oxides and chlorides of Period 3 elements react with water illustrates further trends in properties across this period. Explanations of these reactions offer opportunities to develop an in-depth understanding of how and why these reactions occur.

3.2.4.1 Properties of Period 3 elements and their oxides and chlorides
The reactions of Na and Mg with water.
The trends in the reactions of the elements Na, Mg, Al, Si, P and S with oxygen, limited to the formation of Na2O, MgO, Al2O3, SiO2, P4O10, SO2 and SO3
The trends in the reactions of the elements Na, Mg, Al, Si, P and S with chlorine, limited to the formation of NaCl, MgCl2, Al2Cl6, SiCl4 and PCl5
The reactions of the oxides Na2O, MgO, Al2O3, SiO2, P4O10, SO2 and SO3 and the chlorides NaCl, MgCl2, AlCl3, SiCl4 and PCl5 with water and the pH of the solutions formed.
The structures of the acids and the anions formed when the oxides P4O10, SO2 and SO3 and the chlorides SiCl4and PCl5 react with water.

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Students should be able to:

• explain the trend in the melting point of the specified oxides and chlorides in terms of their structure and bonding
• explain the trends in the reactions of the oxides and chlorides with water in terms of the type of bonding present in each specified oxide and chloride and write equations for these reactions
• write equations for the reactions that occur between the specified oxides and chlorides, and given acids and bases.
  3.2.5 TRANSITION METALS (INTERNATIONAL A2)
  The 3d block contains 10 elements, all of which are metals. Unlike the metals in Groups 1 and 2, the transition metals Ti to Cu form coloured compounds and compounds where the transition metal exists in different oxidation states. Some of these metals are familiar as catalysts. The properties of these elements are studied in this section with opportunities for a wide range of practical investigations.

3.2.5.1 General properties of transition metals
Transition metal characteristics of elements Ti–Cu arise from an incomplete sub-level in atoms or ions. The characteristic properties include:

• complex formation
• formation of coloured ions
• variable oxidation state
• catalytic activity.
  A ligand is a molecule or ion that forms a co-ordinate bond with a transition metal ion by donating a pair of electrons.
  A complex is a central metal atom or ion surrounded by ligands.
  Co-ordination number is number of co-ordinate bonds to the central metal atom or ion.

3.2.5.2 Substitution reactions
H2O, NH3 and Cl− can act as monodentate ligands.
The ligands NH3 and H2O are similar in size and are uncharged.
Exchange of the ligands NH3 and H2O occurs without change of co-ordination number (eg Co2+ and Cu2+). Substitution may be incomplete (eg the formation of [Cu(NH3)4(H2O)2]2+).
The Cl− ligand is larger than the uncharged ligands NH3 and H2O
Exchange of the ligand H2O by Cl– can involve a change of co-ordination number (eg Co2+, Cu2+ and Fe3+). Ligands can be bidentate (eg H2NCH2CH2NH2 and C2O42–).
Ligands can be multidentate (eg EDTA4–).
Haem is an iron(II) complex with a multidentate ligand.
Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin, enabling oxygen to be transported in the blood. Carbon monoxide is toxic because it replaces oxygen co-ordinately bonded to Fe(II) in haemoglobin.

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Bidentate and multidentate ligands replace monodentate ligands from complexes. This is called the chelate effect.
Students should be able to:
• explain the chelate effect, in terms of the balance between the entropy and enthalpy change in these reactions.
3.2.5.3 Shapes of complex ions
Transition metal ions commonly form octahedral complexes with small ligands (eg H2O and NH3).
Octahedral complexes can display cis–trans isomerism (a special case of E–Z isomerism) with monodentate ligands and optical isomerism with bidentate ligands.
Transition metal ions commonly form tetrahedral complexes with larger ligands (eg Cl–). Square planar complexes are also formed and can display cis–trans isomerism.
Cisplatin is the cis isomer.
Ag+ forms the linear complex [Ag(NH3)2]+ as used in Tollens’ reagent.
3.2.5.4 Formation of coloured ions
Transition metal ions can be identified by their colour.
Colour arises when some of the wavelengths of visible light are absorbed and the remaining wavelengths of light are transmitted or reflected.
d electrons move from the ground state to an excited state when light is absorbed.
The energy difference between the ground state and the excited state of the d electrons is given by:
∆E = hv = hc/λ
Changes in oxidation state, co-ordination number and ligand alter ∆E and this leads to a change in colour.
The absorption of visible light is used in spectroscopy.
A simple colorimeter can be used to determine the concentration of coloured ions in solution.
3.2.5.5 Variable oxidation states
Transition elements show variable oxidation states.
Vanadium species in oxidation states IV, III and II are formed by the reduction of vanadate(V) ions by zinc in acidic solution.
The redox potential for a transition metal ion changing from a higher to a lower oxidation state is influenced by pH and by the ligand.
The reduction of [Ag(NH3)2]+ (Tollens’ reagent) to metallic silver is used to distinguish between aldehydes and ketones.
The redox titrations of Fe2+ and C2O42– with MnO4–Students should be able to:• perform calculations for these titrations and similar redox reactions.

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3.2.5.6 Catalysts
Transition metals and their compounds can act as heterogeneous and homogeneous catalysts.
A heterogeneous catalyst is in a different phase from the reactants and the reaction occurs at active sites on the surface.
The use of a support medium to maximise the surface area of a heterogeneous catalyst and minimise the cost.
V2O5 acts as a heterogeneous catalyst in the Contact process. Fe is used as a heterogeneous catalyst in the Haber process.
Heterogeneous catalysts can become poisoned by impurities that block the active sites and consequently have reduced efficiency; this has a cost implication.
A homogeneous catalyst is in the same phase as the reactants.
When catalysts and reactants are in the same phase, the reaction proceeds through an intermediate species. Students should be able to:

• explain the importance of variable oxidation states in catalysis
• explain, with the aid of equations, how V2O5 acts as a catalyst in the Contact process
• explain, with the aid of equations, how Fe2+ ions catalyse the reaction between I− and S2O82–
• explain, with the aid of equations, how Mn2+ ions autocatalyse the reaction between C2O42– and MnO4–
  3.2.6 REACTIONS OF IONS IN AQUEOUS SOLUTION (INTERNATIONAL A2)
  The reactions of transition metal ions in aqueous solution provide a practical opportunity for students to show and to understand how transition metal ions can be identified by test-tube reactions in the laboratory.

3.2.6.1 Reactions of ions in aqueous solution
In aqueous solution, the following metal-aqua ions are formed:
[M(H2O)6]2+, limited to M = Fe and Cu
[M(H2O)6]3+, limited to M = Al and Fe
The acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+
Some metal hydroxides show amphoteric character by dissolving in both acids and bases (eg hydroxides of Al3+).
Students should be able to:

• explain, in terms of the charge/size ratio of the metal ion, why the acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+
• describe and explain the simple test-tube reactions of: M2+(aq) ions, limited to M = Fe and Cu, and of M3+(aq) ions, limited to M = Al and Fe, with the bases OH–, NH3 and CO32–
  Required practical 9:
  Carry out simple test-tube reactions to identify transition metal ions in aqueous solution.

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3.3 ORGANIC CHEMISTRY
3.3.1 INTRODUCTION TO ORGANIC CHEMISTRY (INTERNATIONAL AS)
Organic chemistry is the study of the millions of covalent compounds of the element carbon. These structurally diverse compounds vary from naturally occurring petroleum fuels to DNA and the molecules in living systems. Organic compounds also demonstrate human ingenuity in the vast range of synthetic materials created by chemists. Many of these compounds are used as drugs, medicines and plastics.
Organic compounds are named using the International Union of Pure and Applied Chemistry (IUPAC) system and the structure or formula of molecules can be represented in various different ways. Organic mechanisms are studied, which enable reactions to be explained.
In the search for sustainable chemistry, for safer agrochemicals and for new materials to match the desire for new technology, chemistry plays the dominant role.
3.3.1.1 Nomenclature
Organic compounds can be represented by:

• empirical formula
• molecular formula
• general formula
• structural formula
• displayed formula
• skeletal formula.
  The characteristics of a homologous series, a series of compounds containing the same functional group. IUPAC rules for nomenclature.
  Students should be able to:

• draw structural, displayed and skeletal formulas for given organic compounds
• apply IUPAC rules for nomenclature to name organic compounds limited to chains and rings with up to six carbon atoms each
• apply IUPAC rules for nomenclature to draw the structure of an organic compound from the IUPAC name limited to chains and rings with up to six carbon atoms each.

3.3.1.2 Reaction mechanisms
Reactions of organic compounds can be explained using mechanisms. Free-radical mechanisms:

• the unpaired electron in a radical is represented by a dot
• the use of curly arrows is not required for radical mechanisms. Students should be able to:
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Other mechanisms:

• the formation of a covalent bond is shown by a curly arrow that starts from a lone electron pair or from another covalent bond
• the breaking of a covalent bond is shown by a curly arrow starting from the bond.
  Students should be able to:

• outline mechanisms by drawing the structures of the species involved and curly arrows to represent the movement of electron pairs.
3.3.1.3 Isomerism
Structural isomerism. Stereoisomerism.
E–Z isomerism is a form of stereoisomerism and occurs as a result of restricted rotation about the planar carbon–carbon double bond.
Cahn–Ingold–Prelog (CIP) priority rules. Students should be able to:

• define the term structural isomer
• draw the structures of chain, position and functional group isomers
• define the term stereoisomer
• draw the structural formulas of E and Z isomers
• apply the CIP priority rules to E and Z isomers.
  3.3.2 ALKANES (INTERNATIONAL AS)
  Alkanes are the main constituent of crude oil, which is an important raw material for the chemical industry. Alkanes are also used as fuels and the environmental consequences of this use are considered in this section.

3.3.2.1 Fractional distillation of crude oil
Alkanes are saturated hydrocarbons.
Petroleum is a mixture consisting mainly of alkane hydrocarbons that can be separated by fractional distillation.
3.3.2.2 Modification of alkanes by cracking
Cracking involves breaking C–C bonds in alkanes.
Thermal cracking takes place at high pressure and high temperature and produces a high percentage of
alkenes (mechanism not required).
Catalytic cracking takes place at a slight pressure, high temperature and in the presence of a zeolite catalyst
and is used mainly to produce motor fuels and aromatic hydrocarbons (mechanism not required). Students should be able to:• explain the economic reasons for cracking alkanes.

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3.3.2.3 Combustion of alkanes
Alkanes are used as fuels.
Combustion of alkanes and other organic compounds can be complete or incomplete.
The internal combustion engine produces a number of pollutants including NOx, CO, carbon and unburned hydrocarbons.
These gaseous pollutants from internal combustion engines can be removed using catalytic converters. Combustion of hydrocarbons containing sulfur leads to sulfur dioxide that causes air pollution. Students should be able to:• explain why sulfur dioxide can be removed from flue gases using calcium oxide or calcium carbonate.3.3.2.4 Chlorination of alkanes
The reaction of methane with chlorine.
Students should be able to:
• explain this reaction as a free-radical substitution mechanism involving initiation, propagation and termination steps.
3.3.3 HALOGENOALKANES (INTERNATIONAL AS)
Halogenoalkanes are much more reactive than alkanes. They have many uses, including as refrigerants, as solvents and in pharmaceuticals. The use of some halogenoalkanes has been restricted due to the effect of chlorofluorocarbons (CFCs) on the atmosphere.
3.3.3.1 Nucleophilic substitution
Halogenoalkanes contain polar bonds.
Halogenoalkanes undergo substitution reactions with the nucleophiles OH–, CN– and NH3
Students should be able to:

• outline the nucleophilic substitution mechanisms of these reactions
• explain why the carbon–halogen bond enthalpy influences the rate of reaction.
  3.3.3.2 Elimination
  The concurrent substitution and elimination reactions of a halogenoalkane (eg 2-bromopropane with potassium hydroxide).
  Students should be able to:

• explain the role of the reagent as both nucleophile and base
• outline the mechanisms of these reactions.
  3.3.4 ALKENES (INTERNATIONAL AS)
  In alkenes, the high electron density of the carbon–carbon double bond leads to attack on these molecules by electrophiles. This section also covers the mechanism of addition to the double bond and introduces addition polymers, which are commercially important and have many uses in modern society.

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3.3.4.1 Structure, bonding and reactivity
Alkenes are unsaturated hydrocarbons.
Bonding in alkenes involves a double covalent bond, a centre of high electron density.
3.3.4.2 Addition reactions of alkenes
Electrophilic addition reactions of alkenes with HBr, H2SO4 and Br2
The use of bromine to test for unsaturation.
The formation of major and minor products in addition reactions of unsymmetrical alkenes.
Students should be able to:

• outline the mechanisms for these reactions
• explain the formation of major and minor products by reference to the relative stabilities of primary, secondary and tertiary carbocation intermediates.

3.3.4.3 Addition polymers
Addition polymers are formed from alkenes and substituted alkenes. The repeating unit of addition polymers.
IUPAC rules for naming addition polymers.
Addition polymers are unreactive.
Appreciate that knowledge and understanding of the production and properties of polymers has developed over time.
Typical uses of poly(chloroethene), commonly known as PVC, and how its properties can be modified using a plasticiser.
Students should be able to:

• draw the repeating unit from a monomer structure
• draw the repeating unit from a section of the polymer chain
• draw the structure of the monomer from a section of the polymer
• explain why addition polymers are unreactive
• explain the nature of intermolecular forces between molecules of polyalkenes.
  3.3.4.4 Epoxyethane
  The production of epoxyethane by the partial oxidation of ethene and understand the hazards of this process. The reactions with water and alcohols and the uses of the products formed.
  Students should be able to:

• explain the high reactivity of epoxyethane
• write equations for the reactions of epoxyethane with water and with alcohols and outline the mechanism for these reactions
• explain the economic and environmental importance of products including, surfactants and antifreeze, formed in these reactions.

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3.3.5 ALCOHOLS (INTERNATIONAL AS)Alcohols have many scientific, medicinal and industrial uses.
3.3.5.1 Oxidation of alcohols
Alcohols are classified as primary, secondary and tertiary.
Primary alcohols can be oxidised to aldehydes which can be further oxidised to carboxylic acids.
Secondary alcohols can be oxidised to ketones.
Tertiary alcohols are not easily oxidised.
Acidified potassium dichromate(VI) is a suitable oxidising agent.
Students should be able to:

• write equations for these oxidation reactions (equations showing [O] as oxidant are acceptable)
• explain how the method used to oxidise a primary alcohol determines whether an aldehyde or carboxylic acid is obtained
• use chemical tests to distinguish between aldehydes and ketones including Fehling’s solution and Tollens’ reagent.
  3.3.5.2 Elimination
  Alkenes can be formed from alcohols by acid-catalysed elimination reactions.
  Alkenes produced by this method can be used to produce addition polymers without using monomers derived from crude oil.
  Students should be able to:

• outline the mechanism for the elimination of water from alcohols.
Required practical 4:
Distil a product from a reaction.
3.3.6 ORGANIC ANALYSIS (INTERNATIONAL AS)
Our understanding of organic molecules, their structure and the way they react, has been enhanced by organic analysis. This section considers some of the analytical techniques used by chemists, including test- tube reactions and spectroscopic techniques.
3.3.6.1 Identification of functional groups by test tube reactions
The reactions of functional groups listed in the specification.
Students should be able to:
• identify the functional groups using reactions in the specification.
Required practical 5:
Carry out tests for alcohols, aldehydes, alkenes and carboxylic acids.

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3.3.6.2 Mass spectrometry
Mass spectrometry can be used to determine the molecular formula of a compound.
Students should be able to:
• use precise atomic masses and the precise molecular mass to determine the molecular formula of a compound.
3.3.6.3 Infrared spectroscopy
Bonds in a molecule absorb infrared radiation at characteristic wavenumbers.
‘Fingerprinting’ allows identification of a molecule by comparison of spectra.
Students should be able to:
• use infrared spectra and the Chemistry data booklet to identify particular bonds, and therefore functional groups, and also to identify impurities.
The link between absorption of infrared radiation by bonds in CO2, methane and water vapour and global warming.
3.3.7 OPTICAL ISOMERISM (INTERNATIONAL A2)
Compounds that contain an asymmetric carbon atom form stereoisomers that differ in their effect on plane polarised light. This type of isomerism is called optical isomerism.
3.3.7.1 Optical isomerism
Optical isomerism is a form of stereoisomerism and occurs as a result of chirality in molecules, limited to molecules with a single chiral centre.
An asymmetric carbon atom is chiral and gives rise to optical isomers (enantiomers), which exist as non super- imposable mirror images and differ in their effect on plane polarised light.
A mixture of equal amounts of enantiomers is called a racemic mixture (racemate). Students should be able to:

• draw the structural formulas and displayed formulas of enantiomers
• understand how racemic mixtures (racemates) are formed and why they are optically inactive.
  3.3.8 ALDEHYDES AND KETONES (INTERNATIONAL A2)
  Aldehydes, ketones, carboxylic acids and their derivatives all contain the carbonyl group which is attacked by nucleophiles. This section includes the addition reactions of aldehydes and ketones.

3.3.8.1 Aldehydes and ketones
Aldehydes are readily oxidised to carboxylic acids.
Chemical tests to distinguish between aldehydes and ketones including Fehling’s solution and Tollens’ reagent.
Aldehydes can be reduced to primary alcohols, and ketones to secondary alcohols, using NaBH4 in aqueous solution. These reduction reactions are examples of nucleophilic addition.
The nucleophilic addition reactions of carbonyl compounds with KCN, followed by dilute acid, to produce hydroxynitriles.

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Aldehydes and unsymmetrical ketones form mixtures of enantiomers when they react with KCN followed by dilute acid.
The hazards of using KCN.
Students should be able to:

• write overall equations for reduction reactions using [H] as the reductant
• outline the nucleophilic addition mechanism for reduction reactions with NaBH4 (the nucleophile should be shown as H–)
• write overall equations for the formation of hydroxynitriles using HCN
• outline the nucleophilic addition mechanism for the reaction with KCN followed by dilute acid
• explain why nucleophilic addition reactions of KCN, followed by dilute acid, can produce a mixture of enantiomers.
  3.3.9 CARBOXYLIC ACIDS AND DERIVATIVES (INTERNATIONAL A2)
  Carboxylic acids are weak acids but strong enough to liberate carbon dioxide from carbonates. Esters occur naturally in vegetable oils and animal fats. Important products obtained from esters include biodiesel, soap and glycerol.

3.3.9.1 Carboxylic acids and esters
The structures of:• carboxylic acids• esters.
Carboxylic acids are weak acids but will liberate CO2 from carbonates.
Carboxylic acids and alcohols react, in the presence of an acid catalyst, to give esters. Common uses of esters (eg in solvents, plasticisers, perfumes and food flavourings). Vegetable oils and animal fats are esters of propane-1,2,3-triol (glycerol).
Esters can be hydrolysed in acid or alkaline conditions to form alcohols and carboxylic acids or salts of carboxylic acids.
Vegetable oils and animal fats can be hydrolysed in alkaline conditions to give soap (salts of long-chain carboxylic acids) and glycerol.
Biodiesel is a mixture of methyl esters of long-chain carboxylic acids.
Biodiesel is produced by reacting vegetable oils with methanol in the presence of a catalyst.
3.3.9.2 Acylation
The structures of:

• acid anhydrides
• acyl chlorides
• amides.

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The nucleophilic addition–elimination reactions of water, alcohols, ammonia and primary amines with acyl chlorides and acid anhydrides.
The industrial advantages of ethanoic anhydride over ethanoyl chloride in the manufacture of the drug aspirin.
Students should be able to:
• outline the mechanism of nucleophilic addition–elimination reactions of acyl chlorides with water, alcohols, ammonia and primary amines.
Required practical 10:
Prepare a pure organic solid and test its purity.
3.3.10 AROMATIC CHEMISTRY (INTERNATIONAL A2)
Aromatic chemistry takes benzene as an example of this type of molecule and looks at the structure of the benzene ring and its substitution reactions.
3.3.10.1 Bonding
The nature of the bonding in a benzene ring, limited to planar structure and bond length intermediate between single and double.
Delocalisation of p electrons makes benzene more stable than the theoretical molecule cyclohexa-1,3,5-triene. Students should be able to:

• use thermochemical evidence from enthalpies of hydrogenation to account for this extra stability
• explain why substitution reactions occur in preference to addition reactions.
  3.3.10.2 Electrophilic substitution
  Electrophilic attack on benzene rings results in substitution, limited to monosubstitutions.
  Nitration is an important step in synthesis, including the manufacture of explosives and formation of amines.
  Sulfonation is an important step in synthesis including the manufacture of surfactant and sulfonamides.
  Friedel–Crafts acylation reactions are also important steps in synthesis.
  Free-radical attack by chlorine results in ring addition to benzene and side chain substitution in methylbenzene.
  Students should be able to:

• outline the electrophilic substitution mechanisms of nitration and sulfonation on benzene and methylbenzene; identity of the products formed in these reactions
• outline the electrophilic substitution mechanisms of acylation and alkylation using AlCl3 as a catalyst; identity of the products formed in these reactions
• outline the free-radical attack of chlorine on benzene and methylbenzene; identity of the products formed in these reactions
• explain the relative reactivities of chlorine substituted in the ring and in the side chain.

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3.3.11 AMINES (INTERNATIONAL A2)
Amines are compounds based on ammonia where hydrogen atoms have been replaced by alkyl or aryl groups. This section includes their reactions as nucleophiles.
3.3.11.1 Preparation
Primary aliphatic amines can be prepared by the reaction of ammonia with halogenoalkanes and by the reduction of nitriles.
Aromatic amines, prepared by the reduction of nitro compounds, are used in the manufacture of dyes.
3.3.11.2 Base properties
Amines are weak bases.
The difference in base strength between ammonia, primary aliphatic and primary aromatic amines.
Students should be able to:
• explain the difference in base strength in terms of the availability of the lone pair of electrons on the N atom.
3.3.11.3 Nucleophilic properties
Amines are nucleophiles.
The nucleophilic substitution reactions of ammonia and amines with halogenoalkanes to form primary, secondary, tertiary amines and quaternary ammonium salts.
The use of quaternary ammonium salts as cationic surfactants.
The nucleophilic addition–elimination reactions of ammonia and primary amines with acyl chlorides and acid anhydrides.
Students should be able to outline the mechanisms of:

• these nucleophilic substitution reactions
• the nucleophilic addition–elimination reactions of ammonia and primary amines with acyl chlorides.
  3.3.12 POLYMERS (INTERNATIONAL A2)
  The study of polymers is extended to include condensation polymers. The ways in which condensation polymers are formed are studied, together with their properties and typical uses. Problems associated with the reuse or disposal of both addition and condensation polymers are considered.
  3.3.12.1 Condensation polymers
  Condensation polymers are formed by reactions between:

• dicarboxylic acids and diols
• dicarboxylic acids and diamines
• amino acids.
  The repeating units in polyesters (eg Terylene) and polyamides (eg nylon 6,6 and Kevlar) and the linkages between these repeating units.
  Typical uses of these polymers.

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OxfordAQA International AS and A-level Chemistry (9620).
For International AS exams May/June 2017 onwards. For International A-level exams May/June 2018 onwards. Version 4.1
Students should be able to:

• draw the repeating unit from monomer structure(s)
• draw the repeating unit from a section of the polymer chain
• draw the structure(s) of the monomer(s) from a section of the polymer
• explain the nature of the intermolecular forces between molecules of condensation polymers.
  3.3.12.2 Biodegradability and disposal of polymers
  Polyalkenes are chemically inert and non-biodegradable.
  Polyesters and polyamides can be broken down by hydrolysis and are biodegradable.
  The advantages and disadvantages of different methods of disposal of polymers, including recycling. Students should be able to:

• explain why polyesters and polyamides can be hydrolysed but polyalkenes cannot.
3.3.13 AMINO ACIDS AND PROTEINS (INTERNATIONAL A2)
Amino acids, proteins and DNA are the molecules of life. In this section, the structure and bonding in these molecules and the way they interact is studied. Drug action is also considered.
3.3.13.1 Amino acids
Amino acids have both acidic and basic properties, including the formation of zwitterions. Students should be able to:• draw the structures of amino acids as zwitterions and the ions formed from amino acids:

• in acid solution
• in alkaline solution.3.3.13.2 Proteins
  Proteins are sequences of amino acids joined by peptide links.
  The importance of hydrogen bonding and sulfur–sulfur bonds in proteins.
  The primary, secondary (a-helix and b–pleated sheets) and tertiary structure of proteins.
  Hydrolysis of the peptide link produces the constituent amino acids.
  Amino acids can be separated and identified by thin-layer chromatography.
  Amino acids can be located on a chromatogram using developing agents such as ninhydrin or ultraviolet light and identified by their Rf values.
  Students should be able to:

• draw the structure of a peptide formed from up to three amino acids
• draw the structure of the amino acids formed by hydrolysis of a peptide
• identify primary, secondary and tertiary structures in diagrams
• explain how these structures are maintained by hydrogen bonding and S–S bonds
• calculate Rf values from a chromatogram.

Visit oxfordaqaexams.org.uk/9620 for the most up-to-date specification, resources, support and administration 39

OxfordAQA International AS and A-level Chemistry (9620). For International AS exams May/June 2017 onwards. For International A-level exams May/June 2018 onwards. Version 4.1
3.3.13.3 Action of anticancer drugs
The Pt(II) complex cisplatin is used as an anticancer drug.
Cisplatin prevents DNA replication in cancer cells by a ligand replacement reaction with DNA in which a bond is formed between platinum and a nitrogen atom on a protein.
Appreciate that society needs to assess the balance between the benefits and the adverse effects of drugs, such as the anticancer drug cisplatin.
Students should be able to:

• explain why cisplatin prevents DNA replication
• explain why such drugs can have adverse effects.
  3.3.14 ORGANIC SYNTHESIS (INTERNATIONAL A2)
  The formation of new organic compounds by multi-step syntheses using reactions included in the specification is covered in this section.

3.3.14 Organic synthesis
The synthesis of an organic compound can involve several steps.
Students should be able to:

• explain why chemists aim to design processes that do not require a solvent and that use non-hazardous starting materials
• explain why chemists aim to design production methods with fewer steps that have a high percentage atom economy
• use reactions in this specification to devise a synthesis, with up to four steps, for an organic compound.
  3.3.15 NUCLEAR MAGNETIC RESONANCE SPECTROSCOPY (INTERNATIONAL A2)
  Chemists use a variety of techniques to deduce the structure of compounds. In this section, nuclear magnetic resonance spectroscopy is added to mass spectrometry and infrared spectroscopy as an analytical technique. The emphasis is on the use of analytical data to solve problems rather than on spectroscopic theory.
  3.3.15.1 Nuclear magnetic resonance spectroscopy
  Appreciation that scientists have developed a range of analytical techniques which together enable the structures of new compounds to be confirmed.
  Nuclear magnetic resonance (NMR) gives information about the position of 13C or 1H atoms in a molecule.13C NMR gives simpler spectra than 1H NMR.
  The use of the δ scale for recording chemical shift.
  Chemical shift depends on the molecular environment.
  Integrated spectra indicate the relative numbers of 1H atoms in different environments.1H NMR spectra are obtained using samples dissolved in deuterated solvents or CCl4The use of tetramethylsilane (TMS) as a standard.

40 Visit oxfordaqaexams.org.uk/9620 for the most up-to-date specification, resources, support and administration

OxfordAQA International AS and A-level Chemistry (9620).
For International AS exams May/June 2017 onwards. For International A-level exams May/June 2018 onwards. Version 4.1
Students should be able to:

• explain why TMS is a suitable substance to use as a standard
• use 1H NMR and 13C NMR spectra and chemical shift data from the Chemistry data booklet to suggest possible structures or part structures for molecules
• use integration data from 1H NMR spectra to determine the relative numbers of equivalent protons in the molecule
• use the n+1 rule to deduce the spin–spin splitting patterns of adjacent, non-equivalent protons, limited to doublet, triplet and quartet formation in aliphatic compounds.
  3.3.16 CHROMATOGRAPHY (INTERNATIONAL A2)
  Chromatography provides an important method of separating and identifying components in a mixture. Different types of chromatography are used depending on the composition of mixture to be separated.
  3.3.16.1 Chromatography
  Chromatography can be used to separate and identify the components in a mixture.
  Types of chromatography include:
  
  • thin-layer chromatography (TLC) – a plate is coated with a solid and a solvent moves up the plate
  • column chromatography (CC) – a column is packed with a solid and a solvent moves down the column
  • gas chromatography (GC) – a column is packed with a solid or with a solid coated by a liquid, and a gas is passed through the column under pressure at high temperature.
    Separation depends on the balance between solubility in the moving phase and retention by the stationary phase.
    Retention times and Rf values are used to identify different substances.
    The use of mass spectrometry to analyse the components separated by GC. Students should be able to:
  • calculate Rf values from a chromatogram
  • compare retention times and Rf values with standards to identify different substances.