AQA CHEMISTRY A-LEVEL Class 6
ACADEMIC YEAR 2020-2021
Teachers: ………… AZAS MICHAEL, RANIA MAVROMMATI……………………………………..
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PERIODS |
LEARNING OBJECTIVES
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METHODOLOGY/ ACTIVITIES | Textbook Reference/ Assignments/Homework
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INDUCTION 25-27/8
Term 1 1/9 – 30/11
Week 1 (Rania) 2 periods + 2 periods practical preliminary
Week 2-4 Rania 4 periods + 4 periods preliminary practical
Week 1 Mr. Azas 5 periods
Week 5+6 Rania Periods 4 + 4 periods preliminary practical
Week 7+8 Rania 4 periods + 4 periods preliminary practicals
Week 9+10 Rania Periods 4 + 4 periods assessed practical 1
Week 2+3 Mr. Azas 10 periods
Week 4 Mr. Azas 5 periods
Week 6 Mr. Azas 5 periods
Week 5 Mr. Azas 5 periods
Week 7 Mr. Azas 5 periods
Week 11-13 Rania Periods 6 + 6 periods preliminary practical and assessed practical 2
Week 14 Rania 2 periods + 2 preliminary practical
Term 2 1/12 – 28/2
Week 15 Rania 2 periods
CHRISTMAS HOLIDAYS
Week 16-18 Rania 6 periods + 6 periods preliminary practical and assessed practical 3
Week 19-20 Rania 4 periods + 4 preliminary practical, assessed practical 4
Week 8 Mr. Azas Periods 5
Week 10 Mr. Azas 5 periods
Week 9 Mr. Azas 5 periods
Week 10 Mr. Azas 5 periods
Week 11-12 Mr. Azas 10 periods
Week 13 Mr. Azas 5 periods
Week 14 +15 Mr. Azas 10 periods
Weeks 16-17 Mr. Azas 10 periods
Week 18-19 Mr. Azas 10 periods
Week 20 Mr. Azas 5 periods
Week 21-30 Mr. Azas and Rania
Term 3 1/3 – 31/5
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INDUCTION LESSONS
3.1 Physical chemistry 3.1.1 Atomic structure 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
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.
Students should be able to: • interpret simple mass spectra of elements • calculate relative atomic mass from isotopic abundance, limited to mononuclear ions.
3.1.1.3 Electron configuration
Students should be able to: • define first ionisation energy • write equations for first and successive ionisation energies • explain how first and successive ionisation energies in Period 3 (Na–Ar) and in Group 2 (Be–Ba) give evidence for electron configuration in sub-shells and in shells.
3.1.2 Amount of substance 3.1.2.1 Relative atomic mass and relative molecular mass
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
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
Students should be able to use the equation in calculations. Students will not be expected to recall the value of the gas constant, R.
3.1.2.4 Empirical and molecular formula
Students should be able to: • calculate empirical formula from data giving composition by mass or percentage by mass • calculate molecular formula from the empirical formula and relative molecular mass.
3.1.2.5 Balanced equations and associated calculations
Students should be able to: • write balanced equations for reactions studied • balance equations for unfamiliar reactions when reactants and products are specified. 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.
3.1.3 Bonding
3.1.3.1 Ionic bonding
Students should be able to: • predict the charge on a simple ion using the position of the element in the Periodic Table • construct formulas for ionic compounds.
3.1.3.2 Nature of covalent and dative covalent bonds
Students should be able to represent: • a covalent bond using a line • a co-ordinate bond using an arrow.
3.1.3.3 Metallic bonding
3.1.3.4 Bonding and physical properties
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
Students should be able to explain the shapes of, and bond angles in, simple molecules and ions with up to six electron pairs (including lone pairs of electrons) surrounding the central atom.
3.1.3.6 Bond polarity
Students should be able to: • use partial charges to show that a bond is polar • explain why some molecules with polar bonds do not have a permanent dipole.
3.1.3.7 Forces between molecules
Students should be able to: • explain the existence of these forces between familiar and unfamiliar molecules • explain how melting and boiling points are influenced by these intermolecular forces.
3.1.4 Energetics
3.1.4.1 Enthalpy change
Students should be able to: • define standard enthalpy of combustion (∆c HƟ) • define standard enthalpy of formation (∆f HƟ).
3.1.4.2 Calorimetry
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.
3.1.4.3 Applications of Hess’s law
Students should be able to use Hess’s law to perform calculations, including calculation of enthalpy changes for reactions from enthalpies of combustion or from enthalpies of formation.
3.1.4.4 Bond enthalpies
Students should be able to: • define the term mean bond enthalpy • use mean bond enthalpies to calculate an approximate value of ∆H for reactions in the gaseous phase • explain why values from mean bond enthalpy calculations differ from those determined using Hess’s law.
3.1.5 Kinetics
3.1.5.1 Collision theory 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
Students should be able to draw and interpret distribution curves for different temperatures.
3.1.5.3 Effect of temperature on reaction rate
Students should be able to use the Maxwell–Boltzmann distribution to explain why a small temperature increase can lead to a large increase in rate.
3.1.5.4 Effect of concentration and pressure
Students should be able to explain how a change in concentration or a change in pressure influences the rate of a reaction.
3.1.5.5 Catalysts Students should be able to use a Maxwell–Boltzmann distribution to help explain how a catalyst increases the rate of a reaction involving a gas.
3.1.6 Chemical equilibria, Le Chatelier’s principle and K c
3.1.6.1 Chemical equilibria and Le Chatelier’s principle
Students should be able to: • use Le Chatelier’s principle to predict qualitatively the effect of changes in temperature, pressure and concentration on the position of equilibrium • explain why, for a reversible reaction used in an industrial process, a compromise temperature and pressure may be used.
3.1.6.2 Equilibrium constant Kc for homogeneous systems
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
3.2 Inorganic chemistry 3.2.1 Periodicity
3.2.1.1 Classification
3.2.1.2 Physical properties of Period 3 elements Students should be able to: • explain the trends in atomic radius and first ionisation energy • explain the melting point of the elements in terms of their structure and bonding.
3.2.2 Group 2, the alkaline earth metals 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 TiCl4 The 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.
3.2.4 Properties of Period 3 elements and their oxides
Students should be able to: • explain the trend in the melting point of the oxides of the elements Na–S in terms of their structure and bonding • explain the trends in the reactions of the oxides with water in terms of the type of bonding present in each oxide • write equations for the reactions that occur between the oxides of the elements Na–S and given acids and bases.
3.1.8 Thermodynamics (A-level only)
3.1.8.1 Born–Haber cycles (A-level only)
Students should be able to: • define each of the above terms and lattice enthalpy • construct Born–Haber cycles to calculate lattice enthalpies using these enthalpy changes • construct Born–Haber cycles to calculate one of the other enthalpy changes • compare lattice enthalpies from Born–Haber cycles with those from calculations based on a perfect ionic model to provide evidence for covalent character in ionic compounds. Cycles are used to calculate enthalpies of solution for ionic compounds from lattice enthalpies and enthalpies of hydration. 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 (A-level only)
Students should be able to: • calculate entropy changes from absolute entropy values • use the relationship ∆G = ∆H – T∆S to determine how ∆G varies with temperature • use the relationship ∆G = ∆H – T∆S to determine the temperature at which a reaction becomes feasible.
3.2.3 Group 7(17), the halogens
3.2.3.1 Trends in properties
Students should be able to: • explain the trend in electronegativity • explain the trend in the boiling point of the elements in terms of their structure and bonding. The trend in oxidising ability of the halogens down the group, including displacement reactions of halide ions in aqueous solution. 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)
3.1.7 Oxidation, reduction and redox equations
Students should be able to: • work out the oxidation state of an element in a compound or ion from the formula • write half-equations identifying the oxidation and reduction processes in redox reactions • combine half-equations to give an overall redox equation.
3.1.9 Rate equations
3.1.9.1 Rate equations (A-level only) 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.
3.1.9.2 Determination of rate equation
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.
3.1.10 Equilibrium constant K p for homogeneous systems (A-level only) Students should be able to: • derive partial pressure from mole fraction and total pressure • construct an expression for K p for a homogeneous system in equilibrium • perform calculations involving K p • 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 K p • understand that, whilst a catalyst can affect the rate of attainment of an equilibrium, it does not affect the value of the equilibrium constant.
3.3 Organic chemistry 3.3.1 Introduction to organic chemistry
3.3.1.1 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.
3.3.1.3 Isomerism Students should be able to: • define the term structural isomer • draw the structures of chain, position and functional group isomers • define the term stereoisomer • draw the structural formulas of E and Z isomers • apply the CIP priority rules to E and Z isomers.
3.3.2 Alkanes
3.3.2.1 Fractional distillation of crude oil
3.3.2.2 Modification of alkanes by cracking
Students should be able to explain the economic reasons for cracking alkanes.
3.3.2.3 Combustion of alkanes
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
Students should be able to explain this reaction as a free-radical substitution mechanism involving initiation, propagation and termination steps.
3.3.3 Halogenoalkanes
3.3.3.1 Nucleophilic substitution
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
Students should be able to: • explain the role of the reagent as both nucleophile and base • outline the mechanisms of these reactions.
3.3.3.3 Ozone depletion
Students should be able to use equations, such as the following, to explain how chlorine atoms catalyse decomposition of ozone: Cl• + O3 → ClO• + O2 and ClO• + O3 → 2O2 + Cl•
3.3.4 Alkenes
3.3.4.1 Structure, bonding and reactivity
3.3.4.2 Addition reactions of alkenes Students should be able to: •• outline the mechanisms for these reactions •• explain the formation of major and minor products by reference to the relative stabilities of primary, secondary and tertiary carbocation intermediates.
3.3.4.3 Addition polymers 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.
3.3.5.1 Alcohol production Students should be able to: •• explain the meaning of the term biofuel •• justify the conditions used in the production of ethanol by fermentation of glucose •• write equations to support the statement that ethanol produced by fermentation is a carbon-neutral fuel and give reasons why this statement is not valid •• outline the mechanism for the formation of an alcohol by the reaction of an alkene with steam in the presence of an acid catalyst •• discuss the environmental (including ethical) issues linked to decision making about biofuel use.
3.3.5.2 Oxidation of alcohols Students should be able to: •• write equations for these oxidation reactions (equations showing [O] as oxidant are acceptable) •• explain how the method used to oxidise a primary alcohol determines whether an aldehyde or carboxylic acid is obtained •• use chemical tests to distinguish between aldehydes and ketones including Fehling’s solution and Tollens’ reagent.
3.3.5.3 Elimination Students should be able to outline the mechanism for the elimination of water from alcohols.
3.3.6.1 Identification of functional groups by test-tube reactions Students should be able to identify the functional groups using reactions in the specification.
3.3.9 Carboxylic acids and derivatives
3.3.9.1 Carboxylic acids and esters |
3.1 Physical chemistry 3.1.1 Atomic structure 3.1.1.1 Fundamental particles
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 in which 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.
3.1.1.2 Mass number and isotopes
Mass number (A) and atomic (proton) number (Z).
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.
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.
3.1.2 Amount of substance 3.1.2.1 Relative atomic mass and relative molecular mass
When chemists measure out 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.
Relative atomic mass and relative molecular mass in terms of 12C. The term relative formula mass will be used for ionic compounds.
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.
3.1.2.3 The ideal gas equation
The ideal gas equation pV = nRT with the variables in SI units.
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.
3.1.2.5 Balanced equations and associated calculations
Equations (full and ionic). Percentage atom economy is: molecular mass of desired product sum of molecular masses of all reactants × 100 Economic, ethical and environmental advantages for society and for industry of developing chemical processes with a high atom economy
Required practical 1: Make up a volumetric solution and carry out a simple acid–base titration.
3.1.3 Bonding The physical and chemical properties of compounds depend on the ways in which the compounds are held together by chemical bonds and by 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.
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.
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.
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.
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.
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.
3.1.4 Energetics 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 ∆H298 Ɵ).
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
Required practical 2: Measurement of an enthalpy change.
3.1.4.3 Applications of Hess’s law Hess’s law.
3.1.4.4 Bond enthalpies
Mean bond enthalpy
3.1.5 Kinetics 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
3.1.5.2 Maxwell–Boltzmann distribution
Maxwell–Boltzmann distribution of molecular energies in gases.
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.
Required practical 3: Investigation of how the rate of a reaction changes with temperature.
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.
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.
3.1.6 Chemical equilibria, Le Chatelier’s principle and K c 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.
3.1.6.2 Equilibrium constant Kc for homogeneous systems The 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.
3.2 Inorganic chemistry 3.2.1 Periodicity 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.
3.2.2 Group 2, the alkaline earth metals 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. The trends in atomic radius, first ionisation energy and melting point of the elements Mg–Ba
3.2.4 Properties of Period 3 elements and their oxides The reactions of the Period 3 elements with oxygen are considered. The pH of the solutions formed when the oxides 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. 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 Na2 O, MgO, Al2 O3 , SiO2 , P4 O10, SO2 and SO3 The trend in the melting point of the highest oxides of the elements Na–S The reactions of the oxides of the elements Na–S with water, limited to Na2 O, MgO, Al2 O3 , SiO2 , P4 O10, SO2 and SO3 , and the pH of the solutions formed. The structures of the acids and the anions formed when P4 O10, SO2 and SO3 react with water.
3.1.8 Thermodynamics (A-level only) 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 (A-level only) 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.
3.1.8.2 Gibbs free-energy change, ∆G, and entropy change, ∆S (A-level only) ∆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.
3.2.3 Group 7(17), the halogens 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.
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 that 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 4: Carry out simple test-tube reactions to identify: • cations – Group 2, NH4 + • anions – Group 7 (halide ions), OH– , CO3 2–, SO4 2–
3.1.7 Oxidation, reduction and redox equations 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.
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.
3.1.9 Rate equations 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]n where 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.
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.
Required practical 7: Measuring the rate of reaction: • by an initial rate method • by a continuous monitoring method.
3.1.10 Equilibrium constant K p for homogeneous systems (A-level only) The further study of equilibria considers how the mathematical expression for the equilibrium constant K p enables 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. The equilibrium constant K p is deduced from the equation for a reversible reaction occurring in the gas phase. K p is the equilibrium constant calculated from partial pressures for a system at constant temperature.
3.3 Organic chemistry 3.3.1 Introduction to organic chemistry 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.
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.
3.3.2 Alkanes 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).
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.
3.3.2.4 Chlorination of alkanes
The reaction of methane with chlorine.
3.3.3 Halogenoalkanes 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
3.3.3.2 Elimination The concurrent substitution and elimination reactions of a halogenoalkane (eg 2-bromopropane with potassium hydroxide).
3.3.3.3 Ozone depletion
Ozone, formed naturally in the upper atmosphere, is beneficial because it absorbs ultraviolet radiation. Chlorine atoms are formed in the upper atmosphere when ultraviolet radiation causes C–Cl bonds in chlorofluorocarbons (CFCs) to break. Chlorine atoms catalyse the decomposition of ozone and contribute to the hole in the ozone layer. Appreciate that results of research by different groups in the scientific community provided evidence for legislation to ban the use of CFCs as solvents and refrigerants. Chemists have now developed alternative chlorine-free compounds.
3.3.4 Alkenes 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.
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.
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.
3.3.5 Alcohols Alcohols have many scientific, medicinal and industrial uses. Ethanol is one such alcohol and it is produced using different methods, which are considered in this section. Ethanol can be used as a biofuel.
3.3.5.1 Alcohol production Alcohols are produced industrially by hydration of alkenes in the presence of an acid catalyst. Ethanol is produced industrially by fermentation of glucose. The conditions for this process. Ethanol produced industrially by fermentation is separated by fractional distillation and can then be used as a biofuel.
3.3.5.2 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.
3.3.5.3 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.
Required practical 5 Distillation of a product from a reaction.
3.3.6.1 Identification of functional groups by test-tube reactions
The reactions of functional groups listed in the specification.
Required practical 6 Tests for alcohol, aldehyde, alkene and carboxylic acid.
3.3.9 Carboxylic acids and derivatives 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. |
Reference to book pages
Students carry out calculations using numbers in standard and ordinary form eg using the Avogadro constant. Students carry out calculations using the Avogadro constant. Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. Students understand that calculated results can only be reported to the limits of the least accurate measurement.
AT a, b and k Students could be asked to find the Mr of a volatile liquid. Students understand that the correct units need to be in pV = nRT. Students carry out calculations with the ideal gas equation, including rearranging the ideal gas equation to find unknown quantities.
AT a and k Students could be asked to find the empirical formula of a metal oxide.
AT a, d, e, f and k Students could be asked to find: • the concentration of ethanoic acid in vinegar • the mass of calcium carbonate in an indigestion tablet • the Mr of MHCO3 • the Mr of succinic acid • the mass of aspirin in an aspirin tablet • the yield for the conversion of magnesium to magnesium oxide • the Mr of a hydrated salt (eg magnesium sulfate) by heating to constant mass. AT a and k Students could be asked to find the percentage conversion of a Group 2 carbonate to its oxide by heat. AT d, e, f and k Students could be asked to determine the number of moles of water of crystallisation in a hydrated salt by titration. Students construct and/or balance equations using ratios. Students calculate percentage yields and atom economies of reactions. Students select appropriate titration data (ie identify outliers) in order to calculate mean titres. Students determine uncertainty when two burette readings are used to calculate a titre value.
AT a, b, h and k Students could be asked to find the type of structure of unknowns by experiment (eg to test solubility, conductivity and ease of melting).
Students could be given familiar and unfamiliar examples of species and asked to deduce the shape according to valence shell electron pair repulsion (VSEPR) principles.
AT d and k Students could try to deflect jets of various liquids from burettes to investigate the presence of different types and relative size of intermolecular forces.
Students understand that the correct units need to be used in q = mc∆T Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. Students understand that calculated results can only be reported to the limits of the least accurate measurement.
AT a and k Students could be asked to find ∆H for a reaction by calorimetry. Examples of reactions could include: • dissolution of potassium chloride • dissolution of sodium carbonate • neutralising NaOH with HCl • displacement reaction between CuSO4 + Zn • combustion of alcohols
Students carry out Hess’s law calculations. AT a and k Students could be asked to find ∆H for a reaction using Hess’s law and calorimetry, then present data in appropriate ways. Examples of reactions could include: • thermal decomposition of NaHCO3 • hydration of MgSO4 • hydration of CuSO4
Students understand that bond enthalpies are mean values across a range of compounds containing that bond.
AT a, b, k and l Students could investigate the effect of temperature on the rate of reaction of sodium thiosulfate and hydrochloric acid by an initial rate method. Research opportunity Students could investigate how knowledge and understanding of the factors that affect the rate of chemical reaction have changed methods of storage and cooking of food.
AT a, e, k and i Students could investigate the effect of changing the concentration of acid on the rate of a reaction of calcium carbonate and hydrochloric acid by a continuous monitoring method.
Students could carry out test-tube equilibrium shifts to show the effect of concentration and temperature (eg Cu(H2 O)6 2+ with concentrated HCl).
Students estimate the effect of changing experimental parameters on a measurable value eg how the value of Kc would change with temperature, given different specified conditions. Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. Students understand that calculated results can only be reported to the limits of the least accurate measurement. Students calculate the concentration of a reagent at equilibrium. Students calculate the value of an equilibrium constant Kc PS 2.3 Students could determine the equilibrium constant, Kc , for the reaction of ethanol with ethanoic acid in the presence of a strong acid catalyst to ethyl ethanoate.
AT c and k Students could test the reactions of Mg–Ba with water and Mg with steam and record their results. AT d and k Students could test the solubility of Group 2 hydroxides by mixing solutions of soluble Group 2 salts with sodium hydroxide and record their results. Students could test the solubility of Group 2 sulfates by mixing solutions of soluble Group 2 salts with sulfuric acid and record their results. Students could test for sulfate ions using acidified barium chloride and record their results. Research opportunity Students could investigate the use of BaSO4 in medicine.
AT a, c and k Students could carry out reactions of elements with oxygen and test the pH of the resulting oxides.
Students could be asked to find ∆S for vaporization of water using a kettle. Students rearrange the equation ∆G = ∆H – T∆S to find unknown values. Students determine ∆S and ∆H from a graph of ∆G versus T
AT d and k Students could carry out test-tube reactions of solutions of the halogens (Cl2 , Br2 , I2 ) with solutions containing their halide ions (eg KCl, KBr, KI). Students could record observations from reactions of NaCl, NaBr and NaI with concentrated sulfuric acid. Students could carry out tests for halide ions using acidified silver nitrate, including the use of ammonia to distinguish the silver halides formed.
Research opportunity Students could investigate the treatment of drinking water with chlorine. Students could investigate the addition of sodium fluoride to water supplies.
Students use given rate data and deduce a rate equation, then use some of the data to calculate the rate constant including units. Rate equations could be given and students asked to calculate rate constant or rate. Students use a graph of concentration–time and calculate the rate constant of a zero-order reaction by determination of the gradient.
Students could determine the order of reaction for a reactant in the iodine clock reaction. Students could be given data to plot and interpret in terms of order with respect to a reactant. Alternatively, students could just be given appropriate graphs and asked to derive order(s). Students calculate the rate constant of a zero-order reaction by determining the gradient of a concentration–time graph. Students plot concentration–time graphs from collected or supplied data and draw an appropriate best-fit curve. Students draw tangents to such curves to deduce rates at different times.
Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. Students understand that calculated results can only be reported to the limits of the least accurate measurement. Students calculate the partial pressures of reactants and products at equilibrium. Students calculate the value of an equilibrium constant K p
Students could be given the structure of one isomer and asked to draw further isomers. Various representations could be used to give the opportunity to identify those that are isomeric. Students understand the origin of E–Z isomerism. Students draw different forms of isomers.
AT a, d and k Fractional distillation of a crude oil substitute.
AT a, b and k Students could follow instructions when carrying out test-tube hydrolysis of halogenoalkanes to show their relative rates of reaction. AT d, g and k Students could prepare a chloroalkane, purifying the product using a separating funnel and distillation.
Research opportunity Students could investigate the role of chemists in the introduction of legislation to ban the use of CFCs and in finding replacements.
AT d and k PS 4.1 Students could test organic compounds for unsaturation using bromine water and record their observations.
AT k PS 1.2 Making poly(phenylethene) from phenylethene.
AT a, d and k PS 1.2 Students could produce ethanol by fermentation, followed by purification by fractional distillation.
AT b, d and k Students could carry out the preparation of an aldehyde by the oxidation of a primary alcohol. Students could carry out the preparation of a carboxylic acid by the oxidation of a primary alcohol.
AT b, d, g and k PS 4.1 Students could carry out the preparation of cyclohexene from cyclohexanol, including purification using a separating funnel and by distillation.
AT b, d and k PS 2.2, 2.3 and 4.1 Students could carry out test-tube reactions in the specification to distinguish alcohols, aldehydes, alkenes and carboxylic acids.
AT b, d, g and k PS 4.1 Students could make esters by reacting alcohols with carboxylic acids, purifying the product using a separating funnel and by distillation. AT b, d, g, h and k Students could identify an ester by measuring its boiling point, followed by hydrolysis to form the carboxylic acid, which is purified by recrystallisation, and determine its melting point. AT b, c, d and k Students could make soap. AT b and k Students could make biodiesel.
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