The mole as applied to electrons, atoms, molecules, ions, formulae 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, relative molecular mass 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. MS 0.4: Students carry out calculations using the Avogadro constant. MS 1.1: 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. 3 Mole Calculations Recap of calculations involving moles, masses and concentration. Calculate the mass given the number of moles for a range of different compounds Calculate the numbers of moles given the mass in a range of different compounds Calculate concentrations of liquids from masses and moles. 2.2 4 Solutions Describe how to make up a standard solution. Make up different solutions of known concentration. Become familiar with a range of calculations involving solutions in terms of masses, moles and concentration. Describe how to make a standard solution Explain how to identify the concentration of an unknown solutions by comparing the colour against solutions of concentration out different types of calculations involving solutions. 2.7 5 Ideal gas equation Introduction to the idea of an gas Relate the ideal gas equation to experiments by Charles and Boyle. Describe the differences between an gas and a real gas Explain the effects of temperature and pressure on a gas Carry out calculations using the ideas gas equation 3.1.2.3 The 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. AT a, b and k PS 3.2: Students could be asked to find the Mr of a volatile liquid. MS 0.0: Students understand that the correct units need to be in pV = nRT MS 2.2, 2.3 and 2.4: Students carry out calculations with the ideal gas equation, including 2.3 6 Relative molecular mass of a volatile liquid 2.3
rearranging the ideal gas equation to find unknown quantities. Use an experimental method to make measurements to calculate the Mr of hexane. Describe an experiment to determine the molecular mass of a volatile liquid Make precise measurements and use sensitive equipment like gas syringes with care. Use experimental data and the ideal gas equation to calculate the Mr of a volatile liquid 7 Calculating empirical formula Recap of the definition of empirical formula Calculation of empirical formula using an experimental methods Explain the difference between empirical formula and molecular formula Describe how to find the empirical mass of a compound by direct reaction and by continuous variation. Use empirical data to calculate the empirical formula. 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: empirical formula from data giving composition by mass or data giving percentage by mass; calculate molecular formula from the empirical formula and relative molecular mass. AT a and k PS 2.3 and 3.3: Students be asked to find the empirical formula of a metal oxide. 2.4 8 Determine the formulae of compounds Introduction to techniques to calculate molecular formula from empirical formula, molecular mass, data from CHNS analysers and high level mass spectrometry Calculate molecular formula from the empirical formula and relative molecular mass Describe how CHNS elemental analysers and high level resolution mass spectrometry are used to find the molecular mass Calculate molecular formula from data from CHNS and high level mass spectrometry. 2.4 9 Balanced chemical equations Describe how to use titrations to determine the stoichiometry of neutralisation reactions 3.1.2.5 Balanced equations and associated calculations 2.5
Recap of features of balanced symbol equations. How to write more complicated balanced symbol and ionic equations. Write balanced equations for a range of different reactions given the formula of reactants and products Write balanced symbol and ionic equations from word equations 10 Yields from reactions percentage yield and atom economy. Carry out a number of calculations of percentage yield from empirical and given data. Use empirical data to calculate percentage yield Calculate the theoretical and percentage yield from balanced symbol equations. Calculate percentage atom economy from balanced symbol equations 2.5 2.6 11 Thermal decomposition Carry out the thermal decomposition of calcium carbonate, hydrated copper (II) sulfate and calculate percentage yields and percentage atom economies. Identify and describe what is by a thermal decomposition reactions Calculate percentage yields from thermal decomposition reactions Calculate the formula of hydrates salts from thermal decomposition reactions. 2.5 12 Titrations Learn how to prepare a standard solution and carry out a titration accurately. Use a titration to calculate concentration. Use equipment with care and precision to prepare a standard solution Carry out a titration to produce accurate and reliable data, calculating the percentage measurement error Use a titration to calculate the concentration of a solution 2.7 Required Practical 1: Make up a volumetric solution and carry out a simple titration. base
13 Ethanoic acid in vinegar Use a titration to determine the percentage composition of ethanoic acid in vinegar Describe how to determine the concentration of ethanoic acid in vinegar out a titration to produce accurate and reliable data, calculating the percentage measurement error Use data from the titration to calculate the percentage concentration of ethanoic acid in vinegar 2.7 14 Relative formula of a hydrated salt Use a titration between hydrated sodium carbonate and hydrochloric acid to determine the relative formula of the hydrated salt Describe how use a titration to determine the formula of hydrated sodium carbonate Carry out the titration to produce accurate and reliable data Use results to calculate the formula of the hydrated sodium carbonate 2.7 15 Relative formula mass of succinic acid out a titration between succinic acid and sodium hydroxide to determine the molecular formula of succinic acid. Describe how use a titration to calculate the relative formula of succinic acid Carry out the titration to produce accurate and reliable data Use reliable results and the balanced equation to determine the molecular formula of succinic acid 2.7
1 Ionic bonds Recap of what is meant by an ionic bond and use dot and cross diagrams of electronic structures and formulae to represent ionic compounds explain what is meant by “ionic bonding” draw appropriate dot and cross diagrams to show electronic structures of ions in an ionic compound determine the formula of simple and complex ionic compounds 3.1.3.1 Ionic bonding Ionic bonding involves electrostatic attraction between oppositely charged ions in a lattice. The formulas of compound ions e.g. 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 2 Covalent Use appropriate conventions to display the bonds in covalent compounds. Introduce the idea of bond polarity and dative covalent/ co-ordinate bonds Draw appropriate dot and cross diagrams to represent covalent bonds compound, including dative covalent bonds. use the idea of electronegativity to explain why electron density in covalent bonds may not be symmetrical Explain a molecule may contain polar bonds, but the molecules itself may or may not be polar. 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. 3.2 3 Metals Recap on metallic bonding. Compare properties of substances which are bonded by ionic, covalent or metallic bonding. identify whether elements or compounds have covalent or metallic bonding explain the bonding and structure of metals use ideas about bonding to explain the properties of metals 3.3 4 Crystal structure Describe different crystal structures. Carry out a practical to grow make observations of the growth of a range of different crystals including silver, 3.4 Bonding and physical properties 3.4
and density to explain the formation of hydrogen bonds explain the formation of hydrogen bonds explain why hydrogen bonds are the strongest of the intermolecular bonds (van der Waals’, dispersion and London) forces; hydrogen bonding. The melting and boiling points of molecular substances are influenced by the strength of 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. investigate the presence of different types and relative size of intermolecular forces. 9 Revision of ideas about bonding Groups work together to produce a large display about bonding. The features of each type of bonding with suitable examples, without external resources if possible. models should be made from everyday materials to accompany the display. give examples of substances which exhibit each type of bonding discussed Explain in detail how each type of bonding works and which types of elements are involved Compare and contrast the properties of substance which exhibit different types of bonding. 3.1 to 3.7 One hour lessons Learning Outcomes Specification Content Skills Covered Student Book Section Required Practicals CHAPTER 4 –The Periodic Table (3 hours) 1 Classifying elements Recap ideas about ionisation energy and energy shells. Use data about first ionisation to make predictions about describe how the Periodic Table is organised classify an element as s, p, d or f according to its position in the Periodic Table 3.2.1.1 Classification An element is classified as s, p, d or f according to its position in the Periodic Table, which is determined by its proton number 4.1
One hour lessons Learning Outcomes Specification Content Skills Covered Student Book Section Required Practicals CHAPTER 5 – Introduction to organic chemistry (9 hours) 1 Organic and inorganic Introduce the idea of organic compounds Explain why C has its own branch of chemistry using ideas about valency and mean bond enthalpy. Identify some organic compounds Explain the key differences between organic compounds and inorganic compounds explain why carbon is able to form a number of different compounds 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. 5.1 5.2 2 Types of formulae Introduce the and conventions for drawing different types of formulae and explain when they are appropriate to use. use empirical formulae, molecular formulae, general formulae, structural formulae, displayed formulae and skeletal formulae to represent organic compounds Explain the differences between the different types of formulae Convert from one type of formulae to another 5.3 3 Molecular models and drawing structures Use 3-D models to represent simple organic molecules and use these to draw 3-D representations of these molecules. use molecular kits to build models of organic molecules recognise the dimensional aspect of molecules use appropriate notation to represent the three dimensional aspect of molecules 5.2 4 Functional groups 5.4
Introduce the stages of the free radical mechanism. Explain the rules for using curly arrows in reaction mechanisms. explain free radical mechanisms write balanced equations for the steps in a free radical mechanism use curly arrows to show the movement of electrons during a chemical reaction 8 Isomers Introduce what an isomer is by showing different examples. Explain the key differences between structural and stereoisomers. explain the differences between structural and stereoisomers Draw the structures of chain, position, functional group and E and Z isomers Apply CIP priority rules to E and Z isomers 5.7 9 Melting points and boiling points of isomers Practical investigation of the melting points of propan-1-ol and and use of data books to analyse the melting and boiling points of different isomers. These are related to the structure of the isomers to explain the differences. out a practical investigation safely measure the boiling point of an organic compound explain why isomers have different boiling and melting points 5.7 One hour lessons Learning Outcomes Specification Content Skills Covered Student Book Section Required Practicals CHAPTER 6 – The alkanes (8 hours)
for complete combustion. Recap the differences between complete and incomplete combustion write balanced chemical equations for complete and incomplete combustion reactions 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 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. 5 Measuring energy when alkanes burn Use calorimetry to measure the energy changes during the combustion of n hexane, heptane and octane and calculate the enthalpy change for each. describe how to carry out a practical investigation to measure the energy released when alkanes burn calculate the enthalpy change from experimental data calculate the measurement uncertainties and identify the systematic errors 6.3 6 Problems with alkane combustion Explore the different products formed by the internal combustion engine, the problems they cause and ways to limit their impact. Name the pollutants produced by the internal combustion engine and describe the problems they cause. explain how catalytic converters remove the pollutants from the internal combustion engine explain why sulfur dioxide can be removed from flue gases using calcium oxide or calcium carbonate 6.4
7 Chlorination of alkanes Explain the stages of the free radical mechanism during homolytic fission of a chloroalkane. write equations to show the free radical substitution of methane with chlorine identify the initiation, propagation and termination in the free radical mechanism Explain why ultraviolet light can be used to break the C-Cl bond using E= hv 3.3.2.4 Chlorination of alkanes The reaction of methane with chlorine. Students should be able to explain this reaction as a substitution mechanism involving initiation, propagation and termination steps. 6.5 8 Debating the uses of alkanes are fuels Work in groups to prepare a debate sources to explain whether or alkanes should remain our primary source fuel, assuming that supplies do not run out. Evidence from chemical reactions and other sources must be used. Research can be carried out to find further evidence. Groups should prepare arguments for both sides of the debate. Explain the importance of alkanes as and how they are derived Explain all the problems surrounding their use and the impact these have on humans and other organisms Use evidence well in conjunction with reasoned arguments 6.1 – 6.5 One hour lessons Learning Outcomes Specification Content Skills Covered Student Book Section Required Practicals
CHAPTER 7 – Energetics (9 hours) 1 Enthalpy change Recap energy level diagrams and introduce the idea of enthalpy. Define standard enthalpies of combustion and formation and use them to calculate relative molecular mass. describe reactions as endothermic or exothermic explain that enthalpy change (ΔH) is the energy transferred by heating or cooling, measured under conditions of constant pressure define standard enthalpy of combustion (ΔcHƟ) and standard enthalpy of formation (ΔfHƟ) 3.1.4.1 Enthalpy change Reactions can be endothermic or exothermic. Enthalpy change (ΔH) is the heat energy change measured under conditions of constant pressure. Standard enthalpy changes to standard conditions, i.e. 100 kPa and a stated temperature (e.g. ΔH298Ɵ). Students should be able to define: standard enthalpy of combustion (ΔcH Ɵ); standard enthalpy of formation (ΔfH Ɵ). 7.1 7.2 7.3 7.4 2 Calorimetry ideas of finding the energy of burning fuels. Use calorimetry to investigate and evaluate the enthalpy of combustion of methanol. Describe how calorimetry can be used to measure the ΔcHƟ of methanol use q = mcΔT to calculate the molar enthalpy change for a reaction using data collected by calorimetry Evaluate calorimetric methods and make suggestions for improvement 3.1.4.2 Calorimetry The heat change, q, in a reaction is given by the equation q = mcΔT. Where m is the mass of the substance that has a temperature change ΔT 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 7.5 Required practical 2: Measurement of an enthalpy change. 3 Enthalpy of solution out experiments to measure the enthalpy of solution of a range of different solids describe dissolution as an endothermic or exothermic process explain how to and carry out experiments to determine enthalpy change of solution use q = mcΔT to calculate the enthalpy change of solution 7.5 Required practical activity 2: Measurement of an enthalpy change. Student Book 7.5 page 152
of the specific heat capacity, c, of a substance. AT a and k PS 2.4, 3.1, 3.2, 3.3 and 4.1: Students could be asked to find ΔH 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. 4 Enthalpy of neutralisation Calculate expected temperature changes from given neutralisation reactions and compare predictions to empirical values. Compare temperature changes obtained in different reaction vessels. explain the term enthalpy of neutralisation Explain how to measure the enthalpy of neutralisation Predict temperature changes from the equation q = mcΔT 7.5 5 Enthalpy of combustion Introduce the bomb calorimeter and compare values for enthalpy of combustion obtained using simple calorimetry with a bomb calorimeter. Explain the differences in findings. define calculate standard enthalpies of combustion using the equation q = mcΔT describe a bomb calorimeter explain how a bomb calorimeter is different to simpler calorimeters that are used to determine enthalpies of combustion 7.5 6 Hess’s law Introduce Hess’s law and recap the definitions of enthalpy of formation, Explain Hess’s law Identify enthalpies of combustion and formation Hess’s cycles 3.1.4.3 Applications of Hess’s Law Hess’s Law. 7.6
combustion and reaction. Use Hess’s law to calculate values of any one of these enthalpies use Hess’s Law to perform calculations, including calculation of enthalpy changes for reactions from enthalpies of combustion or from enthalpies of formation 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. AT a and k PS 2.4, 3.2 and 4.1: Students could be asked to find ΔH 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 7 Hydration of a metal salt Determine the enthalpy of solution for anhydrous and hydrated magnesium sulfate experimentally and use these values and Hess’s Law to calculate the enthalpy of reaction to form the hydrated salt. collect experimental data to calculate enthalpy changes use Hess’s law to calculate enthalpy changes for reactions that cannot be determined directly Explain how Hess’s law can be used to calculate enthalpy changes for reactions 7.6 8 A Thermal decomposition reaction Work out how to measure the enthalpy change for the thermal decomposition of sodium hydrogen carbonate using enthalpy of neutralisation. Plan and carry out the investigation. collect experimental data to calculate enthalpy changes work out how Hess’s law to calculate enthalpy changes for the decomposition of sodium hydrogen carbonate explain and plan an investigation to use Hess’s law to calculate enthalpy changes for this reaction 7.6 9 Bond enthalpies Recap energy level diagrams and ideas of bond making and bond define the term mean bond enthalpy 3.1.4.4 Bond enthalpies Mean bond enthalpy. MS 1.2: Students understand that bond enthalpies are mean values across a range of 7.7
breaking. Use data on bond enthalpies to calculate enthalpies of reaction, combustion and formation. use mean bond enthalpies to calculate an approximate value of ΔH for reactions in the gaseous phase explain why values from mean bond enthalpy calculations differ from those determined using Hess’s Law Students should be able to: define the term mean bond enthalpy; use mean bond enthalpies to calculate an approximate value of ΔH for reactions in the gaseous phase; explain why values from mean bond enthalpy calculations differ from those determined using Hess’s Law. compounds containing that bond. One hour lessons Learning Outcomes Specification Content Skills Covered Student Book Section Required Practicals CHAPTER 8 – Kinetics (10 hours) 1 Rate of reaction Recap of what is meant by the term rate of reaction and how it is calculated. Explain how different factors affect the rate. Recall the definition of rate of reaction Describe how different factors affect the rate of reaction Explain how reactions take place in terms of collisions, and how various factors increase the rate. 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. 8.1 2 Collision theory Introduces the idea of orientation and activation. Uses these ideas to explain why most collision do not lead to a reaction. describe how moving particles store energy (translational, vibrational and rotational) define the term activation energy explain why most collisions do not lead to a reaction 8.2 8.3
Questions and graphs are used to determine the effect of concentration and pressure on the rate of reaction and how this effect can be measured. Describe how concentration and pressure affect the rate of reaction Make predictions how concentration and pressure affect the rate of a reaction on a graph Use collision theory to explain the effect of concentration and pressure on the rate of reaction 7 Continuous monitoring method Data for different reactions are given from which graphs are to be plotted. Rate of reactions are calculated from these continuous monitoring methods. Describe what is meant by continuous monitoring Plot and range of graphs including concentration of product formed and concentration of reactant lost for different reactions Explain how continuous monitoring may be used to investigate the rate of reaction 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. 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. 8.6 Required practical activity 3: Investigation of how the rate of a reaction changes with temperature. 8 Investigating concentration and reaction rate Investigate the reaction between calcium carbonate and hydrochloric acid at different concentrations to calculate the rate of reaction. Describe how to carry out an investigation using continuous monitoring to find the rate of reaction, using different concentration of acid Collect accurate and reliable data, plot graphs of the data and use to calculate the rate of reaction Peer review results and account for anomalies 8.6 9 Investigating catalysis Investigate the effect of different substances to Explain the term catalyst 3.1.5.5 Catalysts 8.8
catalyse the reactions of the decomposition of hydrogen peroxide and the reaction between sodium thiosulfate and iron (III) nitrate. Recognise that catalysts are specific to particular reactions Describe the effect of different catalysts on the rate of reaction 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 of lower activation energy. Students should be able to use a distribution to help explain how a catalyst increases the rate of a reaction involving a gas. 10 Catalysts and reaction rate Introduce the idea that some catalysts do take part in a reactions and investigate an example with the reaction between potassium sodium tartrate and hydrogen peroxide with cobalt (II) chloride as a catalyst. Relate catalysis to the Maxwell distribution. Recognise that some catalysts actually take part in a chemical reaction to form an intermediate step Make careful observations from investigations use a Maxwell–Boltzmann distribution to help explain how a catalyst increases the rate of a reaction involving a gas 8.8 One hour lessons Learning Outcomes Specification Content Skills Covered Student Book Section Required Practicals CHAPTER 9 – Equilibria (7 hours) 1 Dynamic Equilibrium Recap ideas about reversible reactions. Use the effect of heat on iodine crystals, the dissociation of Give examples of different reversible changes Explain how reversible reactions and dynamic equilibrium are linked Apply ideas of dynamic equilibrium to practical investigations 3.1.6.1 Chemical equilibria and Le Chatelier’s principle In a reversible reaction at equilibrium: forward and reverse reactions proceed 9.1
ammonium chloride, carbon dioxide in water and chromate (VI) and dichromate (VI) and to investigate dynamic equilibrium. at equal rates; the concentrations of reactants and products remain constant. 2 Investigating dynamic equilibrium Investigate what happens at equilibrium through the changes between iodine in cyclohexane and potassium iodide Use observations from reactions in equilibrium to make inferences explain that in a reversible reaction at equilibrium forward and reverse reactions proceed at equal rates explain that in a reversible reaction at equilibrium the concentrations of reactants and products remain constant 9.1 3 The equilibrium constant Kc Introduce what is meant by a homogenous and heterogeneous system. Introduce the expression for Kc and use it to perform calculations 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 9.2 4 Determining Kc for the reaction of ethanol and ethanoic acid Use the reaction between ethanol and ethanoic acid to determine the value of Kc describe how to carry out an investigation to determine Kc Determine the value of Kc from data obtained from reliable and accurate observations. compare the value to accepted values and account for differences 9.2 5 Factors affecting equilibrium 9.3
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 systems The equilibrium constant Kc is deduced from the equation for a reversible reaction. The concentration, in mol of a species X involved MS 0.3: Students estimate the effect of changing experimental parameters on a measurable value, e.g. how the value of Kc would change with temperature, given different specified conditions. MS 1.1: 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. MS 2.2 and 2.3: Students calculate: the concentration of a reagent at equilibrium; 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 7 Equilibria and the chemical industry Recap ideas about the Haber Process and introduce the manufacture of ethyl butanoate and nitric acid as industrial reactions in dynamic equilibrium. Analyse these reactions to determine the optimum conditions for product yield and compare to the actual reaction conditions Describe industrial reactions which are in dynamic equilibrium Use Le Châtelier’s principle to predict what the optimum conditions for the reactions are likely to be Explain why compromises in temperature and pressure need to be made. 9.3
1 Oxidation and reduction Recap ideas about reduction and oxidation, using the electrolysis of sodium chloride solution. Use the reaction between sodium and chlorine to investigate redox reactions, writing half and full equations. Describe oxidation and reduction processes in electrolysis and other reactions in terms of electron transfer Explain what happens in a redox reaction like the reaction between sodium and chlorine Use simple half equations to write full equations 3.1.7 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 for assigning oxidation states. Students should be able to: work out the oxidation of an element in a compound or ion from the formula; write half-equations identifying the oxidation and reduction processes in redox reactions combine to give an overall redox equation. 10.1 2 Oxidation states Introduce the term oxidation state and how to determine what they are for different elements in a compound. Use reactions of manganese to investigate the different oxidation states it exists in. explain the term oxidation state work out the oxidation state of an element in a compound or ion from the formula make accurate observations to predict the oxidation state of manganese in different reactions 10.2 3 Redox reactions Recap identify the oxidation states of an element as a reactant and product in a reaction Identify “atoms” are oxidised and reduced during chemical changes Explain which are the reducing and oxidizing agents in given reactions 10.3 4 equations Use metal displacement reactions to further investigate redox reactions. Practice write complex half-equations identifying the oxidation and reduction processes in redox reactions combine half-equations to give an overall redox equation 10.3
excess concentrated ammonia solution. 4 Required practical 4 Carry out simple tests to identify cations – group 2 M2+; ammonium NH4+ and anions group7 (17), X- halide ions; hydroxide OH-; carbonate CO3 2-; sulfate SO4 2-. work safely in a laboratory use laboratory apparatus for a variety of experimental techniques, including qualitative tests for ions draw conclusions from observations made 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. Research opportunity: Students could investigate the treatment of drinking water with chlorine. Students could investigate the addition of sodium fluoride to water supplies. 12.3 Required practical activity 4: Carry out simple reactions to identify: • cations – Group 2, NH4 + • anions – 7 (halide ions), OH– , CO3 2–, SO4 2– 5 Chlorine and water supplies Introduce the term disproportionation and it to the reaction between chlorine and water. Use data to explore the use of chlorine in water supplies describe how chlorine reacts with water to form chloride and chlorate ions and chloride ions and oxygen gas Use ideas about oxidation and reduction to explain how chlorine reacts with water explain the benefits and disadvantages of adding chlorine to water supplies 12.4 6 Chlorine in bleach Use a demonstration to show the bleaching action of sodium chlorate (I) food dyes. Use a class practice to investigate the describe the chemical reaction between chlorine and cold dilute sodium hydroxide solution carry out an experiment and process the results to estimate the concentration of sodium chlorate(I) in different commercial bleaches 12.4 tube
3 Hydrolysis of halogenoalkanes Compare the rate of hydrolysis of and using a practical investigation explain the term hydrolysis carry out test tube hydrolysis reactions of halogenoalkanes and describe methods to compare the rate of reaction explain why the carbon-halogen bond enthalpy influences the rate of reaction construct 13.2 4 Mechanisms for the hydrolysis of primary and tertiary halogenoalkanes Draw examples of primary, secondary and tertiary halogenoalkanes and explain the differences between them. Explain the reaction mechanisms of primary and tertiary halogenoalkanes with water using curly arrows identify and draw displayed formulas for primary, secondary and tertiary halogenoalkanes explain why there are differences in the reaction mechanisms of primary, secondary and tertiary halogenoalkanes explain the mechanisms for the hydrolysis of primary and tertiary halogenoalkanes 13.2 5 Preparation of a halogenoalkane Follow instructions, prepare a sample of follow instructions and work safely in the lab explain how to prepare a sample of a halogenoalkane explain how to purify the product using a separating funnel and distillation 13.2 6 Elimination reactions of halogenoalkanes outline the mechanism of elimination reactions of halogenoalkanes 3.3.3.2 Elimination 13.2
Explain the different between substitution and elimination reactions and why elimination reactions occur instead of substitution reactions. Explain the mechanism of the reaction using curly arrows. explain how hydroxide ions act as a in elimination reactions predict the products of competing substitution and elimination reactions between hydroxide ions and halogenoalkanes The concurrent substitution and elimination reactions of a halogenoalkane (e.g. 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. 7 Ozone depletion Describe the role of ozone and evidence for its depletion. Use reactions to explain how it is being depleted. Explain some of the measures used to restore its concentrations describe the role of naturally formed ozone in the upper atmosphere use equations to explain how chlorine atoms catalyse the decomposition of ozone explain the role of the scientific community in providing evidence for legislation to ban the use of CFCs as solvents and refrigerants and find alternatives 3.3.3.3 Ozone depletion 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 formation of a hole in the ozone layer. Appreciate that results of research by chemists provided evidence for legislation to ban the use of CFCs as solvents and refrigerants. Chemists have now developed alternative chlorine-free compounds. Students should be able to: use equations, such as the following, to explain how chlorine atoms catalyse decomposition of ozone: Research opportunity: Students investigate the role of chemists in the introduction of legislation to ban the use of CFCs and in finding replacements. 13.3
Cl• + O3 → ClO• + O2 and ClO• + O3 → 2O2 + Cl• One hour lessons Learning Outcomes Specification Content Skills Covered Student Book Section Required Practicals CHAPTER 14 – Alkenes (7 1 Alkenes Recap the main features of alkenes, including their ability to form geometric and stereoisomers. Explain the differences between sigma and pi bonds. Describe the essential features of alkenes as a homologous series Recall the types of isomers present in alkenes Explain the nature of the bonding in alkenes including pi and sigma bonds 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. 14.1 2 Electrophilic addition reaction of alkenes Introduce the term electrophile. Using displayer formulae, explore the electrophilic addition reactions of different electrophiles with alkenes. name the features of an electrophile and give some examples explain why alkenes are susceptible to attack by electrophiles describe the mechanisms of the electrophilic addition reactions of alkenes with HBr, H2SO4 and Br2 using curly arrows 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 and minor products in addition reactions of unsymmetrical alkenes. Students should be able to: outline the mechanisms for reactions; explain the formation of major and minor products by reference to the relative 14.2 3 Hydration of alkenes Carry out a practical to investigate the electrophilic addition reaction between Use equipment safely and follow a procedure Describe the procedure to produce from hex-1 ene 14.2
stabilities of primary, secondary and tertiary carbocation intermediates. and sulphuric acid with water. Explain the reaction mechanism for the addition reaction, explaining why hexan-2-ol was made instead of 4 The addition of reactions of unsymmetrical alkenes Use molecular models to explore the possible products from asymmetrical alkenes. Explain the stability of primary, secondary and tertiary carbocations, and use these ideas to predict the products formed from addition reactions of unsymmetrical alkenes identify symmetrical and unsymmetrical alkenes recognise primary, secondary and tertiary carbocations predict the major and minor products in addition reactions of unsymmetrical alkenes 14.2 5 Determining the degree of unsaturation in oils describe the test for unsaturation describe a procedure to test the unsaturation of different oils bromine water explain the reaction mechanism of the reaction between bromine water and less familiar alkenes using curly arrows 14.2 6 Polymers Recap the terms monomer, polymer, polymerisation and repeat unit, and the use of the repeating unit to show the structure of the polymer from the monomer. Research describe the term polymer and use the IUPAC rules for naming polymers explain how polymers are formed from alkenes and substituted alkenes, drawing the repeating unit of a polymer from the monomer and from the polymer chain Understand the nature of the intermolecular forces in polyalkenes. 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. 14.3
different common addition polymers and explain their properties and nature of their intermolecular forces 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 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. 7 Making polymers Demonstrate the preparation of poly (phenylethene) and describe the process. Explain the chemical changes taking place and the role of pieces of equipment Identify the monomer for making poly (phenylethene) and write the equation for the polymerisation reaction. describe how to make poly(phenylethene) also known as polystyrene in the laboratory Compare the laboratory preparation with industrial methods 14.3 One hour lessons Learning Outcomes Specification Content Skills Covered Student Book Section Required Practicals CHAPTER 15 – Alcohols (8
1 Structure and physical properties alcohols Recap the structure of alcohols and introduce primary, secondary and tertiary alcohols. Use data to analyse and compare the boiling points of alkanes, alkenes and alcohols, explain this in terms of hydrogen bonding use IUPAC rules to name alcohols and identify primary, secondary and tertiary alcohols interpret data on boiling points of the homologous series studied so far explain how hydrogen form in alcohols and the impact this has on their properties 3.3.5.1 Ethanol 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. 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. AT a, d and k PS 1.2: Students produce ethanol by fermentation, followed by purification by fractional distillation. 15.1 2 Industrial manufacture of ethanol Introduce the commercial production of ethanol from the hydration of ethane. Discuss the reasons for the reaction conditions and outline the mechanism for the reaction describe how alcohols are produced industrially by the hydration of ethene use le Chatelier’s principle to explain the conditions used for its manufacture outline the mechanism for the reaction between ethene and steam 15.2 3 Ethanol by fermentation out a practical to produce ethanol by fermentation. Discuss the features and uses of the apparatus and reagents used. Discuss the use of ethanol as a biofuel and the ethical implications arising from its use. explain the term biofuel and the ethical issues associated to their use explain the conditions and processes required for the production of ethanol by fermentation use equations to support the statement that ethanol produced by fermentation is a carbon neutral fuel and explain why this statement is not valid 15.2
7 Elimination reactions Use molecular models to demonstrate elimination reactions. Use curly arrow to explain the mechanism of an elimination reaction in detail. Describe how alkenes can be formed from alcohols in acid catalysed reactions Explain how alkenes produced from alcohols can be used to produce addition polymers without the use of monomers derived from crude oil Outline the mechanism for the elimination of water from alcohols 3.3.5.3 Elimination Alkenes can be formed from alcohols by acid-catalysed elimination reactions (mechanism not required). Alkenes produced by this method can be used to produce polymers without using monomers derived from crude oil. 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. 15.4 8 Preparation and purification of cyclohexene from cyclohexanol Class carry out required practical to prepare and purify cyclohexene from cyclohexanol Calculate the atom economy and percentage yield for the reaction. Describe how to carry out the preparation of cyclohexene from cyclohexanol Explain how to purify the product using a separating funnel and distillation Calculate a percentage yield 15.3 15.4 Required practical 5: Distillation of a product from a reaction. One hour lessons Learning Outcomes Specification Content Skills Covered Student Book Section Required Practicals CHAPTER 16 – Organic analysis (6 hours) 1 identifying functional groups Carry out a variety of different tests to recall the reactions of functional groups in this specification describe and carry out test reactions to identify alcohol, 3.3.6.1 Identification of functional groups by test-tube reactions 16.1
identify functional groups studied in this specification aldehydes, alkene, halogenoalkane and carboxylic acid on known samples, either by using given methods, or by researching their own methods record their observations in a table and make suitable deductions The reactions of functional groups listed in the specification. Students should be able to identify the functional groups reactions in the specification. in the specification to distinguish alcohols, aldehydes, alkenes and carboxylic acids. 2 Required practical activity Students carry out their own tests to determine the identity of unknown functional groups as part of required practical 6. use suitable chemical tests to identify the functional groups present in organic compounds Decide which tests to carry out to determine the identity of the following functional groups:- alcohol, aldehydes, alkene and carboxylic acid explain the chemistry of each test Required practical 6: for alcohol, aldehyde, alkene and carboxylic acid. 3 Mass spectrometry Recap of the time of flight spectrometer. Interpret a range of mass spectra. Compare the differences between low and high resolution mass spectrometers. explain how mass spectrometry can be used to determine the molecular formula of a compound use precise atomic masses to calculate precise molecular masses use precise molecular masses to determine the molecular formula of a compound 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. 16.2 4 Infrared spectroscopy Introduce infrared radiation and explain its effect on the bonds in organic molecules. Explain how the spectrometer works explain the link between wavelength and wavenumber know that in a molecule absorb infrared radiation at characteristic wavenumbers explain the principles of infrared spectrometry 3.3.6.3 Infrared spectroscopy 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 16.3 5 Interpreting infrared spectra use infrared spectra and the Chemistry Data sheet to identify 16.3