3.2.2 Group 2, the alkaline earth metals 1.4 Energetics 4 3.1.4 Enthalpy change, calorimetry, Hess’s law, bond enthalpies 4.0 Required practical 2 How fast and how far? Rates and equilibrium 2.6 3.1.5 Kinetics 1.6 Required practical 3 3.1.6 Equilibrium 1.0 Teacher 2 (15.4 Themes Topic Number of Structure and bonding 3 weeks 3.1.3.Types of bonding and physical properties 1.0 3.1.3. Shapes of molecules and intermolecular forces 1.8 3.2.1.2 Physical properties of Period 3 elements 0.2 Hydrocarbons 5 weeks 3.3.1.1 Nomenclature 3.3.2. Fractional distillation, cracking and Alkanes 2.4 3.1.2.3 The gas equation 0.4 3.3.1.3 E/Z Isomerism 3.3.4.Alkenes and addition polymers 2.2 Required practical 5 reactions and Group 7 2.4 weeks 3.1.7 Redox reactions 0.8 3.2.3. Halogens 1.6 Required practical 4 Further organic chemistry 3.3.3 Halogenoalkanes 1.4 3.3.5 Alcohols 1.4
Relative atomic mass and relative molecular mass. Relative atomic mass and relative molecular mass in of 12C. Relative formula mass. Calculate relative atomic mass from isotopic abundance, limited to mononuclear ions. Look at mass spectra of compounds to determine the relative formula 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 (IEs). 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. 1.2 weeks Give electron configurations in terms of shells and sub-shells for atoms and ions up to Z=36 Define first ionisation energy. Write equations for first and successive ionisation energies. Explain how first and successive ionisation energies in Period 3 and in Group 2 give evidence for electron configuration in sub-shells and in shells. Students: Write electron configurations for atoms and ions up to Z=36 in terms of shells and sub-shells. Find out about the historical development of the Periodic Table. Plot graphs of both first IE and successive IEs and use them to identify and explain the trends. Use successive IE data to determine group an element is in. January 2012 Unit 1 Q5a and 5b June 2013 Unit 1 Q6b,6c and 6d January 2010 Unit 1 Q1c, 1d and 1e January 2009 Unit 1 Q1a and 1b of the Periodic Table Molymod models to show shapes of orbitals.
3.1.2.5 Balanced equations. Formulae and equations. 0.6 weeks Write balanced equations for reactions studied. Balance equations for unfamiliar reactions when reactants and products are specified. Students learn the formulae and charges of common ions and use them to write formulae of compounds and balanced equations. Students use their observations from simple test-tube reactions to identify the products and write balanced equations. Precipitation, acid-base and thermal decomposition reactions as examples of writing formulae and balancing equations. 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. 0.2 weeks Carry out calculations: using the Avogadro constant using mass of substance, Mr, and amount in moles(concentrations and gas volume calculations are covered later) Students practice carrying out appropriate calculations. Practical activity: the number of molecules in a mouthful of water. 3.1.2.4 Empirical and molecular formula. Definition of empirical formula and molecular formula, and the 0.6 weeks Calculate empirical formula from data giving composition by mass or percentage by mass. Calculate molecular formula from the empirical Students carry out appropriate calculations. Practical activity: Empirical formula of magnesium June 2010 Unit 1 Q4a June 2009 Unit 1 Q2c
relationship between them. formula and relative molecular mass. 3.1.2.5 Balanced equations and associated calculations. Reacting masses, percentage and atom economy. 0.8 weeks Use balanced equations to calculate: masses percentage yields atom economies. Students carry out appropriate calculations. Practical activities: Water of crystallisation of hydrated MgSO4 by heating to constant mass. Percentage conversion of magnesium carbonate to magnesium June 2013 Unit 1 Q7a, 7b and 7c January 2013 Unit 1 Q5b June 2012 Unit 1 Q5d and 5ee January 2011 Unit 1 Q3d Atom economy and percentage yield Structure and bonding (3
Learning objective Time taken Learning outcome Learning activity with opportunity to develop skills Assessment opportunities Resources 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. 0.2 weeks Predict the charge on a simple ion using the position of the element in the Periodic Table. Construct formulas for ionic compounds. Students draw dot and cross diagrams for simple ionic compounds. 3.1.3.2 Covalent bonding. Covalent bonding (single and multiple), co-ordinate bonding. 0.4 weeks Describe the nature of a covalent bond and use a line to represent it. Describe a co-ordinate bond and use an arrow to represent it. Students draw dot and cross diagrams to represent covalent molecules, including adducts and ions (eg BF3.NH3 and NH4+) which involve the use of co-ordinate bonds. 3.1.3.3 Metallic bonding. Metallic bonding. 0.2 weeks Describe the nature of metallic bonding. Spot the bonding 3.1.3.4 Bonding and physical properties. 0.6 weeks Relate the melting point and conductivity of materials to the type of Practical activity: relationship between physical properties and structure and bonding. June 2013 Unit 1 Q3 June 2011 Unit 1 Q4 June 2010 Unit 1 Q7 June 2006 Unit 1 Q2 Structure and bonding
do not have a permanent dipole. 3.1.3.7 Forces between molecules and their effect on melting point and boiling point. Permanent forces. Induced dipole–dipole (van der Waals) forces. Hydrogen bonding including low density of ice and anomalous boiling points of compounds. 0.6 weeks Explain the existence of forces between familiar and unfamiliar molecules. Explain how melting and boiling points are influenced by these intermolecular forces. Students predict and explain the relative boiling points of a range of covalent substances. Students interpret the graph of boiling point against Mr of covalent hydrides in groups 4 to 7. Practical activity: making slime. Practical demonstration: density of ice. June 2013 Unit 1 Q4 January 2012 Unit 1 Q1 June 2011 Unit 1 Q3d January 2011 Unit 1 Q1 January 2010 Unit 1 Q3 Hydrogen bonding activity Practical: Slime Practical: Density of ice 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 0.2 weeks Explain the trends in atomic radius and first ionisation energy. Explain the melting point of the elements in terms of their structure and bonding. Students plot data on graphs for atomic radius, first ionisation energy and melting point and explain those trends. Students give explanations for these trends. January 2011 Unit 1 Q5 January 2009 Unit 1 Q4
The reasons for these trends in terms of the structure of and bonding in the elements. Opportunities for skills development Acids, bases and Group 2 (4 weeks) Learning objective Time taken Learning outcome Learning activity with opportunity to develop skills Assessment opportunities Resources 3.1.2.5 Balanced equations and associated calculations. The concentration of a substance in solution, measured in mol dm–3 and calculations concentration, volume and amount of substance in a solution. 2.6 weeks Use balanced equations to calculate concentrations and volumes for reactions in solutions. Construct and/or balance equations using ratios. Select appropriate titration data (ie identify outliers) in order to calculate mean titres. Determine uncertainty when two burette readings are used to calculate a value. Practical activities: Concentration of ethanoic acid in vinegar. Mr of succinic acid Percentage calcium carbonate in indigestion tablets. Mass of aspirin in an aspirin tablet. Students perform a range of different types of calculations involving reacting masses and volume of solutions. June 2004 Unit 1 Q2 January 2004 Unit 1 Q3 January 2002 Unit 1 Q7 January 2009 Unit 1 Q3b January 2011 Unit 1 Q3 June 2010 Unit 1 Q3 Indigestion tablets
Required practical 1 Make up a volumetric solution and carry out a simple acid–base titration. 3.2.2 2, the alkaline earth metals. The trends in atomic radius, first ionisation energy and melting point of the elements Mg–Ba 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. The relative solubilities of the sulfates. Uses of some Group 2 compounds. 1.4 Explain the trends in atomic radius and first ionisation energy. Explain the melting point of the elements in terms of their structure and bonding. Describe the reactions of Mg– Ba with water. Describe the relative solubilities of the hydroxides and sulfates. Recall the use of magnesium in the extraction of Ti Give the uses of Ca(OH)2. Mg(OH)2, CaO and BaSO4 Explain why BaCl2 solution is used to test for sulfate ions and why it is acidified. Students plot graphs of atomic radius and first ionisation energy, and explain the trends. Practical activities: Reactions of Mg–Ba with water and Mg with steam. Testing solubilities of hydroxides and sulfates by precipitation reactions Test a variety of solutions with NaOH/H2SO4 to identify unknown compounds. Students write full and ionic equations for these reactions and for a variety of other precipitation reactions. June 2012 Unit 2 Q5 June 2006 Unit 1 Q5a January 2005 Unit 1 Q5b January 2012 Unit 2 Q7 Limestone (RSC) 2 reactions: making a mind map
Fractional distillation of crude oil. Crude oil is a mixture of saturated hydrocarbons which can be separated by fractional distillation. 3.3.2.2 Cracking. Thermal and catalytic cracking: why cracking is useful and the type of products formed. hydrocarbons in oil are separated on the basis of their different boiling points. Recall the conditions for thermal and catalytic cracking. Understand that cracking involves breaking C-C bonds, and explain why it is economically important. Students write balanced equations for a variety of cracking reactions. June 2004 Unit 3 Q5 Practical: fractional distillation of crude oil RSC Videos and animations on fractional distillation of crude oil Animations of fractional distillation 3.3.2.3 Combustion of alkanes. Alkanes as Complete and incomplete combustion. Use of catalytic converter in internal combustion engines to remove pollutants. 0.6 weeks Write equations for the complete and incomplete combustion of alkanes. Explain how a number of pollutants including NOx, CO, C and unburned hydrocarbons are formed in the internal combustion engine and how their emissions can be reduced. Why SO2 may be formed when fuels are burned and Practical demonstration: methane bubbles. Students write balanced equations for a variety of combustion reactions. June 2010 Unit 1 Q4 June 2010 Unit 1 Q5 January 2004 Unit 3 Q2 January 2009 Unit 1 Q6a, 6b,6c and 6d Methane bubbles about catalytic converters
Unsaturated hydrocarbons with an electron rich double bond. 3.3.1.3 Isomerism. E–Z isomerism is a form of stereo-isomerism and occurs as a result of restricted rotation about the planar double bond. Cahn–Ingold–Prelog (CIP) priority Draw the structural formulas of E and Z isomers. Apply the CIP priority rules to E and Z isomers. Students draw and name isomers of and C6H12, and identify which exhibit E–Z isomerism. Required practical 5 Distillation of a product from a reaction. 3.3.4.2 Addition reactions of alkenes. Electrophilic addition reactions of alkenes with HBr, H2SO4 and Br2 0.8 weeks Outline the mechanisms for these reactions by drawing the structures of the species involved and curly arrows to represent the movement of electron pairs. Understand that the formation of a covalent bond is shown by a curly Students write balanced equations for electrophilic addition reactions. Students write balanced mechanism steps for different alkenes, and identify (and explain the existence of) major and minor products where appropriate. June 2012 Unit 2 Q7b June 2010 Unit 2 Q6a January 2003 Unit 3 Q2a January 2005 Reaction mechanism 1:
Energetics (4 weeks) Learning objective Time taken Learning outcome Learning activity with opportunity to develop skills Assessment opportunities Resources 3.1.4.1 Enthalpy change. Endothermic and exothermic. Define enthalpy change (ΔH) and standard conditions. 0.8 Define standard enthalpy of combustion (ΔcHƟ). Define standard enthalpy of formation (ΔfHƟ). Practical activity: enthalpy of combustion of alcohols. Demonstration of a selection of exothermic and endothermic reactions. Students write equations for formation and combustion reactions. Students sketch and label enthalpy profile diagrams. June 2006 Unit 2 Q1a Spectacular demonstrations: An endothermic reaction: An exothermic reaction: 3.1.4.2 Calorimetry q = mcΔT calculations and associated experiments. 1.4 weeks Use q = mcΔT to calculate the molar enthalpy change for a reaction. Use this equation in related calculations. Practical activity: Measuring enthalpy changes of various reactions using the cooling curve method (eg neutralisation of NaOH with HCl, displacement reaction between Zn and CuSO4, dissolving KCl in water). Students perform a variety of calculations. June 2009 Unit 2 Q3 June 2002 Unit 2 Q2 June 2006 Unit 2 Q1d January 2011 Unit 2 Q9b Enthalpy practical: 3.1.4.3 Hess’s Law. 1.2 weeks Use Hess’s Law to perform calculations, including calculation of enthalpy Practical activity: Hess’s Law to determine ΔH for the thermal decomposition of NaHCO3 or the June 2011 Unit 2 Q2
changes for reactions from enthalpies of combustion or from enthalpies of formation. hydration of MgSO4 or formation of CaCO3 Students practice a range of calculations using Hess’s Law. June 2002 Unit 2 Q1 June 2009 Unit 2 Q2a January 2013 Unit 2 Q3a and 4 Required practical 2 Measurement of an enthalpy change. 3.1.4.4 Bond enthalpies. Understand the term mean bond enthalpy. Use mean bond enthalpies to calculate approximate values for ΔH for reactions. Understand why most bond enthalpies are mean values. 0.6 weeks Recall the meaning of the term mean bond enthalpy. Calculate enthalpy changes using mean bond enthalpies. Understand why most bond enthalpies are mean values. Students practice a range of calculations using bond enthalpies (students could compare the calculation of the enthalpy of combustion of a fuel using bond enthalpies with both the theoretical and experimental values). January 2013 Unit 2 Q6 January 2006 Unit 2 Q1 June 2005 Unit 2 Q1 January 2003 Unit 2 Q2 reactions and Group 7 (2.4 weeks)
Learning objective Time Learning outcome Learning activity with opportunity to develop skills Assessment opportunities Resources 3.1.7 Oxidation, reduction and redox equations. Oxidation and reduction in terms of electron transfer. Rules for assigning oxidation states. Writing and full equations. 0.8 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. Students work out oxidation states of elements in a of compounds. Students write half-equations for different oxidation and reduction reactions, and then combine them. June 2013 Unit 2 Q4a January 2012 Unit 2 Q5a and 5b June 2011 Unit 2 Q5a January 2005 Unit 2 Q2 January 2002 Unit 2 Q4 Thermit reaction demonstration: 3.2.3.1 Trends in the properties of the Halogens. Trends in electronegativity and boiling point down Group 7 Trends in oxidising power of halogens and reducing power of halide ions. 1.2 weeks Explain the trend in electronegativity and the boiling point of the elements in terms of their structure and bonding. Explain why silver nitrate is acidified in the halide test, and why ammonia is used to distinguish between the halides. Recall observations and write equations for the Students plot data on graphs for electronegativity and boiling point and explain those trends. Write full and ionic equations for the reactions studied. Write half equations for the redox reactions studied. Practical activities: Test-tube reactions of solutions of the halogens (Cl2, Br2, I2) with solutions containing their halide ions (eg KCl, KBr, KI). January 2002 Unit 2 Q8 January 2013 Unit 2 Q9 June 2002 Unit 2 Q4 June 2012 Unit 2 Q9 January 2011
Learning objective Time taken Learning outcome Learning activity with opportunity to develop skills Assessment opportunities Resources 3.1.5.1 Collision theory. 3.1.5.2 Maxwell–Boltzmann distribution. Define Activation energy. Maxwell–Boltzmann distribution of molecular energies in 0.4 weeks Define the term activation energy. Explain why most collisions do not lead to a reaction. Draw and interpret distribution curves for different temperatures. Practical demonstration: A selection of demonstrations to illustrate some fast and slow reactions. June 2013 Unit 2 Q3 January 2012 Unit 2 Q3 June 2006 Unit 2 Q2 January 2002 Unit 2 Q7 Spectacular demonstrations 3.1.5.3 Effect of temperature on reaction rate. The qualitative effect of temperature changes on the rate of reaction. 0.4 weeks Use the distribution to explain why a temperature increase can lead to a large increase in rate. June 2006 Unit 2 Q2 January 2004 Unit 2 Q2 January 2012 Unit 2 Q3 Effect of temperature on rate: Required practical 3 Investigation of how the rate of a reaction changes with temperature. 3.1.5.4 Effect of concentration and pressure. 0.4 weeks Explain how a change in concentration or a change in pressure influences the rate of a reaction. Practical activity: The kinetics of the reaction between CaCO3 and HCl by continuous monitoring.
The qualitative effect of changes in concentration and pressure on collision frequency. Students plot graphs of results from other continuous monitoring experiments, and interpret the data. 3.1.5.5 Catalysts. Define catalyst. Catalysts work by providing an alternative reaction route of lower activation energy. 0.4 weeks Use a Maxwell–Boltzmann distribution to help explain how a catalyst increases the rate of a reaction involving a gas. Practical activity: Reaction between hydrogen peroxide and tartrate ions catalysed by Co(II) ions. June 2012 Unit 2 Q1 June 2011 Unit 2 Q1 January 2003 Unit 2 Q3 January 2011 Unit 2 Q2b Hydrogen peroxide and tartrate ions catalysed by Co(II) ions. 3.1.6.1 Chemical equilibria and Le Chatelier’s principle. Know that at equilibrium the forward and reverse reactions proceed at equal rates and the concentrations of reactants and products remain constant. State Le Chatelier’s principle. 0.6 weeks 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. Practical activity: The effect of concentration and temperature changes on the equilibrium between [Co(H2O)6] 2+ and CoCl42- Students predict and explain the effect of changes in temperature, pressure and concentration on the position of an equilibrium. Students explain the use of compromise conditions in a of industrial processes June 2013 Unit 2 Q10a June 2013 Unit 2 Q1a January 2013 Unit 2 Q2 January 2012 Unit 2 Q2 Ammonia production Equilibrium between CoCl42-and [Co(H2O)6 ]2+
3.1.6.2 Equilibrium constant Kc for homogeneous systems. The equilibrium constant Kc can be deduced from the equation for a reversible reaction. 0.4 weeks Construct an expression for Kc for a homogeneous system in equilibrium. Calculate a value for Kc from the equilibrium concentrations. Perform calculations involving Kc Predict the qualitative effects of changes of temperature on the of Kc Students practice working out the expression, units and value of Kc for a variety of different equilibrium reactions. Students practice calculations to determine equilibrium concentrations and use them to calculate Kc Practical activity: Determination of the value of Kc of an esterification reaction (eg ISA Q: CHM6T June 2011). June 2013 Unit 4 Q2 June 2012 Unit 4 Q1 January 2013 Unit 4 Q3 Further organic chemistry (5 weeks) Learning objective Time taken Learning outcome Learning activity with opportunity to develop skills Assessment opportunities Resources
3.3.3.1 Nucleophilic substitution of halogenoalkanes with OH–, CN– and NH3 0.6 weeks Outline the nucleophilic substitution mechanisms of reactions. Explain why the bond enthalpy influences the rate of reaction. Practical activities: Preparation of from of hydrolysis of halogenobutanes. Students practice writing balanced equations and mechanism steps for different substitution reactions. January 2011 Unit 2 Q8a and 8b June 2010 Unit 2 Q2 Importance of organic chemistry: Making medicines 3.3.3.2 Elimination. The concurrent substitution and elimination reactions of a halogenoalkane (eg with potassium hydroxide). 0.4 weeks Explain the role of the reagent as both nucleophile and base. Outline the mechanisms of these reactions Students practice writing balanced equations and mechanism steps for different elimination reactions. Students compare substitution and elimination with different halogenoalkanes. June 2013 Unit 2 Q5 January 2011 Unit 2 Q8 January 2010 Unit 2 Q7 June 2009 Unit 2 Q8 June 2002 Unit 3 Q6 3.3.3.3 depletion. 0.4 Understand the role of ozone in the atmosphere. Understand how chlorine free-radicals can be formed Students write equations and mechanisms for the formation of chlorine free-radicals and the destruction of ozone. January 2013 Unit 2 Q7 June 2011 Unit 2 Q7 RSC resource on CFCs and ozone:
The role of CFCs in depletion and the search for alternatives. in the atmosphere from compounds such as CFCs. Understand the mechanism for the depletion of ozone by chlorine free-radicals. Evaluate the role of chemists in the introduction of legislation to ban the use of CFCs and to find replacements. Students find out about the role of chemists in the introduction of legislation to ban the use of CFCs and in finding replacements. June 2009 Unit 2 Q11 3.3.5.1 Alcohol production. Hydration of alkenes to make alcohols. Compare hydration of ethene with fermentation of glucose in the production of ethanol. 0.4 weeks Explain the meaning of the term biofuel and discuss the environmental issues of biofuel usage. Write equations for ethanol produced by fermentation. Outline the mechanism for the formation of ethanol by the reaction of ethene with steam in the presence of an acid catalyst. Practical demonstrations: Fermentation of glucose followed by fractional distillation. Whoosh bottle. Students write balanced equations and mechanism steps for the dehydration of a variety of alcohols. Students compare hydration and fermentation processes. Students research the pros and cons of the use of biofuels. January 2005 Unit 3 Q5a, 5b and 5c January 2002 Unit 3 Q7 January 2012 Unit 2 Q9 Ethanol as a fuel: 3.3.5.2 Oxidation of alcohols. 0.8 weeks equations for these oxidation reactions. Practical activities: Test-tube reactions to illustrate January 2013 Unit 2 Q5 Giant silver mirror:
Understand how addition polymers can be made from alkenes made this way without using monomers derived from crude oil. 3.3.6.1 Identification of functional groups by test-tube reactions. 0.8 weeks Identify alcohols, aldehydes, alkenes and carboxylic acid functional groups using reactions in the specification. Interpret observations from these reactions. Practical activity: Revise tests for the appropriate functional groups. Students plan a series of tests to distinguish between unknown compounds. January 2010 Unit 2 Q6 June 2011 Unit 2 Q6 Required practical 6 Tests for alcohol, aldehyde, alkene and carboxylic acid. 3.3.6.2 Mass spectrometry to determine the molecular formula of a compound. 0.4 weeks Use precise atomic masses and the precise molecular mass to determine the molecular formula of a compound. Students use relevant data to determine the Mr from mass spectra. June 2012 Unit 2 Q3c January 2010 Unit 2 Q6e June 2011 Unit 2 Q8c January 2012 Unit 2 Q6
3.3.6.3 Infrared spectroscopy. Use infrared absorptions to identify functional groups. Know how the “fingerprint” region can be used. The role of infrared absorption by molecule in global warming. 0.6 Use infrared spectra and the Chemistry Data Sheet to identify particular bonds, and functional groups, and also to identify impurities. Students use relevant data to determine the bonds and functional groups from IR spectra. June 2012 Unit 2 Q8bii June 2011 Unit 2 Q6e January 2012 Unit 2 Q10 June 2009 Unit 2 Q9 January 2010 Unit 2 Q9b Spectra database Spectraschool: 3.3.1.2 Reaction mechanisms. The formation of a covalent bond is shown by a curly arrow that starts from a lone electron pair or from another covalent bond. The breaking of a covalent bond is shown by a curly arrow starting from the bond. 0.2 weeks Outline mechanisms by drawing the structures of the species involved and curly arrows to represent the movement of electron pairs. Students revise the five mechanisms: free radical substitution, electrophilic addition, nucleophilic substitution and elimination (halogenoalkane) and elimination (alcohol). January 2013 Unit 2 Q7 and 8 January 2011 Unit 2 Q7 and 8 RSC mechanisms resource: Mechanism Inspector:
3.3.1.3 Isomerism. Structural isomerism. Stereoisomerism. 0.2 weeks 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. Students revise the different types of isomerism using different examples from the specification. June 2013 Unit 2 Q8 June 2012 Unit 2 Q7 June 2011 Unit 2 Q3fi and Q6 June 2003 Unit 3 Q3 January 2003 Unit 3 Q1 June 2002 Unit 3 Q4 Nomenclature and isomerism in organic chemistry: Useful resources and websites Chemistry Demonstrations Class practicals and other resources Royal Society of Chemistry Nuffield Foundation
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