Teaching scheme Week 1 1. Revision of AS 2.1.3 2. Benzene s structure historical to modern 3. The evidence for and against the delocalised model 4. The extra stability of the benzene molecule and its reluctance to undergo addition reactions 5. The mononitration of benzene 6. The monohalogenation of benzene 7. The mechanisms for the two electrophilic substitution reactions in 5 and 6 above Explain the terms: arene and aromatic. Describe and explain the models used to depict the structure of benzene. Review the evidence for a delocalised model of benzene. Describe the delocalised model of benzene. Describe the electrophilic substitution of arenes with concentrated nitric acid. Describe the electrophilic substitution of arenes with a halogen in the presence of a halogen carrier. Outline the mechanism of electrophilic substitution in arenes. Outline the mechanism for the mononitration and monohalogenation of benzene. 1.1.1 5 Practical activity 1: The nitration of methyl benzoate to form methyl 3-nitrobenzoate 4.1.1 Arenes Structure of benzene Electrophilic substitution of arenes 1
Week 2 1. A comparison of the reaction of bromine with benzene and an alkene such as cyclohexene 2. The structures of phenol and other phenols 3. The acidic properties of phenol, i.e. with sodium and sodium hydroxide 4. The explanation of the acidic properties of phenol in terms of the delocalisation of oxygen s lone pairs 5. The reaction of phenol with bromine 6. The explanation of phenol s reactivity compared with benzene in terms of the delocalisation of oxygen s lone pairs 7. The uses of phenols Explain the relative resistance to bromination of benzene compared with alkenes. Describe the reactions of phenol with aqueous alkalis and with sodium to form salts. Discuss the role of phenol as an early antiseptic. Describe the reactions of phenol with bromine to form 2,4,6-tribromophenol. Explain the relative ease of bromination of phenol compared with benzene. State the uses of phenols. 1.1.6 8 4.1.1 Arenes Phenols 2
Week 3 1. The carbonyl group and the difference between aldehydes and ketones 2. The oxidation of primary alcohols to aldehydes and carboxylic acids 3. The oxidation of secondary alcohols to ketones 4. The oxidation of aldehydes to carboxylic acids 5. The reduction of aldehydes and ketones 6. The mechanism for the reaction of carbonyls with the H ion in NaBH 4 7. The use of 2,4-DNPH to detect and identify carbonyl compounds 8. The use of Tollens reagent to distinguish between aldehydes and ketones Recognise and name aldehydes and ketones. Describe the oxidation of primary alcohols to form aldehydes and carboxylic acids. Describe the oxidation of secondary alcohols to form ketones. Describe the oxidation of aldehydes to form carboxylic acids. Describe the use of 2,4-dinitrophenylhydrazine (2,4-DNPH) to detect and identify a carbonyl compound. Describe the use of Tollens reagent to detect the presence of an aldehyde group. Describe the reduction of carbonyl compounds to form alcohols. Outline the mechanism for nucleophilic addition reactions of aldehydes and ketones with hydride. 1.1.9 12 Practical activity 2: The characteristic test for a carbonyl compound and the use of the 2,4-DNPH derivative to identify an unknown carbonyl compound Practical activity 3: Oxidation and reduction reactions of carbonyl compounds 4.1.2 Carbonyl compounds Naming of carbonyls and formation via oxidation of primary and secondary alcohols Reactions of carbonyl compounds Mechanism of nucleophilic addition 3
Week 4 1. The names and structures of carboxylic acids 2. The solubility in water due to hydrogen bonding 3. The acidic reactions e.g. with metals, bases and carbonates 4. Making esters 5. The hydrolysis of esters 6. The uses of esters Name common carboxylic acids. Explain the water solubility of carboxylic acids. Describe the reactions of carboxylic acids with metals, carbonates and bases. Describe the esterification of carboxylic acids with alcohols in the presence of an acid catalyst. Describe the reaction of acid anhydrides with alcohols to form esters. Describe the hydrolysis of esters. State the uses of esters in perfumes and flavourings. 1.1.13 14 Practical activity 4: The preparation of two esters ethyl ethanoate and methyl 2-hydroxybenzoate Practical activity 6: Hydrolysis of an ester the hydrolysis of methyl benzoate to produce benzoic acid Practical activity 9: Reactions of carboxylic acids and those of glycine 4.1.3 Carboxylic acids and esters Properties of carboxylic acids 4
Week 5 1. The structure of a triol such as propane-1, 2,3-triol 2. The structure of fatty acids such as hexadecanoic acid 3. The formation of an ester (triglyceride) from the above compounds 4. Saturated and unsaturated fats 5. Cis and trans unsaturated fats 6. The comparative healthiness of unsaturated especially trans fats 7. The increased use of fatty acid esters as biodiesel Describe a triglyceride as a triester of glycerol (propane-1,2,3-triol) and fatty acids. Compare the structures of saturated fats, unsaturated fats and fatty acids. Compare the structures of cis and trans isomers of unsaturated fatty acids. Compare the link between trans fatty acids, the possible increase in bad cholesterol and the resultant increased risk of coronary heart disease and strokes. Describe and explain the increased use of fatty acid esters as biodiesel 1.1.15 16 Practical activity 7: Reactions of ethanoic anhydride and the synthesis of aspirin (acetylsalicylic acid) 4.1.3 Carboxylic acids Esters, triglycerides, unsaturated and saturated fats 5
Week 6 1. The structural formulae of some simple amines 2. Define a base as a proton acceptor. 3. Explain that amines are bases because nitrogen s lone pair can accept a proton. 4. Examples of amines reacting with acids to form salts 5. The preparation of aliphatic amines from halogenoalkanes 6. The preparation of phenylamine by the reduction of nitrobenzene 7. The synthesis of an azo dye 8. Uses of reactions such as in (7) to form dyestuffs Explain the basicity of amines in terms of proton acceptance by the nitrogen lone pair. Describe the reactions of amines with acids to form salts. Describe the preparation of aliphatic amines by the substitution of halogenoalkanes. Describe the preparation of aromatic amines by the reduction of nitroarenes. Describe the synthesis of an azo dye by diazotisation and coupling. State the use of the azo dye reactions in the formation of dyestuffs. 1.1.17 18 Practical activity 5: The synthesis of antifebrin Practical activity 8: The reactions of amines and the preparation of azo dyes 4.1.4 Amines Reactions/formation of amines Azo dyes Uses of azo dyes 6
Week 7 1. The general formula of an α- amino acid 2. Some simple examples and structures plus common and systematic names 3. The formation of zwitterions 4. The isoelectric point and the affect of different R groups on this point 5. The acid-base properties of amino acids at different phs 6. The condensation of amino acids to form polypeptides and proteins 7. The alkaline hydrolysis of polypeptides and proteins 8. The acidic hydrolysis of polypeptides and proteins 9. Optical isomerism and chiral carbons 10. E/Z isomers and optical isomers as stereoisomers State the general formula for an α-amino acid such as RCH(NH 2 )COOH. State that an amino acid exists as a zwitterion at a ph value called the isoelectric point. State that different R groups in α-amino acids may result in different isoelectric points. Describe the acid base properties of α- amino acids at different ph values. Explain the formation of a peptide (amide) linkage between α-amino acids to form polypeptides and proteins. Describe the acid and alkaline hydrolysis of proteins and peptides. Describe optical isomers as nonsuperimposable mirror images about an organic chiral centre. Identify chiral centres in a molecule of given structural formula. Explain that optical isomerism and EIZ isomerism are types of stereoisomerism. 1.2.1 3 Practical activity 9: Reactions of carboxylic acids and those of glycine 4.2.1 Amino acids and chirality Amino acids 7
Week 8 1. Explain the term condensation polymerisation. 2. Explain polyesters with some examples, including Terylene and poly(lactic acid). 3. Explain polyamides with some examples, including Nylon-6,6 and Kevlar. 4. Practise working out the type and structure of a polymer given its monomers, and vice versa. 5. Give the use of polyesters and polyamides as fibres in clothing. 6. Compare and contrast condensation and addition polymerisation. 7. Describe the acid and base hydrolysis of condensation polymers. 8. Minimising environmental waste, e.g. degradable polymers Describe condensation polymerisation to form polyesters and polyamides such as Terylene, poly(lactic acid), Nylon-6,6 and Kevlar. State the use of polyesters and polyamides as fibres in clothing. Compare condensation polymerisation with addition polymerisation. Suggest the type of polymerisation from a given: o monomer or pair of monomers o section of a polymer molecule. Identify the monomer(s) required to form a given section of a polymer, and vice versa. Describe the acid and base hydrolysis of polyesters and polyamides. Outline the role of chemists in the development of degradable polymers. Explain that condensation polymers may be photodegradable and hydrolysed. 1.2.4 6 Practical activity 10: Nylon rope trick and the preparation of a polyester resin and a polyacrylic ester 4.2.2 Polyesters and polyamides Role of chemists in producing biodegradable plastics 8
Week 9 1. Molecules and functional groups 2. Give each student a copy of the flowcharts from the student book. 3. Explain the idea behind synthesis. 4. Discuss the presence of chiral centres in pharmaceuticals and the problems that it can cause. 5. Explain how single optical isomers can be produced and how this increases costs. Identify functional groups in an organic molecule containing several functional groups. Predict properties and reactions of an organic molecule containing several functional groups. Devise multi-stage synthetic routes for preparing organic compounds. Explain that the synthesis of pharmaceuticals often requires the production of a single optical isomer. Explain that synthetic molecules often contain a mixture of optical isomers, whereas natural molecules often only have one optical isomer. Explain that there are increased costs if the synthesised pharmaceutical is a single optical isomer. Describe strategies for the synthesis of a pharmaceutical with a single optical isomer. 1.2.7 9 4.2.3 Synthesis Synthetic routes 9
Week 10 1. Explain the terms: chromatography, mobile phase and stationary phase. 2. Describe separation by adsorption and by relative solubility. 3. Describe thin layer chromatography (TLC). 4. Describe gas chromatography (GC). 5. Explain R f values and retention time. 6. Describe the extra usefulness of GC-MS and the uses to which it can be put Describe chromatography as an analytical technique that separates components in a mixture between a mobile phase and a stationary phase. State that the mobile phase may be a liquid or a gas. State that the stationary phase may be a solid, or either a liquid or solid on a solid support. State that a solid stationary phase separates by adsorption. State that a liquid stationary phase separates by relative solubility. State that the mobile phase in TLC is a liquid and the stationary phase is a solid on a solid support and that the solid stationary phase in TLC separates by adsorption. Explain the term R f value and interpret chromatograms in terms of R f values. Explain the term retention time and interpret gas chromatograms in terms of retention times and the approximate proportions of the components of a mixture. Explain that analysis by gas chromatography 1.3.1 4 Practical activity 11: Thin layer and paper chromatography 4.3.1 Chromatography 10
has limitations. Explain that mass spectrometry can be combined with chromatography in GC-MS to provide a far more powerful analytical tool than from gas chromatography alone. Explain that the mass spectra generated can be analysed or compared with spectral databases for positive identification of a component. State the use of GC-MS in analysis e.g. in forensics, environmental analysis, airport security and space probes. 11
Week 11 1. Introduction and brief explanation of nuclear magnetic resonance (NMR) 2. Tetramethylsilane (TMS) standard and the need for deuterated solvents 3. Carbon-13 NMR different types of carbon and chemical shifts and how to use a data sheet 4. Using carbon-13 NMR to predict possible structures 5. Go through the worked examples for carbon-13 NMR in the student book. 6. Proton NMR different types of proton, relative peak areas and chemical shifts 7. The use of proton NMR to make predictions about structures 8. Go through the worked examples for proton NMR in the student book. State that nuclear magnetic resonance (NMR) spectroscopy involves the interaction of materials with the low-energy radio wave radiation. Describe the use of tetramethylsilane (TMS) as the standard for chemical shift. State the need for deuterated solvents such as CDCl 3 when running an NMR spectrum. Analyse carbon-13 NMR spectra to make predictions about the different types of carbon atoms present. Predict the chemical shifts of carbons in a given molecule. Analyse carbon-13 NMR spectra to make predictions about possible structures for an unknown compound. Analyse a proton NMR spectrum to make predictions about: o the different types of proton present o the relative numbers of each type of proton present from relative peak areas and chemical shifts o possible structures for the molecule. Predict the chemical shifts of the protons in a given molecule. 1.3.5 8 4.3.2 Spectroscopy 12
Week 12 1. Explain what is meant by splitting. 2. How does splitting arise? 3. Explain how to use the n+1 rule to determine the number of protons on the adjacent carbon. 4. Explain how to predict the splitting pattern in a given molecule. 5. Go through the worked example in the student book. 6. Explain the use of deuterium (D 2 O) to identify OH and NH protons. 7. Go through the worked example in the student book. 8. Explain the similarities between NMR spectroscopy and magnetic resonance imaging (MRI). Analyse a high-resolution proton NMR spectrum to make predictions about: o the number of non-equivalent protons adjacent to a given proton o possible structures for the molecule. Predict the splitting patterns of the protons in a given molecule. Describe the identification of OH and NH protons by proton exchange using deuterium (D 2 O). Explain that NMR spectroscopy is the same technology as that used in magnetic resonance imaging (MRI). 1.3.9 12 4.3.2 Spectroscopy NMR Spectroscopy 13
Week 13 1. Go through the infrared (IR) absorption peaks on the data sheets. 2. Use the data sheets to identify the presence or absence of peaks from the data sheet on various IR spectra. 3. Explain the use of molecular ion peaks in mass spectra. 4. Explain the use of fragment peaks in mass spectra. 5. Identify the various peaks in mass spectra and suggest a structure. 6. Explain the limitations and advantages of each spectroscopic technique. 7. Discuss the advantages of combining spectroscopic techniques. 8. Go through some examples and get students to try some. Analyse infrared absorptions in an infrared (IR) spectrum in order to identify the presence of functional groups in an organic compound. Analyse molecular ion peaks and fragmentation peaks in a mass spectrum in order to identify parts of an organic structure. Combine evidence from NMR, IR and mass spectra to deduce organic structures. 1.3.13 14 4.3.2 Spectroscopy Combined techniques 14
Week 14 1. Revise AS work on rates of reaction. 2. Explain and define the rate of a reaction. 3. Describe how some rates are proportional to concentrations i.e. first order. 4. Describe how some rates are proportional to concentrations squared i.e. second order. 5. Define: order of reaction. 6. Deduce rate equations from orders. 7. Explain calculating the rate constant including its units. 8. Explain how concentration time graphs can be plotted from experimental data and used to measure rates. 9. Describe experimental methods for obtaining rate data. Explain and use the terms: rate of reaction, order and rate constant. Deduce the rate of a reaction from a concentration time graph. Plot a concentration time graph from experimental results. Deduce a rate equation from orders. 2.1.1 3 Practical activity 12: The reaction between calcium carbonate and hydrochloric acid solution monitoring gas loss or mass loss Practical activity 13: The rate of reaction between propanone and iodine 5.1.1 How fast? Rate graphs and orders 15
Week 15 1. Define the terms: rate constant; order of reaction; and half-life. 2. Explain that the half-life of a first-order reaction is constant it does not depend on the concentration. 3. Deduce the half-life of a first-order reaction from a concentration time graph. 4. Use rate concentration graphs to deduce the order of a reaction. Explain and use the terms: order and halflife. Deduce the half-life of a first-order reaction from a concentration time graph. State that the half-life of a first-order reaction is independent of the concentration. Deduce the order (0, 1 or 2) with respect to a reactant from a rate concentration graph. 2.1.4 5 Practical activity 14: The rate of the reaction between iodine and the persulfate ion [peroxodisulfate(vi)] Practical activity 15: The rate of the reaction between sodium thiosulfate and hydrochloric acid 5.1.1 How fast? Rate graphs and orders 16
Week 16 1. Use the initial rates method to deduce the order of a reaction. 2. Using orders of reaction, deduce a rate equation. 3. Calculate the rate constant (including units) from a rate equation. 4. Describe and explain the affect of temperature on the rate of a reaction and thus on the rate constant. 5. Explain multi-step reactions and the rate-determining step. 6. Explain why rate equations and the rate-determining step must be consistent. 7. Show how mechanisms can be proposed that are consistent with both the rate equation and the stoichiometric equation. Determine, using the initial rates method, the order (0, 1 or 2) with respect to a reactant. Deduce from orders a rate equation of the form: rate = k[a] m [B] n. Calculate the rate constant k from a rate equation. Explain the effect of temperature change on a rate constant and rate of a reaction. Propose a rate equation that is consistent with the rate-determining step in a multi-step reaction. Propose steps in a reaction mechanism in a multi-step reaction from the rate equation and the balanced equation for the overall reaction. 2.1.6 7 Practical activity 16: The effect of temperature on the rate of a reaction 5.1.1 How fast? Rate equations; Rate constants Rate-determining step 17
Week 17 1. Revise equilibrium from AS course. 2. Explain the characteristics of dynamic equilibria. 3. Explain the equilibrium constant. 4. Derive expressions for the equilibrium constant from equations for equilibria. 5. Give the units for the equilibrium constant for specific equilibria. 6. Use these expressions and concentration data to calculate values for the equilibrium constant. 7. Given initial concentrations and one equilibrium concentration; derive other equilibrium concentrations and a value for the equilibrium constant. Deduce for homogeneous reactions expressions for the equilibrium constant K c. Calculate the values of the equilibrium constant K c including the determination of units. Calculate the concentration or quantities present at equilibrium. 2.1.8 9 Practical activity 17: Determination of the K c for the ethanoic acid/ethyl ethanoate equilibrium 5.1.2 How far? Equilibrium 18
Week 18 1. Discuss what high and low values for K c signify. 2. Revise le Chatelier s principle. 3. Explain the effect of temperature changes on the position of equilibrium and therefore its effect on the value of K c do for both exothermic and endothermic forward reactions. 4. State that K c is unaffected by changes in concentration, pressure or the presence of catalysts. 5. Explain how equilibria shift to keep K c constant when concentrations (or pressures) are changed and how this explains this aspect of le Chatelier s principle. 6. Explain how industry must use conditions which are a compromise between rate and equilibrium yield (as well as safety and cost) e.g. the Haber process. 7. Link the compromise conditions to the values of K c and k. Understand the significance of K c values. Explain the effect of changing temperature on the value of K c. State that the value of K c is unaffected by changes in concentration, pressure or the presence of a catalyst. Make predictions for shifts in equilibrium position from concentration and pressure changes. Understand that compromise conditions rely on a balance between K c and k. 2.1.10 11 5.1.2 How far? Equilibrium 19
Week 19 1. Revise work already covered on acids and bases. 2. Detail the historical development of our model of acids and bases. 3. Appreciate that this model may change again scientific knowledge is always evolving! 4. Provide definitions of acids and bases with a few examples as equations. 5. The reactions of acids with equations 6. Ionic equations for the reactions in point 5 7. Conjugate pairs Describe how the model of an acid has changed over the years. Understand that scientific knowledge is always evolving. Describe an acid as a species that can donate a proton, and a base as a species that can accept a proton. Illustrate the role of H + in the reactions of acids with carbonates, bases, alkalis and metals. Describe and use the term: conjugate acid base pairs. 2.1.12 14 5.1.3 Acids, bases and buffers 20
Week 20 1. Explain the convenience of ph. 2. Define ph as ph = log[h + (aq)]. 3. Define [H + ] = 10 ph. 4. Explain how to calculate ph from [H + ] and vice versa using the students' own calculators. 5. Explain the difference between strong and weak acids. 6. Calculate the ph of a strong acid. 7. Deduce the acid dissociation constant K a. 8. Explain the meaning of pk a and its interconversion with K a. 9. Explain how to calculate the ph of a weak acid. 10. Explain how to calculate K a for a weak acid. Define ph as ph = log[h + (aq)]. Define [H + ] = 10 ph. Convert between ph values and [H + (aq)]. Explain the differences between strong and weak acids. Explain that the acid dissociation constant K a shows the extent of acid dissociation. Deduce expressions for K a and pk a for weak acids. Convert between K a and pk a. Calculate ph from [H + (aq)] and [H + (aq)] from ph for strong and weak acids. Calculate K a for a weak acid. 2.1.15 17 Practical activity 18: Finding the ph of strong and weak acids 5.1.3 Acids, bases and buffers 21
Week 21 1. Define a buffer solution. 2. Give uses of buffer solutions. 3. Describe how a buffer solution can be made. 4. Explain how a buffer solution works. 5. Calculate the ph of buffer solutions. 6. Explain the buffer system in blood. 7. Explain acid base titration curves drawing and interpreting them for strong and weak acids and bases. 8. Discuss how to choose indicators. 9. Define the term: standard enthalpy change of neutralisation. 10. Explain how enthalpy change of neutralisation can be calculated from experimental data. Describe what is meant by a buffer solution. State that a buffer solution can be made from a weak acid and a salt of the weak acid. Explain the role of the conjugate acid base pair in an acid buffer solution. Calculate the ph of a buffer solution from the K a value of a weak acid and the equilibrium concentrations of the conjugate acid base pair. Explain the role of carbonic acid-hydrogencarbonate as a buffer in the control of blood ph. Interpret and sketch acid-base titration ph curves for strong and weak acids and bases. Explain the choice of suitable indicators for acid-base titrations. Define and use the term: standard enthalpy change of neutralisation. Calculate enthalpy changes from appropriate experimental results. 2.1.20 23 Practical activity 19: An investigation of buffer solutions Practical activity 20: Generating acid base curves with a datalogger 5.1.3 Acids, bases and buffers Buffers: Action, uses and calculations Neutralisation 22
Week 22 1. Introduce the concept of lattice enthalpy as an exothermic process. 2. Explain that lattice enthalpies cannot be measured experimentally. 3. Introduce the Born Haber cycle as an example of Hess law that can be used to calculate lattice enthalpies. 4. Go through an example of a Born Haber cycle and the associated calculation. 5. Get the students to practise their own Born Haber cycles. Explain and use the term: lattice enthalpy. Use the lattice enthalpy of a simple ionic solid and relevant energy terms to construct Born Haber cycles. Carry out related calculations to calculate lattice enthalpies. 2.2.1 4 5.2.1 Lattice enthalpy 23
Week 23 1. Explain what is happening when an ionic compound dissolves in water. 2. Discuss hydration and the factors that affect the magnitude of the enthalpy of hydration. 3. Explain how enthalpy of solution depends on the relative values of the lattice enthalpy and the enthalpy of hydration and how this may determine whether or not an ionic compound is soluble. 4. Carry out calculations based on point 3 above. 5. Explain how ionic size and ionic charge effect lattice enthalpies. Use the enthalpy change of solution of a simple ionic solid and relevant energy terms to construct Born Haber cycles. Carry out related calculations. Explain in qualitative terms the effect of ionic charge and ionic radius on lattice enthalpy and enthalpy change of hydration. 2.2.5 6 Practical activity 21: Measuring enthalpy changes of solution Practical activity 22: Measuring the enthalpy of neutralisation using a thermometric titration 5.2.1 Lattice enthalpy 24
Week 24 1. Introduce entropy as a measure of disorder. 2. Explain all spontaneous changes produce an overall increase in entropy. 3. Compare the value of the entropy in different states of the same amount of substance. 4. Explain how to calculate entropy changes for reactions given the relevant entropies of the reactants and the products. 5. Explain that reactions take place only if enthalpy change, entropy change and temperature are correct. 6. Introduce free energy and G = H T S. 7. Explain reactions only take place when the free energy is greater than zero. 8. Explain how endothermic reactions can take place spontaneously. 9. Explain how to calculate the free energy for some reactions at given temperatures. Explain that entropy is a measure of the disorder of a system and that a system becomes energetically more stable when it becomes more disordered. Explain the difference in entropy: (i) of a solid and a gas; (ii) when a solid lattice dissolves; and (iii) for a reaction in which there is a change in the number of gaseous molecules. Calculate the entropy change for a reaction given the entropies of reactants and product. Explain that the tendency of a process to take place depends on: (i) absolute temperature T; (ii) the entropy change in the system S; and (iii) the enthalpy change H with the surroundings. Explain that the balance between entropy and enthalpy changes is the free energy change G. State and use the relationship: G = H T S Explain how endothermic reactions are able to take place spontaneously. 2.2.7 8 5.2.2 Enthalpy and entropy 25
Week 25 1. Revise redox and oxidation numbers from AS. 2. Introduce redox half-equations and use them to create balanced equations for redox reactions. 3. Introduce half-cells include the various types. 4. Explain how half-cells can be put together to form electrochemical cells. 5. Introduce standard electrode potentials. 6. Explain the role of a hydrogen electrode in measuring other standard electrode potentials. 7. Use standard electrode potentials to calculate the e.m.f. of electrochemical cells. 8. Use standard electrode potentials to predict the feasibility of a reaction. 9. Explain the limitations of these predictions. Explain for simple redox reactions the terms: redox; oxidation number; half-reaction; oxidising agent; and reducing agent. Construct redox equations using relevant half-equations or oxidation numbers. Interpret and make predictions for reactions involving electron transfer. Describe simple half-cells made from: ο metals or non-metals in contact with their ions in aqueous solution ο ions of the same element in different oxidation states. Describe how half-cells can be combined to make an electrochemical cell. Define the term: standard electrode (redox) potential, E o. Describe how to measure standard electrode potentials using a hydrogen electrode. Calculate a standard cell potential by combining two standard electrode potentials. Predict, using standard cell potentials, the feasibility of a reaction. Consider the limitations of predictions made using standard cell potentials in terms of kinetics and concentration 2.2.9 12 Practical activity 23: Redox reactions Practical activity 24: Constructing electrochemical cells and measuring electrode potentials 5.2.3 Electrode potentials and fuel cells Electrode potentials 26
Week 26 1. Discuss modern cells and batteries (non-rechargeable and rechargeable) and fuel cells. 2. Explain the hydrogen oxygen fuel cell including the electrode reactions. 3. Discuss hydrogen-rich fuels. 4. Discuss the modern developments in fuel cells including fuel cell vehicles (FCVs). 5. Discuss the advantages of fuel cells. 6. Discuss the limitations of fuel cells. 7. Consider the hydrogen economy. Apply principles of electrode potentials to modern storage and fuel cells. Explain that a fuel cell uses the energy from the reaction between a fuel and oxygen to create a voltage. Explain the changes that take place at each electrode in a fuel cell. Outline the development of fuel cell vehicles (FCVs) fuelled by hydrogen gas and hydrogen-rich fuels. State advantages of FCVs over conventional petrol or diesel-powered vehicles. Understand how hydrogen might be stored in FCVs. Consider the limitations of hydrogen fuel cells. Comment about the contribution of the hydrogen economy to future energy and discuss any limitations. Discuss the hydrogen oxygen fuel cell. 2.2.13 14 Practical activity 25: Making a fuel cell 5.2.3 Electrode potentials and fuel cells Storage and fuel cells 27
Week 27 1. Define transition elements and d-block elements. 2. Use Sc and Zn as examples of d-block elements which are not transition elements. 3. Explain the order of filling the sub-shells and therefore electron configurations of the elements. 4. Show that Cr and Cu are exceptions to the pattern shown in point 3. 5. Explain the electron configurations of the ions. 6. Discuss the typical properties of transition elements physical properties; coloured compounds; variable oxidation states, etc. 7. Discuss transition metals as catalysts with examples. 8. Explain precipitation reactions to form insoluble metal hydroxides. Deduce the electron configurations of atoms and ions of the d-block elements. Describe the elements Ti through to Cu as transition elements. Show more than one oxidation state for a transition element in its compounds. Know that transition metal ions are coloured. Illustrate the catalytic behaviour of the elements and/or their compounds. Describe the simple precipitation reactions of Cu 2+ (aq), Co 2+ (aq), Fe 2+ (aq) and Fe 3+ (aq) with aqueous sodium hydroxide. 2.3.1 3 Practical activity 26: Precipitation of transition metal hydroxides 5.3.1 Transition elements Precipitation reactions 28
Week 28 1. Explain with examples the terms: ligand, complex ion, coordination number and coordinate bonding. 2. Introduce formulae such as [Cu(H 2 O) 6 ] 2+. 3. Explain the conventions for drawing 3D octahedral complexes. 4. Revise stereoisomerism cis trans and optical isomerism. 5. Explain cis trans isomerism in octahedral and square planar complex ions. 6. Discuss the use of cis platin. 7. Discuss bidentate ligands and cis trans isomerism. 8. Discuss bidentate ligands and optical isomerism. Explain the term: ligand in terms of coordinate bonding. State and use the terms: complex ion and coordination number. State and give examples of complexes with six-fold coordination with an octahedral shape. Describe the use of cis platin as an anticancer drug and its action of binding to DNA. Explain and use the term: bidentate ligand (i.e. NH 2 CH 2 CH 2 NH 2, en). Describe stereoisomerism in transition metal multi-dentate complexes using examples of cis trans and optical isomerism. 2.3.4 6 5.3.1 Transition elements Ligands and complex ions 29
Week 29 1. Explain what is meant by ligand substitution. 2. Demonstrate the reactions of the hydrated copper(ii) ion with ammonia and with chloride ions. 3. Demonstrate the reactions of the hydrated cobalt(ii) ion with ammonia and with chloride ions. 4. Discuss haemoglobin and its role in carrying oxygen and carbon dioxide. 5. Discuss carbon monoxide, the 'silent killer'. 6. Explain stability constants. Describe the process of ligand substitution. Describe examples of the ligand substitution of [Cu(H 2 O) 6 ] 2+ and [Co(H 2 O) 6 ] 2+ with ammonia and chloride ions. Explain the biochemical importance of iron in haemoglobin including ligand substitution. Use and understand the term: stability constant, K stab. 2.3.7 8 Practical activity 27: Ligand substitution reactions 5.3.1 Transition elements Ligand substitution 30
Week 30 1. Explain redox in transition metals use both oxidation numbers and transfer of electrons. 2. Explain redox titrations involving potassium permanganate including calculations initially structured and then unstructured. 3. Explain redox titrations involving iodine and sodium thiosulfate including calculations initially structured and then unstructured. Describe, using suitable examples, redox behaviour in transition elements. Understand how to carry out redox titrations and structured calculations involving MnO 4 and I 2 /S 2 O 2 3. Perform non-structured titration calculations based on experimental results. 2.3.9 11 Practical activity 28: The estimation of the percentage of iron in iron tablets Practical activity 29: The estimation of the percentage of copper in brass 5.3.1 Transition elements Redox reactions and titrations 31