Week 1 Weekly learning outcomes Student Practical … to form 2,4,6-tribromophenol. • Explain the...

© Pearson Education Ltd 2009 This document may have been altered from the original 1

Week 1 Weekly learning outcomes Student book links Practical activity links

• 1.1.1–5 • Practical activity 1: The nitration of methyl benzoate – to form methyl 3-nitrobenzoate

OCR Scheme of Work topic outlines

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

Students should be able to: • 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.

4.1.1 Arenes • Structure of benzene • Electrophilic substitution of arenes

Teaching scheme



• 1.1.6–8


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

Students should be able to: • 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.

4.1.1 Arenes • Phenols



• 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


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 NaBH4

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

Students should be able to: • 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.

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



• 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


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

Students should be able to: • 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. 4.1.3 Carboxylic acids and esters

• Properties of carboxylic acids



• 1.1.15–16 • Practical activity 7: Reactions of ethanoic anhydride and the synthesis of aspirin (acetylsalicylic acid)


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

Students should be able to: • 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

4.1.3 Carboxylic acids • Esters, triglycerides, unsaturated and

saturated fats



• 1.1.17–18 • Practical activity 5: The synthesis of antifebrin

• Practical activity 8: The reactions of amines and the preparation of azo dyes


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

Students should be able to: • 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.

4.1.4 Amines • Reactions/formation of amines • Azo dyes • Uses of azo dyes



• 1.2.1–3 • Practical activity 9: Reactions of carboxylic acids and those of glycine


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

Students should be able to: • State the general formula for an α-amino acid

such as RCH(NH2)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 non-superimposable 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.

4.2.1 Amino acids and chirality • Amino acids



• 1.2.4–6 • Practical activity 10: Nylon rope trick and the preparation of a polyester resin and a polyacrylic ester


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

Students should be able to: • 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.

4.2.2 Polyesters and polyamides • Role of chemists in producing biodegradable

plastics



• 1.2.7–9


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.

Students should be able to: • 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.

4.2.3 Synthesis • Synthetic routes



• 1.3.1–4 • Practical activity 11: Thin layer and paper chromatography


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 Rf values and retention time.

6. Describe the extra usefulness of GC-MS and the uses to which it can be put

Students should be able to: • 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 Rf value and interpret chromatograms in terms of Rf 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

4.3.1 Chromatography


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.



• 1.3.5–8


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.

Students should be able to: • 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 CDCl3 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.

4.3.2 Spectroscopy



• 1.3.9–12


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 (D2O) 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).

Students should be able to: • 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 (D2O).

• Explain that NMR spectroscopy is the same technology as that used in magnetic resonance imaging (MRI).

4.3.2 Spectroscopy • NMR Spectroscopy



• 1.3.13–14


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.

Students should be able to: • 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.

4.3.2 Spectroscopy • Combined techniques



• 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


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.

Students should be able to: • 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. 5.1.1 How fast?

• Rate graphs and orders



• 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


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.

Students should be able to: • Explain and use the terms: order and half-

life. • 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. 5.1.1 How fast?

• Rate graphs and orders



• 2.1.6–7 • Practical activity 16: The effect of temperature on the rate of a reaction


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.

Students should be able to: • 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.

5.1.1 How fast? • Rate equations; Rate constants • Rate-determining step



• 2.1.8–9 • Practical activity 17: Determination of the Kc for the ethanoic acid/ethyl ethanoate equilibrium


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.

Students should be able to: • Deduce for homogeneous reactions

expressions for the equilibrium constant Kc. • Calculate the values of the equilibrium

constant Kc including the determination of units.

• Calculate the concentration or quantities present at equilibrium.

5.1.2 How far? • Equilibrium



• 2.1.10–11


1. Discuss what high and low values for Kc 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 Kc – do for both exothermic and endothermic forward reactions.

4. State that Kc is unaffected by changes in concentration, pressure or the presence of catalysts.

5. Explain how equilibria shift to keep Kc 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 Kc and k.

Students should be able to: • Understand the significance of Kc values. • Explain the effect of changing temperature

on the value of Kc. • State that the value of Kc 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 Kc and k.

5.1.2 How far? • Equilibrium



• 2.1.12–14


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

Students should be able to: • 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.

5.1.3 Acids, bases and buffers



• 2.1.15–17 • Practical activity 18: Finding the pH of strong and weak acids


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 Ka . 8. Explain the meaning of pKa and its

interconversion with Ka. 9. Explain how to calculate the pH of

a weak acid. 10. Explain how to calculate Ka for a

weak acid.

Students should be able to: • 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

Ka shows the extent of acid dissociation. • Deduce expressions for Ka and pKa for weak

acids. • Convert between Ka and pKa. • Calculate pH from [H+(aq)] and [H+(aq)] from

pH for strong and weak acids. • Calculate Ka for a weak acid.

5.1.3 Acids, bases and buffers



• 2.1.20–23 • Practical activity 19: An investigation of buffer solutions

• Practical activity 20: Generating acid–base curves with a datalogger


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.

Students should be able to: • 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

Ka 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.

5.1.3 Acids, bases and buffers • Buffers: Action, uses and calculations • Neutralisation



• 2.2.1–4


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.

Students should be able to: • 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.

5.2.1 Lattice enthalpy



• 2.2.5–6 • Practical activity 21: Measuring enthalpy changes of solution

• Practical activity 22: Measuring the enthalpy of neutralisation using a thermometric titration


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.

Students should be able to: • 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.

5.2.1 Lattice enthalpy



• 2.2.7–8


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.

Students should be able to: • 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.

5.2.2 Enthalpy and entropy



• 2.2.9–12 • Practical activity 23: Redox reactions

• Practical activity 24: Constructing electrochemical cells and measuring electrode potentials


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.

Students should be able to: • 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

5.2.3 Electrode potentials and fuel cells • Electrode potentials



• 2.2.13–14 • Practical activity 25: Making a fuel cell


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.

Students should be able to: • 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.

5.2.3 Electrode potentials and fuel cells • Storage and fuel cells



• 2.3.1–3

• Practical activity 26: Precipitation of transition metal hydroxides


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.

Students should be able to: • 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

Cu2+(aq), Co2+(aq), Fe2+(aq) and Fe3+(aq) with aqueous sodium hydroxide.

5.3.1 Transition elements • Precipitation reactions



• 2.3.4–6


1. Explain – with examples – the terms: ligand, complex ion, coordination number and coordinate bonding.

2. Introduce formulae such as [Cu(H2O)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.

Students should be able to: • 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. NH2CH2CH2NH2, en).

• Describe stereoisomerism in transition metal multi-dentate complexes using examples of cis–trans and optical isomerism.

5.3.1 Transition elements • Ligands and complex ions



• 2.3.7–8 • Practical activity 27: Ligand substitution reactions


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.

Students should be able to: • Describe the process of ligand substitution. • Describe examples of the ligand substitution

of [Cu(H2O)6]2+ and [Co(H2O)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, Kstab.

5.3.1 Transition elements • Ligand substitution



• 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


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.

Students should be able to: • Describe, using suitable examples, redox

behaviour in transition elements. • Understand how to carry out redox titrations

and structured calculations involving MnO4–

and I2/S2O32–.

• Perform non-structured titration calculations based on experimental results.

5.3.1 Transition elements • Redox reactions and titrations

Date post:	12-May-2018
Category:	Documents
Upload:	vudan
View:	219 times
Download:	5 times

Week 1 Weekly learning outcomes Student Practical … to form 2,4,6-tribromophenol. • Explain the...

Documents