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OCR Chemistry A Scheme of Work

© Pearson Education Ltd 2016 This document may have been altered from the original

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Teacher Resource Pack

Week 1 Links to prior learning Student Book links Teaching plan links

1. What is meant by ‘rate of reaction’ 2. Order of reactions – the three main types 3. Rate equations and the rate constant, k 4. Calculating a rate constant and determining

units 5. Using graphs to calculate the rate of a

reaction 6. Half-life and its use in determining the order

of a reaction 7. Calculating a rate constant using half-life

● Factors affecting rate of reaction and why they affect the rate of reaction

● How collision theory can be used to predict how changes in conditions influence rates of reactions

● Different reactions occur at different rates ● How to monitor the rate of a reaction by

collecting experimental results ● Calculating a rate using a graph of a

physical quantity changing over time

● 5.1.1 ● 5.1.2

● 5.1.1 ● 5.1.2

Weekly learning outcomes Specification links Practical activity links

Students should be able to demonstrate and apply their knowledge and understanding of: ● the terms: rate of reaction, order, overall

order, rate constant ● the deduction of:

(i) orders from experimental data

(ii) a rate equation from orders in the form: rate = k[A]m [B]n, where m and n are 0, 1 or 2

● calculating a rate constant, k, and related quantities, from a rate equation, including determining its units

● 5.1.1 (a)–(c) ● 5.1.1 (d)–(f) ● 5.1.1 (h)

● Practical 3: Use a clock reaction to determine a rate equation



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● a concentration–time graph: (i) deduction of the order

(0 or 1) with respect to a reactant from the shape of the graph

(ii) calculation of reaction rates from the measurement of gradients.

● a concentration–time graph of a first order reaction and measurement of constant half-life, 1

2t

● a first order reaction and determination of the rate constant, k, from the constant half-

life, 12

t , using the relationship: k = 12

ln 2t

● the techniques and procedures used to investigate reaction rates by the initial rates method and by continuous monitoring, including the use of colorimetry.



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1. Initial rate and how it is determined 2. Changes that take place during a reaction 3. Orders and rate–concentration graphs 4. Determining a rate constant 5. Explanation of ‘rate-determining step’ 6. Using rate equations to predict reaction

mechanisms 7. The affect of the order of a reactant on the

rate-determining step 8. Working out reaction mechanisms

● Calculating a rate of reaction using a graph of a physical quantity changing over time

● The effect of concentration and temperature on the rate of a reaction

● Organic reaction mechanisms ● Rates of reaction ● Balancing equations

● 5.1.3 ● 5.1.4

● 5.1.3 ● 5.1.4


Students should be able to demonstrate and apply their knowledge and understanding of: ● a rate–concentration graph:

(i) deduction of the order (0, 1 or 2) with respect to a reactant from the shape of the graph

● determination of the rate constant for a first order reaction using the gradient

● the use of the term ‘rate-determining step’ ● a multi-step reaction, prediction of:

(i) a rate equation that is consistent with the rate-determining step

(ii) possible steps in a reaction mechanism from the rate equation and the balanced equation for the overall reaction.

● 5.1.1 (g) ● 5.1.1 (i)

● Practical 8: Follow the rate of the iodine–propanone reaction using a titrimetric method



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1. Rate of reaction and the value of a rate constant, k

2. How temperature affects k 3. Understanding the Arrhenius equation 4. Arrhenius plots 5. Equilibrium and how it is achieved 6. The equilibrium constant Kc 7. Determining the units for Kc 8. Practical methods for determining

concentrations at equilibrium 9. Calculations involving Kc and unknown

equilibrium concentrations

● How collision theory can be used to predict the influence of changes in conditions on rates of reactions

● The effect that temperature has on the rate of a reaction

● Boltzmann distributions ● Dynamic equilibrium ● Writing an equilibrium expression ● What the magnitude of Kc indicates about

the position of equilibrium ● Le Chatelier’s principle ● Heterogeneous and homogenous reactions

● 5.1.5 ● 5.1.6

● 5.1.5 ● 5.1.6


Students should be able to demonstrate and apply their knowledge and understanding of: ● a qualitative explanation of the effect of

temperature change on the rate of a reaction, and hence its rate constant

● the Arrhenius equation: (i) the exponential

relationship between rate constant, k, and temperature, T, given by the Arrhenius

equation, k = − a

eERTA

(ii) determine Ea and A graphically using:

ln k = a− ERT

+ ln A

derived from the Arrhenius equation

● 5.1.1 (j), (k) ● 5.1.2 (b)−(e)

● Practical 11: Determine the activation energy for the reaction between bromide ions and bromate(V) ions



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● the quantities of reactants and products present at equilibrium, given appropriate data

● the techniques and procedures used to determine quantities present at equilibrium

● expressions for Kc and Kp for homogeneous and heterogeneous equilibria

● calculations involving Kc and Kp, or related quantities, including determination of units.



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1. Mole fraction 2. Partial pressure 3. Introduction to Kp 4. Calculating Kp 5. The value of the equilibrium constant, K,

and the position of equilibrium 6. The effect of temperature on K 7. The effect of concentration, pressure and

catalysts on K

● Dynamic equilibrium ● Writing an equilibrium expression ● What the magnitude of Kc can tell you about

the position of equilibrium ● Le Chatelier’s principle ● Heterogeneous and homogenous reactions ● The effect of catalysts on the rate of

reaction, including lowering the activation barrier

● The link between the magnitude of Kc and the position of equilibrium

● The effect of temperature on equilibrium

● 5.1.7 ● 5.1.8

● 5.1.7 ● 5.1.8


Students should be able to demonstrate and apply their knowledge and understanding of: ● the terms ‘mole fraction’ and ‘partial

pressure’ ● expressions for Kc and Kp for homogeneous

and heterogeneous equilibria ● calculations involving Kc and Kp, or related

quantities, including determination of units ● the qualitative effect of changing

temperature on equilibrium constants for exothermic and endothermic reactions

● the constancy of equilibrium constants with changes in concentration and pressure, and in the presence of a catalyst

● 5.1.2 (a) ● 5.1.2 (d)−(e) ● 5.1.2 (f)–(h)



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● how an equilibrium constant controls the position of equilibrium in response to changes in concentration, pressure and temperature

● the above principles for Kc and Kp in relation to other equilibrium constants, where appropriate.



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1. The history of scientific understanding of acids and bases

2. What is meant by Brønsted–Lowry acids and bases

3. Examples of Brønsted–Lowry acids and bases

4. Ionic equations 5. Mono-, di- and tribasic acids 6. Reactions of acids with carbonates, bases

and alkalis 7. Redox reactions of acids with metals 8. Acid strength and dissociation 9. Acid dissociation constant, Ka 10. Ka and pKa

● Acid and base behaviour ● Reactions of acids and bases ● Neutralisation ● Equilibrium constants

● 5.1.9 ● 5.1.10

● 5.1.9 ● 5.1.10


Students should be able to demonstrate and apply their knowledge and understanding of: ● a Brønsted–Lowry acid as a species that

donates a proton and a Brønsted–Lowry base as a species that accepts a proton

● the term conjugate acid–base pairs ● monobasic, dibasic and tribasic acids ● the role of H+ in the reactions of acids with

metals and bases, including carbonates, metal oxides and alkalis, using ionic equations

● the acid dissociation constant, Ka, for the extent of acid dissociation

● the relationship between Ka and pKa ● the principles for Kc and Kp to other

equilibrium constants, where appropriate.

● 5.1.3 (a) (i)–(iii) ● 5.1.2 (h) ● 5.1.3 (b), (c)

● Practical 12: Generate acid–base curves using a datalogger



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1. Development of the pH scale 2. Defining pH as pH = −log [H+(aq)] 3. Defining [H+] as [H+] = 10−pH 4. Converting pH values to hydrogen ion

concentration and vice versa 5. Calculating the pH of strong acids 6. Calculating the pH of weak acids 7. The limitations of using approximations for

the pH of weak acids 8. Definition and expression of the ionic

product of water 9. Units of Kw 10. Kw of pure water 11. The role of Kw in determining the

concentration of H+(aq) and OH−(aq) ions in solution

12. Using the relationship between H+ and OH− to calculate their concentrations for different pH values

13. Calculating the pH of strong bases

● The pH scale ● Acids and bases ● Calculating Ka values ● Calculating values for pKa ● Le Chatelier’s principle ● pH and [H+(aq)]

● 5.1.11 ● 5.1.12

● 5.1.11 ● 5.1.12


Students should be able to demonstrate and apply their knowledge and understanding of: ● the expression for pH as pH = −log[H+]

[H+] = 10−pH ● calculations involving pH, or related

quantities, for strong monobasic acids ● calculations involving pH, Ka or related

quantities, for a weak monobasic acid using approximations

● the limitations of using approximations to Ka-related calculations for ‘stronger’ weak acids

● the expression for the ionic product of water, Kw

● calculations involving pH, or related quantities, for strong monobasic acids and strong bases using Kw

● the principles for Kc and Kp to other equilibrium constants, where appropriate.

● 5.1.3 (d) ● 5.1.3 (f)–(h) ● 5.1.3 (e),(f)

● Practical 13: Determine Ka for a weak acid.



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1. What is a buffer? 2. Making buffer solutions 3. How buffers work 4. Calculating the pH of a buffer solution 5. Buffering in the blood 6. Explanation of acid–base titration curves,

including drawing and interpreting them for strong against weak acids and bases

7. Discussion of equivalence points 8. Indicators and end points 9. How to choose a suitable indicator

● Acids and bases ● pH ● Strength of acids and alkalis ● Ka ● Indicators

● 5.1.13 ● 5.1.14

● 5.1.13 ● 5.1.14


Students should be able to demonstrate and apply their knowledge and understanding of: ● a buffer solution as a system that minimises

pH changes on addition of small amounts of an acid or a base

● the formation of a buffer solution from: (i) a weak acid and a salt

of the weak acid, e.g. CH3COOH/CH3COONa

(ii) excess of a weak acid and a strong alkali, e.g. excess CH3COOH/NaOH

● the role of the conjugate acid–base pair in an acid buffer solution, e.g. CH3COOH/CH3COO−, in the control of pH

● calculations involving the pH of buffer solutions, from the Ka value of a weak acid and the equilibrium concentrations of the conjugate acid–base pair; calculations of related quantities

● control of blood pH by the carbonic acid–hydrogencarbonate buffer system

● 5.1.3 (i)–(m) ● 5.1.3 (n), (o)

● Practical 12: Generate acid–base curves using a datalogger



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● pH titration curves for combinations of strong and weak acids against strong and weak bases, including:

(i) sketching and interpreting their shapes

(ii) explaining the choice of suitable indicators, given the pH range of the indicator

(iii) explaining indicator colour changes in terms of equilibrium shift between the HA and A− forms of the indicator

● the techniques and procedures used when measuring pH using a pH meter.



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1. Recap ionic compounds, including their formation

2. Definitions and explanations of the key enthalpy changes

3. What does lattice enthalpy tell us? 4. Key features of a Born–Haber cycle 5. How to construct a Born–Haber cycle 6. Constructing a Born–Haber cycle 7. Calculating enthalpy values from a Born–

Haber cycle

● Ionic bonding ● Enthalpy ● Exothermic and endothermic reactions ● Hess’s law ● Lattice enthalpy ● Enthalpy equations

● 5.2.1 ● 5.2.2 ● 5.2.3

● 5.2.1 ● 5.2.2-3


Students should be able to demonstrate and apply their knowledge and understanding of: ● the term lattice enthalpy (formation of one

mole of ionic lattice from gaseous ions, ∆ LEH ) and use it as a measure of the strength of ionic bonding in a giant ionic lattice

● the use of the lattice enthalpy of a simple ionic solid, such as, NaCl and MgCl2, and relevant energy terms for the construction of Born–Haber cycles

● the use of the lattice enthalpy of a simple ionic solid, such as NaCl, MgCl2, and relevant energy terms, for:

(i) the construction of Born–Haber cycles

(ii) related calculations.

● 5.2.1 (a), (b)(i) ● 5.2.1 (b)



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1. The processes that take place when solids dissolve

2. How enthalpy change of lattice breakdown compares and contrasts to lattice enthalpy

3. The factors that affect the size of lattice enthalpy

4. What happens during hydration 5. The factors that affect the size of the

enthalpy of hydration 6. Calculating lattice enthalpy from enthalpy of

solution and hydration 7. Explanation of entropy and standard

entropy 8. Factors affecting entropy, such as

temperature, dissolution and the number of gas molecules

9. Explanation of how to calculate entropy changes for reactions given the relevant entropies of the reactants and the products

● Enthalpy ● Hess’s law ● Born–Haber cycles

● 5.2.4 ● 5.2.5

● 5.2.4 ● 5.2.5


Students should be able to demonstrate and apply their knowledge and understanding of: ● the explanation and use of the terms:

(i) enthalpy change of solution (dissolving of one mole of solute, ∆sol H)

(ii) enthalpy change of hydration (dissolving of one mole of gaseous ions in water, ∆hydH)

● the use of the enthalpy change of solution of a simple ionic solid, such as NaCl, MgCl2, and relevant energy terms (enthalpy change of hydration and lattice enthalpy) for:

(i) the construction of enthalpy cycles

(ii) related calculations ● the qualitative explanation of the effect of

ionic charge and ionic radius on the exothermic value of a lattice enthalpy and enthalpy change of hydration

● the explanation that entropy is a measure of the dispersal of energy in a system, which is higher, the more disordered a system is

● 5.2.1 (c)–(e) ● 5.2.2 (a)–(c)

● Practical 5: Measure an enthalpy change of solution



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● the explanation of the difference in magnitudes of the entropy of a system:

(i) of solids, liquids and gases

(ii) of a reaction in which there is a change in the number of gaseous molecules

• calculation of the entropy change of a system, ∆S, and related quantities for a reaction given the entropies of the reactants and products.



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1. Explanation of how reactions can take place spontaneously

2. Calculating a total change in entropy 3. Factors to consider when determining if

reactions will be spontaneous or not 4. Gibbs free energy and the Gibbs equation 5. Predicting the feasibility of a reaction using

the Gibbs equation 6. Explanation of how to calculate free energy

for some reactions at given temperatures 7. Explanation of redox reactions (OIL RIG) –

reducing agents and oxidation agents 8. Half-equations and how they are

constructed 9. The steps involved in writing balanced

redox equations 10. How oxidation numbers are used to balance

redox equations

● Enthalpy ● Entropy ● Redox ● Oxidation states

● 5.2.6 ● 5.2.7

● 5.2.6 ● 5.2.7


Students should be able to demonstrate and apply their knowledge and understanding of: ● the explanation that the feasibility of a

process depends upon the entropy change and temperature in the system, T∆S, and the enthalpy change of the system, ∆H

● the explanation, and related calculations, of the free energy change, ∆G, as ∆G = ∆H − T∆S (the Gibbs equation) and that a process is feasible when ∆G has a negative value

● the limitations of predictions made using ∆G values about feasibility, in terms of kinetics

● the terms oxidising agent and reducing agent

● the construction of redox equations using half-equations and oxidation numbers

● the interpretation and predictions of reactions involving electron transfer.

● 5.2.2 (d)–(f) ● 5.2.3 (a)–(c)



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1. Recall the steps involved in a successful titration

2. Discuss the use of manganese(VII) and iodine as self-indicating chemicals

3. Titration calculations 4. Explanation of what a half-cell is 5. Examples of half-cells, such as, the

hydrogen half-cell and metal half-cells 6. How to measure the standard electrode

potential for a half-cell 7. How to use the electrochemical series

● Titration ● Redox ● Mole calculations ● Half-equations

● 5.2.8 ● 5.2.9

● 5.2.8 ● 5.2.9


Students should be able to demonstrate and apply their knowledge and understanding of: ● the techniques and procedures used when

carrying out redox titrations, including those involving Fe2+ / MnO4

− and I2 / S2O32−

● structured and non-structured titration calculations, based on experimental results of redox titrations involving:

(i) Fe2+ and MnO4−; I2 and

S2O32−

(ii) non-familiar redox systems

● use of the term standard electrode (redox) potential, Eɵ, including its measurement using a hydrogen electrode

● the techniques and procedures used for the measurement of cell potentials of:

(i) metals or non-metals in contact with their ions in aqueous solution

(ii) ions of the same element in different oxidation states in contact with a platinum electrode.

● 5.2.3 (d)–(e) ● 5.2.3 (f)–(g)

● Practical 1: Construct electrochemical cells and measure electrode potentials

● Practical 2: Find the amount of iron in an iron tablet using redox titration



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1. Using half-cells to make an electrochemical cell

2. How to calculate a standard cell potential 3. How to predict if a reaction will happen

using standard cell potentials 4. Why we may not always be able to predict

reaction feasibility 5. Using electrochemical cells as batteries and

fuel cells

● Redox ● Electrode potentials

● 5.2.10 ● 5.2.10


Students should be able to demonstrate and apply their knowledge and understanding of: ● the calculation of a standard cell potential

by combining two standard electrode potentials

● the prediction of the feasibility of a reaction using standard cell potentials and the limitations of such predictions in terms of kinetics and concentration

● the application of principles of electrode potentials to modern storage cells

● the explanation that a fuel cell transfers the energy from the reaction of a fuel with oxygen to a voltage and the changes that take place at each electrode.

● 5.2.1 (h)–(k)



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1. Introduction to transition metals 2. How to write the electron configuration for

the d-block elements of Period 4 3. Using energy-level diagrams to explain how

transition metals react 4. Properties of the transition metals, including

physical and chemical properties 5. Disproportionation reactions involving

transition metals 6. Why transition metals have variable

oxidation states 7. How transition metals are used as catalysts

● Electron shells ● Energy-level diagrams ● Electron configuration ● The periodic table ● Catalysts ● Oxidation numbers and states ● Redox

● 5.3.1 ● 5.3.2

● 5.3.1 ● 5.3.2


Students should be able to demonstrate and apply their knowledge and understanding of: ● the electron configurations of atoms and

ions of the d-block elements of Period 4 (Sc–Zn), given the atomic number and charge

● the elements Ti–Cu as transition elements, i.e. d-block elements that have ions with an incomplete d-sub-shell

● illustration, using at least two transition elements, of:

(i) the existence of more than one oxidation state for each element in its compounds

(ii) the formation of coloured ions

(iii) the catalytic behaviour of the elements and their compounds and their importance in the industrial manufacture of chemicals.

● 5.3.1 (a), (b) ● 5.3.1 (c)



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1. Explain what is meant by the terms complex ion, coordination number and ligand

2. Identifying the charge on complex ions 3. Working out the oxidation number and

charge of complex ions 4. Shapes of complex ions 5. Explanation of what stereoisomers are,

including both cis–trans and optical isomerism

6. Explanation of cis–trans isomerism in octahedral and square planar complex ions

7. Using transition metal complexes in medicine

8. Discussion of bidentate and multidentate ligands and cis–trans isomerism

9. Discussion of bidentate and multidendate ligands and optical isomerism

● Lewis acids ● Bonding ● Shapes of molecules and ions ● Isomerisation

● 5.3.3 ● 5.3.4

● 5.3.3 ● 5.3.4


Students should be able to demonstrate and apply their knowledge and understanding of: ● the term ligand in terms of coordinate

(dative covalent) bonding to a metal ion or metal, including bidentate ligands

● the terms complex ion and coordination number and examples of complexes with:

(i) six-fold coordination with an octahedral shape

(ii) four-fold coordination with either a planar or tetrahedral shape

● the types of stereoisomerism shown by metal complexes, including those associated with bidentate and multidentate ligands:

(i) cis–trans isomerism e.g. Pt(NH3)2Cl2

(ii) optical isomerism e.g. [Ni(NH2CH2CH2NH2)3]2+

• the use of cis-platin as an anti-cancer drug and its action by binding to DNA preventing cell division.

● 5.3.1 (d), (e) ● 5.3.1 (f), (g)



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1. Explanation of what is meant by ligand substitution

2. Demonstration of the reactions of hydrated copper(II) ions with ammonia and with chloride ions

3. The colours produced by precipitation reactions of given transition metal ions

4. Discussion about haemoglobin and its role in carrying oxygen and carbon dioxide.

5. Discussion about the 'silent killer', carbon monoxide

6. Explanation of redox in transition metals, using both oxidation numbers and transfer of electrons

7. Explanation of redox titrations involving potassium manganate(IV), including calculations initially structured and then unstructured

8. Explanation of redox titrations involving potassium dichromate(VI)

9. Explanation of redox titrations involving iodine and sodium thiosulfate, including calculations initially structured and then unstructured

● Ligands ● Transition metals ● Equilibrium and Le Chatelier’s principle ● Entropy ● Ligand substitution ● Redox ● Disproportionation reactions ● Ligand substitution

● 5.3.5 ● 5.3.6 ● 5.3.7

● 5.3.5-6 ● 5.3.7


Students should be able to demonstrate and apply their knowledge and understanding of: ● ligand substitution reactions and the

accompanying colour changes in the formation of:

(i) [Cu(NH3)4(H2O)2]2+ and [CuCl4]2− from [Cu(H2O)6]2+

(ii) [Cr(NH3)6]3+ from [Cr(H2O)6]3+

● the explanation of the biochemical importance of iron in haemoglobin, including ligand substitution involving O2 and CO

● the reactions, including ionic equations, and the accompanying colour changes of aqueous Cu2+, Fe2+, Fe3+, Mn2+ and Cr3+ with aqueous sodium hydroxide and aqueous ammonia, including:

(i) precipitation reactions

● 5.3.1 (h)–(j) ● 5.3.1 (k), (l)

● Practical 2: Find the amount of iron in an iron tablet using redox titration

● Practical 6: Perform ligand substitution reactions



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(ii) complex formation with excess aqueous sodium hydroxide and aqueous ammonia

● redox reactions and accompanying colour changes for:

(i) interconversions between Fe2+ and Fe3+

(ii) interconversions between Cr3+ and Cr2O7

2− (iii) reduction of Cu2+ to Cu+

and disproportionation of Cu+ to Cu2+ and Cu

● interpretation and prediction of unfamiliar reactions including ligand substitution, precipitation and redox.



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1. The formula and structure of benzene 2. The Kekulé model of benzene 3. Evidence for and against the model 4. Benzene’s delocalised electron structure 5. Naming arenes; single- and double-

substituted

● Electron arrangement ● Naming organic molecules ● Delocalisation of electrons

● 6.1.1 ● 6.2.2

● 6.1.1-2


Students should be able to demonstrate and apply their knowledge and understanding of: ● the comparison of the Kekulé model of

benzene with subsequent delocalised models for benzene in terms of p-orbital overlap forming a delocalised π-system

● the experimental evidence for a delocalised, rather than Kekulé, model for benzene in terms of bond lengths, enthalpy change of hydrogenation and resistance to reaction

● the use of IUPAC rules of nomenclature for systematically naming substituted aromatic compounds.

● 6.1.1 (a)–(c)



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1. Description of electrophilic substitution of aromatic compounds

2. Nitration of benzene and the associated mechanism

3. Halogenation of benzene and the associated mechanism

4. Explain why alkenes will undergo halogenation but benzene will not

5. Friedel–Crafts reaction and its mechanism 6. Using acyl chloride as a halogen carrier in

Friedel–Crafts reactions

● Core organic chemistry ● Reactions of alkenes ● Electrophilic addition ● Bonding between carbon

atoms ● Alkenes ● Lewis acids

● 6.1.3 ● 6.1.4

● 6.1.3 ● 6.1.4


Students should be able to demonstrate and apply their knowledge and understanding of: ● the electrophilic substitution of aromatic

compounds with: (i) concentrated nitric acid in the

presence of concentrated sulfuric acid

(ii) a halogen in the presence of a halogen carrier.

● the interpretation of unfamiliar electrophilic substitution reactions of aromatic compounds, including prediction of mechanisms

● the explanation of the relative resistance to bromination of benzene, compared with alkenes, in terms of the delocalised electron density of the π-system in benzene compared with the localised electron density of the π-bond in alkenes

● 6.1.1 (d), (g) ● 6.1.1 (e), (f)



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● the electrophilic substitution of aromatic compounds with a haloalkane or an acyl chloride in the presence of a halogen carrier (Friedel–Crafts reaction) and its importance to synthesis by formation of a C–C bond to an aromatic ring

● the mechanism of electrophilic substitution in arenes for nitration and halogenation.



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1. Phenols and aromatic alcohols 2. The acidic properties of phenol 3. How the structure of phenol affects its

chemical reactivity 4. Reactions of phenol, including bromination

and nitration 5. Factors affecting the position of substitution 6. Carbonyl compounds, including aldehydes

and ketones 7. Oxidation of aldehydes 8. Nucleophilic addition of aldehydes and the

associated mechanism 9. Reactions of aldehydes with:

(i) sodium tetrahydridoborate(III) (ii) hydrogen cyanide

● Alcohols ● Electrophilic addition ● Halogenation of aromatic compounds ● Oxidation of alcohols ● Organic mechanisms ● Naming organic molecules

● 6.1.5 ● 6.1.6 ● 6.1.7

● 6.1.5-6 ● 6.1.7


Students should be able to demonstrate and apply their knowledge and understanding of: ● the weak acidity of phenols shown by the

neutralisation reaction with NaOH but absence of reaction with carbonates

● the relative ease of electrophilic substitution of phenol compared with benzene, in terms of electron pair donation to the π-system from an oxygen p-orbital in phenol

● the electrophilic substitution reactions of phenol:

(i) with bromine to form 2,4,6-tribromophenol

(ii) with dilute nitric acid to form 2-nitrophenol

● the 2- and 4-directing effect of electron-donating groups (–OH, –NH2 ) and the 3-directing effect of electron-withdrawing groups (–NO2 ) in electrophilic substitution of aromatic compounds

● the prediction of substitution products of aromatic compounds by using directing effects and the importance for organic synthesis

● 6.1.1 (h)–(l) ● 6.1.2 (a)–(c)



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● the oxidation of aldehydes using Cr2O72−/H+

(i.e. K2Cr2O7 / H2SO4) to form carboxylic acids

● the nucleophilic addition reactions of carbonyl compounds with:

(i) NaBH4 to form alcohols (ii) HCN [i.e.

NaCN(aq)/H+(aq)] to form hydroxynitriles

● the mechanism for nucleophilic addition reactions of aldehydes and ketones with NaBH4 and HCN.



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1. Identifying aldehydes and ketones from other carbonyl compounds using: (i) 2,4-dintrophenylhydrazine (2,4-DNPH) (ii) melting points of the

2,4-dinitrophenylhydrazone derivative 2. Distinguishing between

aldehydes and ketones using Tollen’s reagent

3. Naming and drawing carboxylic acids 4. The solubility of carboxylic acids 5. The chemical properties of carboxylic acids

as weak acids 6. Reactions of carboxylic acids with:

(i) metals (ii) metal oxides (iii) metal hydroxides (iv) metal carbonates

● Aldehydes and ketones ● Oxidation reactions ● Redox ● Synthetic routes in organic synthesis ● Properties of acids ● Reactions of acids

● 6.1.8 ● 6.1.9

● 6.1.8 ● 6.1.9


Students should be able to demonstrate and apply their knowledge and understanding of: ● the use of 2,4-dinitrophenylhydrazine to:

(i) detect the presence of a carbonyl group in an organic compound

(ii) identify a carbonyl compound from the melting point of the derivative

● the use of Tollens’ reagent (ammoniacal silver nitrate) to:

(i) detect the presence of an aldehyde group

(ii) distinguish between aldehydes and ketones, explained in terms of the oxidation of aldehydes to carboxylic acids with reduction of silver ions to silver

● 6.1.2 (d), (e) ● 6.1.3 (a), (b)

● Practical 7: Identify an unknown carbonyl compound using 2,4-DNPH

● Practical 14: Reactions of carboxylic acids



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● the explanation of the water solubility of carboxylic acids in terms of hydrogen bonding

● reactions in aqueous conditions of carboxylic acids with metals and bases (including carbonates, metal oxides and alkalis).



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1. Drawing and naming esters 2. Esterification using carboxylic acids 3. Esterification using acid anhydrides 4. Hydrolysis of esters in:

(i) acidic conditions (ii) alkaline conditions

5. Drawing and naming acyl chlorides 6. Preparing acyl chlorides in the laboratory 7. Using acyl chlorides to synthesise organic

molecules: (i) carboxylic acid (ii) primary amines (iii) secondary amines

● IUPAC naming rules ● Carboxylic acids ● Rates of reaction ● Organic synthesis

● 6.1.10 ● 6.1.11

● 6.1.10 ● 6.1.11


Students should be able to demonstrate and apply their knowledge and understanding of: ● esterification of:

(i) carboxylic acids using alcohols in the presence of an acid catalyst (e.g. concentrated H2SO4)

(ii) acid anhydrides using alcohols ● hydrolysis of esters:

(i) in hot aqueous acid to form carboxylic acids and alcohols

(ii) in hot aqueous alkali to form carboxylate salts and alcohols

● the formation of acyl chlorides from carboxylic acids using SOCl2

● the uses of acyl chlorides in synthesis in the formation of esters, carboxylic acids and primary and secondary amides.

● 6.1.3 (c), (d) ● 6.1.3 (e), (f)

● Practical 10: Ester hydrolysis



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1. The structural formulae of some simple amines

2. Defining a base as a proton acceptor 3. Identifying that amines are bases because

nitrogen’s lone pair of electrons can form a dative covalent bond

4. Examples of amines reacting with acids to form salts

5. Preparing aliphatic amines from halogenoalkanes

6. Preparing phenylamine by the reduction of nitrobenzene

7. The general formula and structure of an α-amino acid

8. Explanation of zwitterions and their formation

9. Isoelectric points of zwitterions 10. Reactions of amino acids associated with

the carboxylic acid group 11. Reactions of amino acids associated with

the amine group

● Naming of alcohols ● Acids and bases ● Equilibria ● Carboxylic acids ● Amines

● 6.2.1 ● 6.2.2

● 6.2.1 ● 6.2.2


Students should be able to demonstrate and apply their knowledge and understanding of: ● the basicity of amines in terms of proton

acceptance by the nitrogen lone pair of electrons and the reactions of amines with dilute acids, e.g. HCl(aq), to form salts

● the preparation of: (i) aliphatic amines by substitution

of haloalkanes using excess ethanolic ammonia and amines

(ii) aromatic amines by reduction of nitroarenes using tin and concentrated hydrochloric acid

● the general formula for an α-amino acid as RCH(NH2)COOH

● the following reactions of amino acids: (i) reaction of the

carboxylic acid group (COOH) with alkalis and in the formation of esters

(ii) reaction of the amine group (NH2) with acids.

● 6.2.1 (a), (b) ● 6.2.2 (a)



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1. The amide functional group 2. Structures of amides and classification into

primary, secondary and tertiary 3. Naming primary amides 4. Naming secondary amides and N-

substitution 5. Recall E/Z cis–trans isomerism 6. Optical isomers (enantiomers) and their

properties 7. Explanation of chirality and chiral carbons

● IUPAC naming rules ● Primary, secondary and tertiary structures ● Acyl structures ● Isomerisation

● 6.2.3 ● 6.2.4

● 6.2.3 ● 6.2.4


Students should be able to demonstrate and apply their knowledge and understanding of: ● the structures of primary and secondary

amides ● optical isomerism (an example of

stereoisomerism, in terms of non-superimposable mirror images about a chiral centre)

● the identification of chiral centres in a molecule of any organic compound.

● 6.2.2 (b) ● 6.2.2 (c), (d)



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1. Revision of addition polymers 2. Introduction to condensation polymers 3. Classification of polymers and drawing their

repeat units 4. Explanation of the formation of polyesters 5. Explanation of the formation of polyamides 6. Explanation of hydrolysis, with respect to

polyesters and polyamides 7. Acid and alkali hydrolysis of polyesters 8. Acid and alkali hydrolysis of polyamides

● Polymers ● Addition polymerisation ● Alkenes ● Polymerisation ● Condensation polymers

● 6.2.5 ● 6.2.6

● 6.2.5 ● 6.2.6


Students should be able to demonstrate and apply their knowledge and understanding of: ● condensation polymerisation to form:

(i) polyesters (ii) polyamides

● prediction from addition polymerisation and condensation polymerisation of:

(i) the repeat unit from a given monomer(s)

(ii) the monomer(s) required for a given section of a polymer molecule

(iii) the type of polymerisation ● the acid and base hydrolysis of:

(i) the ester groups in polyesters (ii) the amide groups in polyamides.

● 6.2.3 (a), (c) ● 6.2.3 (b)



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1. Explanation of how carbon–carbon chain length can be increased using the cyanide nucleophile

2. Mechanism for the nucleophilic substitution of a haloalkane using cyanide ions and ethanol

3. Nucleophilic addition of hydrogen cyanide to carbonyl compounds

4. Mechanism for the nucleophilic addition of hydrogen cyanide to carbonyl compounds

5. How optical isomers can be produced by nucleophilic addition

6. Reduction of nitriles 7. Hydrolysis of nitriles by an acid and the

mechanism for the reaction

● Organic reactions ● Nucleophilic addition reactions

● 6.2.7 ● 6.2.8

● 6.2.7-8


Students should be able to demonstrate and apply their knowledge and understanding of: ● the use of C–C bond formation in synthesis

to increase the length of a carbon chain ● formation of C–C≡N by reaction of:

(i) haloalkanes with CN− and ethanol, including the nucleophilic substitution mechanism

(ii) carbonyl compounds with HCN, including the nucleophilic addition mechanism

● the reaction of nitriles: (i) by reduction (e.g. with H2/Ni) to

form amines (ii) by acid hydrolysis to form

carboxylic acids.

● 6.2.4 (a), (b)



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1. The stability of aromatic compounds due to the delocalisation of the electrons in the carbon ring

2. Substitution reactions in aromatic compounds: (i) Alkylation (ii) Acylation

3. How to use a selection of apparatus for: (i) distillation (ii) reflux (iii) purification (iv) recrystallisation (v) melting point measurement

4. Identifying functional groups 5. Recalling types of reactions,

conditions and equations for aliphatic reactions

6. Producing reaction pathways to synthesise organic molecules

● Lewis acids ● Friedel−Crafts reactions ● Substitution reactions

● 6.2.9 ● 6.2.10 ● 6.2.11

● 6.2.9 ● 6.2.10-11


Students should be able to demonstrate and apply their knowledge and understanding of: ● the formation of a substituted aromatic C–C

bond by alkylation (using a haloalkane) and acylation (using an acyl chloride) in the presence of a halogen carrier (Friedel–Crafts reaction)

● the techniques and procedures used for the preparation and purification of organic solids, involving the use of a range of techniques including:

(i) organic preparation (use of Quickfit® apparatus; distillation and heating under reflux)

(ii) purification of an organic solid (filtration under reduced pressure; recrystallisation; measurement of melting points)

● for an organic molecule containing several functional groups:

(i) identification of individual functional groups

(ii) prediction of properties and reactions

● multi-stage synthetic routes for preparing organic compounds.

● 6.2.4 (d) ● 6.2.5 (a) ● 6.2.5 (b), (c)

● Practical 4: Synthesise aspirin from 2-hydroxybenzoic acid



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1. Explanation of the terms ‘chromatography’, ‘mobile phase’ and ‘stationary phase’

2. Description of one-way thin-layer chromatography (TLC)

3. Explanation of Rf values 4. Description of gas chromatography (GC) 5. Explanation of retention times 6. Work through each functional group test:

(i) unsaturation (ii) haloalkanes (iii) carbonyl compounds (iv) aldehydes (v) aliphatic carboxylic acids (vi) phenols (vii) primary alcohols

● Separation techniques ● Organic chemistry

● 6.3.1 ● 6.3.2

● 6.3.1 ● 6.3.2


Students should be able to demonstrate and apply their knowledge and understanding of: ● interpreting one-way TLC chromatograms in

terms of Rf values ● interpreting gas chromatograms in terms of:

(i) retention times (ii) the amounts and proportions of

the components in a mixture ● the qualitative analysis of organic functional

groups on a test-tube scale.

● 6.3.1 (a), (b) ● 6.3.1 (c)

● Practical 9: Identify amino acids using thin layer and paper chromatography

● Practical 15: Analyse organic and inorganic unknowns



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1. Explanation of how nuclear magnetic resonance, NMR, works

2. Use of tetramethylsilane, TMS, as a standard

3. What is meant by ‘chemical shift’ 4. Which solvents can be used 5. Carbon-13 NMR, including the different

types of carbon and chemical shifts 6. Description of how to use a data sheet 7. Interpreting carbon-13 NMR spectra 8. Using carbon-13 NMR to predict structures

● Isotopes ● NMR

● 6.3.3 ● 6.3.4

● 6.3.3 ● 6.3.4


Students should be able to apply their knowledge and understanding of: ● the use of tetramethylsilane, TMS, as the

standard for chemical shift measurements ● the need for deuterated solvents, e.g.

CDCl3, when running a nuclear magnetic resonance, NMR, spectrum

● Students should be able to demonstrate and apply their knowledge and understanding of:

● the analysis of a carbon-13 NMR spectrum of an organic molecule to make predictions about:

(i) the number of carbon environments in the molecule

(ii) the different types of carbon environment present, from chemical shift values

(iii) possible structures for the molecule

● the prediction of a carbon-13 or proton NMR spectrum for a given molecule.

● 6.3.2 (d) (i), (ii) ● 6.3.2 (a), (c)



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1. Proton NMR, including different types of proton, relative peak areas and chemical shifts

2. Spin–spin coupling 3. Using proton NMR to make predictions

about structures 4. Predicting proton NMR for given molecules 5. Introduction to the appearance of the peaks

in –OH and –NH spectra 6. Using D2O in NMR evaluation 7. Production of broad singlet peaks from –OH

and –NH protons 8. Practice questions 9. Revision of the use of mass spectra 10. Revision of the use of infrared spectra 11. Discussion of the information that can be

gathered from all three spectra 12. Practice questions based on combined

techniques

● NMR spectroscopy ● Organic identification ● Carbon-13 NMR ● Proton NMR ● Mass spectrometry ● Infrared spectra ● NMR spectra ● Identification of anions and cations

● 6.3.5 ● 6.3.6 ● 6.3.7 ● 6.3.8

● 6.3.5 ● 6.3.6 ● 6.3.7-8


Students should be able to demonstrate and apply their knowledge and understanding of: ● analysis of a high-resolution proton NMR

spectrum of an organic molecule to make predictions about:

(i) the number of proton environments in the molecule

(ii) the different types of proton environment present from chemical shift values

(iii) the relative numbers of each type of proton present from relative peak areas, using integration traces or ratio numbers, when required

(iv) the number of non-equivalent protons adjacent to a given proton from the spin–spin splitting pattern, using the n + 1 rule

(v) possible structures for the molecule

● 6.3.2 (b) ● 6.3.2 (d) (iii) ● 6.3.2 (a) ● 6.3.2 (e)

● Practical 15: Analyse organic and inorganic unknowns



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● prediction of a proton NMR spectrum for a given molecule

● the identification of O–H and N–H protons by proton exchange using D2O

● deduction of the structures of organic compounds from different analytical data including:

(i) elemental analysis (ii) mass spectra (iii) infrared spectra (iv) NMR spectra

● qualitative analysis of ions on a test-tube scale; processes and techniques needed to identify the following ions in an unknown compound:

(i) anions: CO32−, Cl−, Br−,

I−, SO42−

(ii) cations: NH4+, Cu2+,

Fe2+, Fe3+, Mn2+, Cr3+.

week 1 links to prior learning student book links teaching plan … · 2019-07-19 · clock...

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