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GCE Chemistry

Contents Page

Unit AS 1: Basic Concepts in Physical and Organic Chemistry 5

Unit AS 2: Further Physical and Inorganic Chemistry and Introduction to Organic Chemistry

46

Unit A2 1: Further Physical and Organic Chemistry 91

Unit A2 2: Analytical, Transition Metals Electrochemistry and Organic Nitrogen

133

CCEA Exemplar Scheme of Work: GCE Chemistry

Introduction

CCEA has developed new GCE Chemistry specifications for first teaching from September 2016. This scheme of work has been designed to support you in introducing the new specification.

The scheme of work provides suggestions for organising and supporting students’ learning activities. It is intended to assist you in developing your own scheme of work and should not be considered as being prescriptive or exhaustive.

Please remember that assessment is based on the specification which details the knowledge, understanding and skills that students need to acquire during the course. The scheme of work should therefore be used in conjunction with the specification.

Published resources and web references included in the scheme of work have been checked and were correct at the time of writing. You should check with publishers and websites for the latest versions and updates. CCEA accepts no responsibility for the content of third party publications or websites referred to within this scheme of work.

A Microsoft Word version of this scheme of work is available on the subject microsite on the CCEA website (www.ccea.org.uk/microsites). You will be able to use it as a foundation for developing your own scheme of work which will be matched to your teaching and learning environment and the needs of your students.

I hope you find this support useful in your teaching.

Best wishes

Elaine LennoxSubject OfficerChemistry

E-mail [email protected] 028 9026 1200 (2320)

Updated: August 2016 1

http://www.ccea.org.uk/microsites


CCEA Exemplar Scheme of Work:

GCE Chemistry



Unit AS 1:Basic Concepts in Physical

and Organic Chemistry

Updated: August 20165


Specification: GCE Chemistry

Unit AS 1: Basic Concepts in Physical and Organic Chemistry

Foreword: The scheme of work has been developed in an editable version to enable teachers to adapt to their own teaching order and style. Whilst non-CCEA specific resources have been highlighted along with CCEA past paper references and the new CCEA Factfiles, other resources are available on the Chemistry microsite www.ccea.org.uk/chemistry to support the teaching of the GCE Chemistry specification. Most of the additional resources are practical based and as the practicals conducted in schools and the resources available differ, they have not been directly referenced in the scheme of work. Teachers will be able to add these resources to the scheme of work as a matter of choice. The additional support resources available are listed below:

E-book for AS 1, AS 2, A2 1 and A2 2; A practical manual for AS 3 and A2 3; A Clarification of Terms document; An Acceptable Colours document; A summary of uniform mark grade boundaries; The assessment Objective grids which align with the Specimen Assessment Materials; Videos of some of the practicals listed in the specification; and Higher Education and Industry videos highlighting third level and employment opportunities in the field of

Chemistry.


http://www.ccea.org.uk/chemistry


Specification Reference

Learning Outcomes

Suggested Teaching Activities

Practical Links

Resources/Assessment Opportunities

Suggested Time

1.1 Formulae, equations and amounts of substance

Students should be able to:

1.1 Assessed using practice questions, past paper questions and end of unit test

Past paper AS 1January 2011 Q 8, 12 a, b, 13 c (i), 16 dJune 2011 Q 6, 12January 2012 Q 13 a, b, 15 a–c,June 2012 Q 12 a–c

CCEA Factfile

www.docbrown.info/page01/ElCpdMix/EleCmdMix3.htm

www.rsc.org/learn-chemistry/wiki/Category:Balancing_equations

4 hours

1.1.1 write formulae of ionic compounds by predicting the ionic charge from the position of an element in the Periodic Table and by recalling the following molecular ions and their formulae (the systematic names for these ions are

Revise the principles of writing formulae. Students can list the formulae they know and analyse the charges on the ions and the balance of them. Using the ions in the specification create a sheet with all possible formulae from


http://www.rsc.org/learn-chemistry/wiki/Category:Balancing_equations



http://www.docbrown.info/page01/ElCpdMix/EleCmdMix3.htm




not given) − see Section 1.7.6:

e.g. K+, Mg2+ and Al3+ and the anions in the table together with NH4

+.




Learning Outcomes


Practical Links


Suggested Time

1.1 Formulae, equations and amounts of substance (cont.)


1.1.1 (cont.) sulfate, SO42-;

sulfite, SO32-;

thiosulfate, S2O3

2-; hydrogensulfat

e, HSO4-; hydrogencarbo

nate, HCO3-;

carbonate, CO3

2-; nitrate, NO3-; nitrite, NO2-; phosphate,

PO43-;

chlorate, ClO3-;

hypochlorite, ClO-;

hydroxide, HO-; dichromate,

Cr2O72-;

chromate, CrO4

2-;

The compounds can be named with or without oxidation numbers but alternatively based on the charge of the cation



permanganate, MnO4

-; and ammonium,

NH4+.




Learning Outcomes


Practical Links


Suggested Time



1.1.2 write and balance equations for unfamiliar reactions given appropriate information;

Revise writing balanced equations. Include examples of unfamiliar reactions with unfamiliar compounds. Include complicated examples with large numbers e.g. FeS2 with O2

1.1.3 write balanced equations (full and ionic) including state symbols, for all reactions studied;

Revise writing ionic equations. Can be linked to section 1.10.4 which includes gases and precipitates




Learning Outcomes


Practical Links


Suggested Time



1.1.4 define and understand the terms Avogadro’s constant, the mole and molar mass;

Introduce the idea of the mole and Avogadro’s constant. Students record and learn definitions. Students calculate the molar mass of a variety of substances e.g. ionic and covalent and the number of atoms/ions in the substances

1.1.5 use Avogadro’s constant in calculations;

Link this section to the definition in 1.1.4. Discuss further examples



e.g. the difference between the number of atoms in a mole and number of molecules e.g. diatomic molecules etc.




Learning Outcomes


Practical Links


Suggested Time



1.1.6 calculate reacting masses of substances including examples in which some reactants are in excess; and

Revise calculations involving reacting masses; introduce ‘limiting reagent’ and one or more reactants in excess. Calculate the numbers of moles of products and reactants from given data and relate to the balanced equation




Learning Outcomes


Practical Links


Suggested Time



1.1.7 understand the terms anhydrous, hydrated and water of crystallisation and be able to calculate the moles of water of crystallisation present from percentage composition, mass composition or experimental data.

Students should record definitions and practice both structured and non-structured calculations

Heat a variety of hydrated salts. Describe and explain what happens. Write equations for the decompositions. Add water to the anhydrous salts and show the reactions are reversible

Determine the formula of a hydrated compound by weighing and heating a hydrated salt to constant mass




Learning Outcomes


Practical Links


Suggested Time

1.2 Atomic structure


Revision From GCSE

1.2.8 – 1.2.14 Assessed using practice questions, past paper questions and end of unit test

Past paper AS 1January 2011 Q 3June 2011 Q 1, 2, 3, 7, 8, 9,15 a-cJanuary 2012 Q 7June 2012 Q 3, 9, 11 b (ii), c

CCEA Factfile

www.chemguide.co.uk/atoms/properties/elstructs.html

www.learner.org/interactives/periodic/elementary3.html

www.chemguide.co.uk/atoms/properties/atomorbs.html

12 hours

1.2.1 – 1.2.7 understand that an orbital is a region within an atom that can hold up to two electrons with opposite spins and describe the shape of s and p-orbitals;

Revise electronic structure and shells and its links to the Periodic Table

Introduce the idea of orbitals and the notion of sub-shells to lead into s orbitals and p orbitals. This is often done via ionisation energies e.g.


http://www.chemguide.co.uk/atoms/properties/atomorbs.html



http://www.learner.org/interactives/periodic/elementary3.html



http://www.chemguide.co.uk/atoms/properties/elstructs.html




those of sodium




Learning Outcomes


Practical Links


Suggested Time

1.2 Atomic structure (cont.)


1.2.8 deduce the electronic configuration of atoms and ions up to krypton in terms of shells and sub-shells using the building up principle(s, p and d notation and electrons in boxes notation);

Students record the definition of an orbital and the shapes of s and p orbitals. 3D diagrams of orbitals can be drawn.Explain the filling of shells, the number of sub-shells and the relative energy levels of sub-shells.Explain the reasons for the electrons in boxes notation.Students write electronic configurations up to krypton. Discuss the



differences between an atom and its ions. Explain why atoms and ions have certain structures based on stabilities of filled and half-filled sub-shells




Learning Outcomes


Practical Links


Suggested Time



1.2.10 classify an element as belonging to the s, p, d or f block according to its position in the Periodic Table;

Shade the Periodic Table into different coloured blocks. Write the full and outer electron structures of the elements to show why they fit into the blocks

www.showme.com/sh/?h=IQg5WHg

1.2.11 define and write equations for the first and successive ionisation energies of an element in terms (IE) of one mole of gaseous atoms and ions;

Students record definitions and write equations for IE for metals and non-metals using state symbols.

Advise students to take care with electron balancing on the LHS and RHS of

www.bbc.co.uk/bitesize/higher/chemistry/energy/patterns/revision/3/


http://www.bbc.co.uk/bitesize/higher/chemistry/energy/patterns/revision/3/



http://www.showme.com/sh/?h=IQg5WHg

http://www.showme.com/sh/?h=IQg5WHg


the equation is mentioned




Learning Outcomes


Practical Links


Suggested Time



1.2.12 understand that successive ionisation energies can be used to predict the group of an element, and that graphs of successive ionisation energies against number of electrons removed, for an element, give evidence for the existence of shells;

Introduce graphs of successive IEs and label graphs to show the sub-shells and use them to predict the Group number.

Explain the use of the logarithm of IE in order to use a reasonable scale.

Explain the determination of an element from successive IEs




Learning Outcomes


Practical Links


Suggested Time



1.2.13 explain the trend in the first ionisation energies of atoms down Groups, and across Periods in terms of nuclear charge, distance of outermost electron from the nucleus, shielding and stability of filled and half-filled sub-shells; and

Students discuss their ideas on trends of IE down Groups and across Periods. Students explain IE, based on nuclear charge/effective nuclear charge, distance of outermost electron from the nucleus and the shielding and stability of filled and half-filled sub-shells

http://legacy.chemgym.net/as_a2/topics/periodic_table/quiz_4.html

1.2.14 understand that graphs of first ionisation energies of elements up to

Students plot graphs of 1st IE up to krypton and explain the presence of shells






krypton provide evidence for the existence of shells and sub-shells.

and sub-shells using the terms used in 1.2.13




Learning Outcomes


Practical Links


Suggested Time

1.3 Bonding Students should be able to:


6 hours

1.3.1 understand that ionic bonding is the electrostatic attraction between oppositely charged ions formed by electron transfer;

Record the definition and give examples outlined in 1.3.2 and quiz students from a list of many compounds to ask which are ionic and which are covalent or simply classify into substance types.

Revise ionic

Past Paper AS 1January 2011 12 c, dJune 2011 Q 10, 15 dJanuary 2012 Q 14 dJune 2012 Q 10, 11 b (iii)

CCEA Factfile

www.s-cool.co.uk/a-level/chemistry/atomic-structure/revise-it/chemical-bonding


http://www.s-cool.co.uk/a-level/chemistry/atomic-structure/revise-it/chemical-bonding




forces reattractive and repulsive forces and the effect of the distance between them




Learning Outcomes


Practical Links


Suggested Time

1.3 Bonding (cont.)


1.3.2 construct dot and cross diagrams for ionically bonded compounds, for example elements in Groups I, II, VI and VII, the ions of which have a noble gas structure;

The noble gas structure can be related to the octet rule in 1.3.7. Ask students to explain the formation of ionic compounds from atoms and molecules of reactants and ask them to construct dot and cross diagrams of reactants and products

1.3.3 know that a covalent bond is the electrostatic attraction between a shared pair of electrons and the nuclei of

Students record the definition. Explain that the attraction for the electrons is greater than the repulsion



the bonded atoms; between the nuclei

1.3.4 define the term lone pair;

Students record the definition. Examine a variety of molecules to determine the number of lone pairs e.g. Cl2, H2O, NH3. Draw the dot and cross diagrams




Learning Outcomes


Practical Links


Suggested Time

1.3 Bonding (cont.)


1.3.5 describe the coordinate bond as a shared pair of electrons with both electrons supplied by one atom, for example the ammonium ion, NH4

+;

Introduce the idea of lone pairs being able to make a coordinate bond i.e.NH3 + H+ → NH4

+

Draw the dot cross diagram for the reaction. Write the equation for the reaction of ammonia with hydrochloricacid

1.3.6 construct ‘dot and cross’ diagrams for molecules and ions with coordinate, single, double and triple covalent bonds;

Students construct examples of dot and cross diagrams of molecules and ions with coordinate,



single, double and triple covalent bonds




Learning Outcomes


Practical Links


Suggested Time

1.3 Bonding (cont.)


1.3.7 define the octet rule and state its limitations, for example in BeCl2 and BF3;

Students record the definition of octet rule. Students construct dot and cross diagrams for BeCl2, BF3 and SF6. Explain the octet rule and state that not all molecules/ions follow the rule

1.3.8 define the term electronegativity and explain the trend in the electronegativity of elements across Periods and down Groups;

Students record the definition of electronegativity. Using HCl and Cl2 as examples introduce the concept of unequal sharing of electrons between atoms. Students deduce how

www.chemguide.co.uk/atoms/bonding/electroneg.html


http://www.chemguide.co.uk/atoms/bonding/electroneg.html




electronegativity varies across Periods and down Groups




Learning Outcomes


Practical Links


Suggested Time

1.3 Bonding (cont.)


1.3.9 explain that bond polarity arises when covalently bonded atoms have different electronegativities and use partial charges to show that a bond is polar; and

Using HCl and Cl2 students predict the effect of EN on the charges in the molecule. Introduce the term bond polarity and the symbol δ to denote partial charge. Record Pauling EN values and determine which bonds are least and most polar

Use the deflection of a stream of liquid from a burette to indicate polarity or lack of polarity within a molecule

1.3.10 understand that metallic bonding is the attraction between positive ions and delocalised electrons in a

Students record the definition of metallic bonding and draw a labelled diagram. Compare the diagrams for Na,



lattice. Mg and Al. Explain the different electrical conductivities of metals in Groups I, II and III




Learning Outcomes


Practical Links


Suggested Time

1.4 Intermolecularforces



2 hours

1.4.1 describe intermolecular forces as van der Waals’ forces (viewed as attractions between induced dipoles), permanent dipole-dipole attractions and hydrogen bonding (between molecules

Using xenon as an example students explain van der Waals’ forces and record the definition and a diagram of induced forces

Using HCl as an example, revise bond polarity and explain

Past paper AS 1January 2011 Q 10, 15June 2011 Q 16 (e)January 2012 Q 1, 2, 15 f (ii),15 f (iii)June 2012 Q 5, 15

CCEA Factfile

www.bbc.co.uk/bitesize/higher/chemistry/energy/bsp/revision/2/

www.chemguide.co.uk/atoms/bonding/hbond.html

www.wiley.com/college/boyer/0470003790/reviews/pH/ph_water.htm


http://www.wiley.com/college/boyer/0470003790/reviews/pH/ph_water.htm



http://www.chemguide.co.uk/atoms/bonding/hbond.html



http://www.bbc.co.uk/bitesize/higher/chemistry/energy/bsp/revision/2/




containing N, O or F and the H atom of -OH, -NH or HF); and

(permanent) dipole-dipole attractions.




Learning Outcomes


Practical Links


Suggested Time

1.4 IntermolecularForces (cont.)


1.4.2 understand the relationships between these attractive forces and physical properties, such as melting point, boiling point and solubility of covalent molecular substances − see Section 1.5.2.

Students record the diagram using δ symbols. Using H2O, NH3 and HF explain hydrogen bonding; students record diagrams showing lone pair interaction with a hydrogen bond symbol

Discussion on the relative strengths of intermolecular forces. Generally, H-bonding > dipole-dipole forces > van der Waals’ forces. Explain



in terms of forces what happens when substances melt, boil and dissolve




Learning Outcomes


Practical Links


Suggested Time

1.5 Structure Students should be able to:


Past paper AS 1

3 hours

1.5.1 describe the following types of structure: the giant ionic

lattice of sodium chloride;

the metallic lattice of metals;

the giant covalent structures of graphite and diamond; and

molecular

Students discuss the structures of NaCl, Mg, I2 and diamond and graphite. Students record diagrams of the structures. Explain the bonding within each structure

January 2011 Q 5, 15June 2011 Q 11, 16January 2012 Q 4, 6, 8, 9

CCEA Factfile

Students could create a summary presentation of all structures in 1.5



covalent structures, for example, iodine;




Learning Outcomes


Practical Links


Suggested Time

1.5 Structure (cont.)


1.5.2 explain the characteristic physical properties of these structures including melting and boiling point, hardness (graphite and diamond only) and electrical conductivity in terms of structure and bonding; and

Explain that sodium chloride has strong ionic forces and when molten, or in solution, the ions move and carry charge. Explain that metals have positive charges held together by delocalised electrons which move and carry charge. Explain that diamond has strong bonds and no free electrons. Explain that graphite has strong bonds but weak bonds between layers through which electrons can

Carry out tests of electrical conductivity on solids and liquids and aqueous solutions of ionic and covalent substances



move. Explain that iodine has only van der Waals’ forces and no free electrons




Learning Outcomes


Practical Links


Suggested Time

1.5 Structure (cont.)


1.5.3 explain the trend in melting point across the Period sodium to argon, in terms of structure and bonding.

Graph of melting points of Period III elements versus atomic number. Discussion of variation of melting points based on metallic structure i.e. number of delocalised electrons; then giant covalent; and finally P4, S8, Cl2 and Ar based on van der Waals’ forces




Learning Outcomes


Practical Links


Suggested Time

1.6 Shapes of molecules and ions



Past paper AS 1January 2011 Q 7, 12 dJune 2011 Q 5, 16January 2012 Q 11June 2012 Q 7

CCEA Factfile

www.youtube.com/watch?v=Ml-kdheJHDM

5 hours

1.6.1 understand that the shape of a molecule or ion is determined by the repulsion between the electron pairs surrounding a central atom;

Explain that bonding pairs repel each other to be as far away as possible to minimise repulsion. Use BeCl2 as an example to give a linear shape. Use CH4 to give a tetrahedral shape


http://www.youtube.com/watch?v=Ml-kdheJHDM

http://www.youtube.com/watch?v=Ml-kdheJHDM



Learning Outcomes


Practical Links


Suggested Time

1.6 Shapes of molecules and ions (cont.)


1.6.2 use valence shell electron pair repulsion theory to explain the shapes, and bond angles of molecules and ions with up to six outer pairs of electrons around the central atom to include the following shapes – linear, bent, trigonal planar, tetrahedral, pyramidal, octahedral, square planar, trigonal bipyramid, T-shaped;

Using the following molecules, write the dot and cross electron structure and identify the number of bonding pairs (BP) and lone pairs (LP): CO2 (4BP); H2O (2 BP + 2LP); BF3 (3BP); CCl4 (4BP); NH3 (3BP + 1LP); SF6 (6BP); XeF4 (4BP + 2LP); PCl5 (5BP); ClF3 (3BP + 2LP)Students can make ball and stick models of the molecules or

http://glossary.periodni.com/glossary.php?en=T-shaped+molecular+geometry






use balloons. Record the names of shapes and bond angles for the molecules




Learning Outcomes


Practical Links


Suggested Time

1.6 Shapes of molecules and ions (cont.)


1.6.3 explain the departure of the bond angles in NH3 (107°) and H2O (104.5°) from the predicted tetrahedral (109.5°), in terms of the increasing repulsion between bonding pair-bonding pair, lone pair-bonding pair and lone pair-lone pair electrons; and

Students draw the shapes of NH3, H2O and HF based on a tetrahedron showing the bonding pairs and the lone pairs. Students rationalise and explain in words the reasons for the shapes

1.6.4 understand the difference between polar bonds and polar molecules and be able to use the

Discuss various molecules and explain how to determine whether a molecule is polar

http://preparatorychemistry.com/Bishop_molecular_polarity.htm






shape and dipoles present to predict whether or not a given molecule is polar.

or not. Carbon dioxide has polar bonds but based on symmetry the molecule is non-polar




Learning Outcomes


Practical Links


Suggested Time

1.7 Redox Students should be able to:


Past paper AS 1January 2011 Q 13 a–bJanuary 2012 Q 12 a, b, d (i)June 2012 Q 4 (c)

CCEA Factfile

www.chemguide.co.uk/inorganic/redox/oxidnstates.html

www.chemteam.info/Redox/Redox-Rules.html

4 hours

1.7.1 calculate the oxidation state for an element in a compound or ion including peroxides and metal hydrides;

Use the reaction between Zn and Cu2+ to introduce the idea of oxidation state. Students record rules for calculating oxidation states and practice calculating oxidation states for various elements and ions. Oxygen has an oxidation


http://www.chemteam.info/Redox/Redox-Rules.html

http://www.chemteam.info/Redox/Redox-Rules.html

http://www.chemguide.co.uk/inorganic/redox/oxidnstates.html




number of -1 in peroxides and hydrogen has an oxidation number of -1 in hydrides

1.7.2




Learning Outcomes


Practical Links


Suggested Time

1.7 Redox (cont.)


1.7.3 define the term redox and explain oxidation and reduction in terms of electron transfer and changes in oxidation state;

understand that oxidising agents gain electrons and are reduced, and reducing agents lose electrons and are oxidised;

Students use the reaction between Zn and Cu2+ to label oxidation/reduction changes and oxidising and reducing agents by looking at changes in oxidation state. This is repeated using electron transfer

Students examine and balance equations which show oxidising and reducing agents in terms of electrons added or removed e.g. the equation for the



reduction of MnO4

- ions:

MnO4- + 8H+ +

5e- → Mn2+ + 4H2O

1.7.4




Learning Outcomes


Practical Links


Suggested Time

1.7 Redox (cont.)


1.7.5 define disproportionation and use oxidation numbers to classify a redox reaction as disproportionation;

write half-equations and combine half-equations to give a balanced redox equation; and

Students record the definition of disproportionation. Using the example of Cl2 dissolving in water, develop understanding of disproportionation; students should label reaction with oxidation numbers

Write half-equations e.g.Zn → Zn2+ + 2 e- and Cu2+ + 2 e- → Cu and write a balanced equation. Apply oxidation numbers and electron transfer to show it is



redox

1.7.6 use Roman numerals to indicate the oxidation number when an element has compounds or ions with different oxidation numbers, for example chlorate(I), chlorate(V).

Revise Roman numerals and give further examples e.g. manganate(VII), manganese(II), nitrate (V), nitrate(III)




Learning Outcomes


Practical Links


Suggested Time

1.8 Halogens Students should be able to:


8 hours

1.8.1 recall the colours of the elements and explain the trends within the Group, limited to physical state at room temperature, melting and boiling points;

Students discuss, examine samples, then record colours according to the clarification of terms document and record melting and boiling points. Explain the trends using van der Waals’ forces. Combine

Past paper AS 1January 2011 Q 14June 2011 Q 14January 2012 Q 5, 12 c, d (ii),13 c–f, 15 f (iv)June 2012 Q 14

CCEA Factfile

http://chemstuff.co.uk/2012/12/21/displacement-reactions/






everything in a table. Predict the data for astatine

1.8.2




Learning Outcomes


Practical Links


Suggested Time

1.8 Halogens (cont.)


1.8.3 compare the solubility and colours of the halogens in water and non-aqueous solvents, for example, hexane;

describe the reaction of the halogens with cold dilute and hot concentrated aqueous sodium hydroxide and explain the disproportionation in these reactions;

Students discuss why solubility varies in polar water and non-polar hexane. Students record aqueous and non-aqueous colours according to the clarification of terms document

Students record the equations for the reactions, the colour changes and the changes in oxidation numbers

Determine the solubility of chlorine and iodine in aqueous and non-aqueous solvents

1.8.4 recall the reaction of chlorine with

Students record the reaction and



water to form chloride ions and chlorate(I) ions;

note its use in the sterilisation of water. Chloride ions can be tested for using the test in section 1.10




Learning Outcomes


Practical Links


Suggested Time



1.8.5

1.8.6

describe the trend in oxidising ability of the halogens down the group applied to displacement reactions of the halogens with other halide ions in solution;

understand the reactions of solid halides with concentrated sulfuric and phosphoric acid in relation to the relative reducing ability of the hydrogen halides/halide ions; understand the reactions of solid halides with

Students discuss and record the trend. Revise displacement reactions and carry out the experiments; students record colours as listed in the clarification of terms document

Discuss the differences between the oxidising ability of sulfuric and phosphoric acid.Demonstrate, or students perform, the reactions. Students record observations, list products and

Produce a reactivity order of the halogens using the displacement reactions of the halogens with other halide ions in solution

Carry out the reactions of the halides with concentrated sulfuric and phosphoric acids and perform chemical tests for the products (excluding hydrogen sulfide)

www.youtube.com/watch?v=yOBqSc7dNi0

www.chemguide.co.uk/inorganic/group7/acidityhx.html

www.chemguide.co.uk/inorganic/group7/halideions.html


http://www.chemguide.co.uk/inorganic/group7/halideions.html



http://www.chemguide.co.uk/inorganic/group7/acidityhx.html



http://www.youtube.com/watch?v=yOBqSc7dNi0

http://www.youtube.com/watch?v=yOBqSc7dNi0


concentrated sulfuric and phosphoric acid in relation to the relative reducing ability of the hydrogen halides/halide ions; and

write equations for the reactions. Explanation of the reducing ability of the halide ion




Learning Outcomes


Practical Links


Suggested Time



1.8.7 compare the advantages and disadvantages of adding chlorine or ozone to drinking water.

Students discuss/debate the pros and cons of adding chlorine or ozone to drinking water e.g. costs, reactions, freedom of medication etc.




Learning Outcomes


Practical Links


Suggested Time

1.9 Acid-basetitrations


1.9 Assessed using practice questions, past paper questions and end of unit test. Also, link to practical exam

12 hours

1.9.1 understand the concept of weak and strong acids and bases in terms of dissociation of hydrogen ions and hydroxide ions;

Students revise strong and weak acids and bases and their neutralisation reactions. Students record equations for reactions and transform them to the corresponding ionic equations

Cancellation of

Carry out an acid-base titration to determine the concentration of acid/base, the degree of hydration in a hydrated metal carbonate and the percentage of ethanoic acid in vinegar

See CCEA video

Past Paper AS 1January 2012 Q 15 dJune 2012 Q 8, 12 d, 13

CCEA Factfile



all spectator ions leads to the equationH+ + OH- → H2O




Learning Outcomes


Practical Links


Suggested Time

1.9 Acid-basetitrations (cont.)


1.9.2 understand the techniques and procedures used when experimentally carrying out acid-base titrations involving strong acid/strong base, strong acid/weak base and weak acid/strong base, for example, determination of the degree of hydration in a sample of sodium carbonate, analysis of vinegar;

Demonstrate and explain the full procedure for carrying out titrations. Students record the procedures together with the explanations. They practice a wide variety of acid-base titrations

www.youtube.com/watch?v=9DkB82xLvNE


http://www.youtube.com/watch?v=9DkB82xLvNE

http://www.youtube.com/watch?v=9DkB82xLvNE



Learning Outcomes


Practical Links


Suggested Time



1.9.3 select the correct indicator for each type of titration and recall the colour changes of phenolphthalein and methyl orange at the end point;

Students practice titrations using both indicators and record the colour changes according to the clarification of terms document. A comparison of the two indicators for the same reaction, e.g. a weak acid with a strong base, can be made

www.titrations.info/acid-base-titration-indicators-preparation

1.9.4 identify uncertainties in measurements and calculate the uncertainty when two burette readings are used

Discuss the errors in measurements and practice calculating the error in burette readings before


http://www.titrations.info/acid-base-titration-indicators-preparation




to calculate a titre value;

and after addition




Learning Outcomes


Practical Links


Suggested Time



1.9.5 select appropriate titration data, ignoring outliers, in order to calculate mean titres;

Explain the term “outlier”. Explain the use of a table to present titration data and the correct recording of titres and significant figures. Calculate mean/average titres

1.9.6 calculate concentrations and volumes for titration calculations;

Discuss the use of a rough titration and the need to have two titres within 0.1 cm3. Discuss and practice selecting titres used to calculate mean/average



titre

1.9.7 be familiar with the term molarity, M, and the units of concentration, for example mol dm-3,g dm-3; and

Revise units of concentration and convert from mol dm-3 to g dm-3

and vice versa

Prepare solutions of known concentration




Learning Outcomes


Practical Links


Suggested Time



1.9.8 describe the techniques and procedures used to prepare a standard solution of required concentration.

Students record the procedure and explain the reasons for it having calculated the mass of solute required for a particular molarity required. The use of the volumetric flask is explained and recorded




Learning Outcomes


Practical Links


Suggested Time

1.10 Qualitativetests


1.10 Assessed using practice questions, past paper questions and end of unit test.Also, useful to practice practical exam style questions/lab experiments

6 hours

1.10.1 use a chemical test for the gases H2, O2, Cl2, CO2, HCl and NH3;

Demonstrate the techniques used in the tests especially the variety of ways of testing for carbon dioxide. The clarification of terms document is used to state how the test

Use chemical tests listed in ‘Qualitative tests’ to identify unknown substances

Past paper AS 1January 2011 Q 16 eJune 2011 Q 15 f,January 2012 Q 15 eJune 2012 Q 2, 11 d

CCEA Factfile

www.youtube.com/watch?v=LiAvDpl5aJA

Students could create a summary presentation based on tests in 1.10

www.youtube.com/watch?v=NEUbBAGw14k

Useful for seeing flames,


http://www.youtube.com/watch?v=NEUbBAGw14k

http://www.youtube.com/watch?v=NEUbBAGw14k

http://www.youtube.com/watch?v=LiAvDpl5aJA

http://www.youtube.com/watch?v=LiAvDpl5aJA


results should be described. List the tests in a table. Students practice the tests

but students should use CCEA descriptions for the colours




Learning Outcomes


Practical Links


Suggested Time

1.10 Qualitativetests


1.10.2 understand how to carry out a flame test using nichrome wire;

Demonstrate the flame test procedure. Explain the reasons for the steps. Students record the procedure then practice using the ions stated in 1.10.3

1.10.3 use cation tests including: flame tests to

identify the metal ions Li+, Na+, K+, Ca2+, Ba2+, and Cu2+;

adding sodium hydroxide solution and warming to identify

Using the procedure from 1.10.2 students practise flame tests and record the colours as detailed in the clarification of terms document.Students record the procedure then practise to

www.youtube.com/watch?v=7F4JhrBWdY4


http://www.youtube.com/watch?v=7F4JhrBWdY4

http://www.youtube.com/watch?v=7F4JhrBWdY4


ammonium ion; identify ammonium ion




Learning Outcomes


Practical Links


Suggested Time

1.10 Qualitative tests


1.10.4 use anion tests including: adding barium

chloride solution to identify sulfate ion;

adding acidified silver nitrate solution to distinguish between chloride, bromide and iodide (followed by adding dilute and concentrated ammonia solution);

adding dilute acid to test for carbonate ion,

Students record procedures then practise using the anions listed, paying special attention to the clarification of terms document. A table can be produced for the tests together with the expected results



and identifying the gas produced; and

1.10.5 use starch to identify iodine.

Students record the procedure and practise testing for solid and aqueous iodine using aqueous starch



Unit AS 2Further Physical and

Inorganic Chemistry and an Introduction to Organic

Chemistry




Unit AS 2: Further Physical and Inorganic Chemistry and an Introduction to Organic Chemistry


Learning Outcomes


Practical Links


Suggested Time

2.1 Formulae and amounts of a substance


2.1 Assessed using past paper questions, practice questions and end of unit test

5 hours

2.1.1 define the terms empirical and molecular formula and explain the relationship between them;

Students define empirical and molecular formula and record examples

Past paper AS 2January 2012 Q 10June 2012 Q 8, 11January 2011 Q 3, 6June 2011 Q 15 e

CCEA Factfile

www.bbc.co.uk/schools/gcsebitesize/science/add_aqa_pre_2011/chemcalc/chemcalc_higherrev2.shtml

2.1.2 calculate empirical and molecular formulae using data, given

Students calculate empirical and molecular formulae using

www.youtube.com/watch?v=wnRaBWvhYKY


http://www.youtube.com/watch?v=wnRaBWvhYKY

http://www.youtube.com/watch?v=wnRaBWvhYKY

http://www.bbc.co.uk/schools/gcsebitesize/science/add_aqa_pre_2011/chemcalc/chemcalc_higherrev2.shtml




composition by mass or percentage composition;

composition by mass or percentage composition




Learning Outcomes


Practical Links


Suggested Time

2.1 Formulae and amounts of a substance (cont.)


2.1.3 define molar gas volume and calculate reacting gas volumes from chemical equations;

Discuss why one mole of any gas occupies the same volume at the same temperature and pressure, regardless of RMM. Students record the definition and practise calculating gas volumes from chemical equations

www.bbc.co.uk/bitesize/higher/chemistry/calculations_1/mole/revision/2/

2.1.4 define percentage yield and calculate percentage yields using chemical

Discussion of the importance of percentage yield in industry and in the laboratory.

www.sparknotes.com/chemistry/stoichiometry/realworldreactions/section2.rhtml


http://www.sparknotes.com/chemistry/stoichiometry/realworldreactions/section2.rhtml



http://www.bbc.co.uk/bitesize/higher/chemistry/calculations_1/mole/revision/2/




equations and experimental data;

Students record the definition and practise calculating percentage yields using chemical equations and experimental data




Learning Outcomes


Practical Links


Suggested Time

2.1 Formulae and amounts of a substance (cont.)


2.1.5 use a percentage yield to determine the amount of reagent(s) needed for a reaction; and

Further practice of using percentage yield; this could be linked to planning an experiment

2.1.6 define atom economy and calculate atom economies using chemical equations.

Discussion of the importance of atom economy in industry; examples include comparing the production of ethanol by the hydration of ethane to fermentation

www.bbc.co.uk/schools/gcsebitesize/science/add_ocr_gateway/chemical_economics/atomeconomyrev2.shtml


http://www.bbc.co.uk/schools/gcsebitesize/science/add_ocr_gateway/chemical_economics/atomeconomyrev2.shtml





Learning Outcomes


Practical Links


Suggested Time

2.2 Nomenclature and isomerism in organic compounds


2.2 Assessed as met in later units

CCEA Factfile

www.chemguide.co.uk/basicorg/isomerism/structural.html

Very detailed for reference:

www.chem.ucalgary.ca/courses/351/WebContent/orgnom/index.html

(covered when met in later units)

2.2.1 define and understand the terms structural and geometric, isomerism, homologous series and functional group;

Students define the terms in the specification and through dealing with the structures gain an understanding of the terms

2.2.2 apply IUPAC rules for nomenclature to name organic compounds with up to six carbon atoms and one or

Students discuss the importance of having IUPAC rules and practice naming molecules as they

Quiz:

www2.chemistry.msu.edu/faculty/reusch/VirtTxtJml/Questions/Nomencl/nomencl.htm


http://www2.chemistry.msu.edu/faculty/reusch/VirtTxtJml/Questions/Nomencl/nomencl.htm



http://www.chem.ucalgary.ca/courses/351/WebContent/orgnom/index.html



http://www.chemguide.co.uk/basicorg/isomerism/structural.html




more functional groups;

meet them throughout the unit. The numbers one to ten for alkanes i.e. methane to decane are recorded




Learning Outcomes


Practical Links


Suggested Time

2.2 Nomenclature and isomerism in organic compounds (cont.)


2.2.3 draw and name structural isomers of aliphatic compounds containing up to six carbon atoms, excluding cyclic structures;

Students make models of organic molecules. They should draw and name the compounds as they meet them in the rest of the unit. Record and explain the terms structural, aliphatic and cyclic

2.2.4 draw structural and skeletal formulae for organic compounds;

Students use models to show the ‘zig-zag’ shape of carbon chains and use them to draw

www.ausetute.com.au/skeletal.html


http://www.ausetute.com.au/skeletal.html

http://www.ausetute.com.au/skeletal.html


skeletal formulae of alkanes and organic compounds and the geometry of E/Z isomers in alkenes, students record skeletal formulae of molecules as they meet them




Learning Outcomes


Practical Links


Suggested Time

2.2 Nomenclature and isomerism in organic compounds (cont.)


2.2.5 understand that geometrical isomers result from restricted rotation due to an energy barrier about the carbon-carbon double bond and exist in E and Z forms; and

Students record why E/Z isomers exist and study examples of the isomers

www.chemguide.co.uk/basicorg/isomerism/geometric.html

2.2.6 draw and identify the structural formulae of E and Z isomers.

Students practise drawing and identifying E/Z isomers according to the priority rules which are recorded. A

http://chemwiki.ucdavis.edu/Organic_Chemistry/Fundamentals/Structure_of_Organic_Molecules/The_E-Z_system_for_naming_alkenes





http://www.chemguide.co.uk/basicorg/isomerism/geometric.html




variety of organic compounds may be used to illustrate the rules




Learning Outcomes


Practical Links


Suggested Time

2.3 Alkanes Students should be able to:

2.3 Assessed using past paper questions, practice questions and end of unit test

9 hours

2.3.1 recall that alkanes are described as saturated hydrocarbons;

Students should record the definition of saturated and hydrocarbon. Introduce the terms: homologous series, aliphatic, structural isomer and functional group. Students practice drawing and naming alkanes using

Past paper AS 2January 2012 Q 3June 2012 Q 2January 2011 Q 2, 13 a, bJune 2011 Q 1, 7, 12

CCEA Factfile

www.chemguide.co.uk/organicprops

www.chemguide.co.uk/organicprops/alkanes/background.html


http://www.chemguide.co.uk/organicprops/alkanes/background.html



http://www.chemguide.co.uk/organicprops

http://www.chemguide.co.uk/organicprops


both structural and skeletal formulae




Learning Outcomes


Practical Links


Suggested Time

2.3 Alkanes (cont.)


2.3.2

2.3.3

explain, in terms of van der Waals’ forces, the variation in boiling points between alkanes with different numbers of carbon atoms;

explain, in terms of van der Waals’ forces, the variation in boiling points between structural isomers of an alkane with the same molecular

Students discuss and record the boiling points of alkanes and plot a graph of boiling point against molecular mass. Link to the number of electrons in the molecules. Link to the fractional distillation of alkanes from crude oil

Students make models of the isomers of C5H12 record the boiling points, draw skeletal structures and illustrate the variation of

http://chemistry.elmhurst.edu/vchembook/501hcboilingpts.html

https://en.wikibooks.org/wiki/A-level_Chemistry/OCR/Chains,_Energy_and_Resources/Basic_Concepts_and_Hydrocarbons/Alkanes









formula; boiling points based on the length of chain and branching within the molecule




Learning Outcomes


Practical Links


Suggested Time

2.3 Alkanes (cont.)


2.3.4

2.3.5

describe the complete and incomplete combustion of alkanes in air and link the appearance of the flame to the amount of carbon present;

recall that pollutants such as carbon monoxide, carbon, oxides of nitrogen and sulfur and unburned hydrocarbons are

Students record equations for complete and incomplete combustion involving carbon and carbon monoxide and mixtures of them. Note the appearance of a Bunsen flame with air hole open and closed. Burn alkanes with increasing numbers of carbon atoms and note the flame appearance

Discussion of the combustion of alkanes and their use as fuels.

www.bbc.co.uk/bitesize/standard/chemistry/materialsfromoil/hydrocarbons/revision/3/

Students create summary presentation on 2.3.5 – 2.3.7


http://www.bbc.co.uk/bitesize/standard/chemistry/materialsfromoil/hydrocarbons/revision/3/




produced during the combustion of alkane fuels;

Explain why carbon monoxide, carbon nitrogen oxides, sulfur dioxide and hydrocarbons are regarded as pollutants and the specific harm they cause




Learning Outcomes


Practical Links


Suggested Time

2.3 Alkanes (cont.)


2.3.6

2.3.7

recall that the percentage of carbon dioxide in the atmosphere has risen from 0.03% to 0.04% because of combustion of organic compounds, and is believed to have caused global warming;

explain how a catalytic converter reduces the environmental impact of burning alkane fuels;

Discussion on global warming and the links to carbon dioxide production. Students suggest the ways that carbon dioxide is added and removed from the atmosphere

Students research the impact of catalytic converters and record the equations for removing nitrogen oxides and carbon monoxide from burning alkane

www.chemistry.wustl.edu/~edudev/LabTutorials/CourseTutorials/Tutorials/AirQuality/CatalyticConverter.htm


http://www.chemistry.wustl.edu/~edudev/LabTutorials/CourseTutorials/Tutorials/AirQuality/CatalyticConverter.htm




fuels2.3.8 describe the

substitution reactions of alkanes by chlorine and by bromine;

Explain substitution reactions. Students record equations for chlorine and bromine reacting with alkanes to give a variety of mono and polysubstituted products




Learning Outcomes


Practical Links


Suggested Time

2.3 Alkanes (cont.)


2.3.9

2.3.10

define the terms radical, homolytic and heterolytic fission; and

outline the radical substitution mechanism involved in the photochemical halogenation of alkanes in terms of initiation, propagation and termination steps.

Introduce the idea of a mechanism. Students record the definitions of radical and homolytic and heterolytic fission

Discuss the mechanism of the photochemical halogenation of alkanes. Students record examples of various alkanes reacting with chlorine and bromine writing the initiation, propagation and termination steps

www.chemguide.co.uk/mechanisms/freerad/ch4andcl2.html

Students practice drawing mechanism for various reactions involving chlorine and bromine reacting with different alkanes




Learning Outcomes


Practical Links


Suggested Time

2.4 Alkenes Students should be able to:


9 hours

2.4.1

define the term unsaturated hydrocarbon and explain why alkenes are described as unsaturated hydrocarbons;

Students record the definition of unsaturated hydrocarbon. Students practise drawing and naming structural and skeletal formulae, including E/Z isomers

Past paper AS 2January 2012 Q 13June 2012 Q 13January 2011 Q 5, 13 c, d, eJune 2011 Q 2, 9

CCEA Factfile

www.bbc.co.uk/schools/gcsebitesize/science/ocr_gateway/carbon_chemistry/making_polymersrev2.shtml

2.4.2 recall the qualitative test for alkenes using bromine water;

Students record general equation and discuss why alkenes react

Test for unsaturation using bromine water


http://www.bbc.co.uk/schools/gcsebitesize/science/ocr_gateway/carbon_chemistry/making_polymersrev2.shtml




readily with bromine whereas alkanes do not. The simple explanation ‘alkenes are more reactive’, is developed further in 2.4.4


Learning Outcomes


Practical Links


Suggested Time

2.4 Alkenes (cont.)


2.4.3 use sigma and pi bonds to explain the relative bond strength and relative bond length of the C=C bond;

Using the shapes of orbitals construct sigma and pi bonds. Discussion about why the sigma bond is stronger than a pi bond. Students record diagram of bonding in alkenes and compare bond strength and length to those in alkanes

www.chemguide.co.uk/organicprops/alkenes/background.html


http://www.chemguide.co.uk/organicprops/alkenes/background.html




2.4.4 recall that the C=C bond is a centre of high electron density and use this to explain the difference in reactivity of alkanes and alkenes;

Linking to 2.4.2, students record that the C=C bond is a centre of high electron density and use this to explain the difference in reactivity of alkanes and alkenes. Link this to the attack by electrophiles




Learning Outcomes


Practical Links


Suggested Time

2.4 Alkenes (cont.)


2.4.5 describe the catalytic hydrogenation of alkenes using finely divided nickel;

Discuss the addition of hydrogen to alkenes. Students record the general reaction and the catalyst. Discussion about this reaction’s use in industry for hardening fats

2.4.6 describe the reaction of chlorine, bromine, hydrogen chloride and hydrogen bromide with alkenes;

Introduce the term addition reaction. Students write equations for alkenes reacting with chlorine, bromine, hydrogen chloride and hydrogen

Test for unsaturation using bromine water



bromide

2.4.7 define the terms electrophile and heterolytic fission;

Students record definitions of electrophile and heterolytic fission

www.chemguide.co.uk/mechanisms/eladdmenu.html


http://www.chemguide.co.uk/mechanisms/eladdmenu.html





Learning Outcomes


Practical Links


Suggested Time

2.4 Alkenes (cont.)


2.4.8 recall the mechanism of electrophilic addition between chlorine, bromine, hydrogen chloride and hydrogen bromide with alkenes using curly arrows;

Using bromine reacting with ethene students record the mechanism for electrophilic addition. Explanations are given with regard to the use of curly arrows. Students practice writing the mechanisms for chlorine, hydrogen chloride and hydrogen bromide with alkenes

2.4.9 explain, with reference to the stability of the carbocation

Using propene reacting with hydrogen bromide,

www.youtube.com/watch?v=EWOvFAu8FmA


http://www.youtube.com/watch?v=EWOvFAu8FmA

http://www.youtube.com/watch?v=EWOvFAu8FmA


intermediates involved, the formation of major and minor products during the electrophilic addition of hydrogen bromide to unsymmetrical alkenes; and

students make models of the reactants and the intermediates in the mechanism. Explain why two different products are formed and why one product is favoured over the other




Learning Outcomes


Practical Links


Suggested Time

2.4 Alkenes (cont.)


2.4.10 describe the addition polymerisation of alkenes, for example, ethene and propene.

Students record the definition of addition polymerisation, and write equations to represent polymerisation of three molecules of ethene and propene. Polymers from a variety of alkenes are deduced e.g.buta-1-3-diene

www.chemguide.co.uk/organicprops/alkenes/polymerisation.html


http://www.chemguide.co.uk/organicprops/alkenes/polymerisation.html





Learning Outcomes


Practical Links


Suggested Time

2.5 Halogenoalkanes



Past paper AS 2January 2012 Q 6, 9, 15June 2012 Q 9January 2011 Q 17June 2011 Q 8, 16

CCEA Factfile

9 hours

2.5.1 explain the variation in boiling points of halogenoalkanes with different halogen atoms;

Introduce halogenoalkanes as a new homologous series. Students record structural and skeletal formulae and named examples of structural isomers

Students record the variation in boiling points of halogenoalkanes

www.chemguide.co.uk/organicprops/haloalkanemenu.html


http://www.chemguide.co.uk/organicprops/haloalkanemenu.html




with different halogen atoms and explain them using a combination of dipolar forces and van der Waals’ forces




Learning Outcomes


Practical Links


Suggested Time

2.5 Halogenoalkanes(cont.)


2.5.3 classify a halogenoalkane as primary, secondary or tertiary;

Students record explanations of primary, secondary and tertiary carbon atoms and the resulting halogenoalkane structures. Using structural isomers students classify halogenoalkanes as primary, secondary or tertiary

2.5.2 explain the variation in boiling points of structural isomers of a halogenoalkane

Students discuss, record and explain the variation in boiling points of structural



with the same molecular formula;

isomers of a halogenoalkane. Explanations are given according to chains and branching and van der Waals’ forces




Learning Outcomes


Practical Links


Suggested Time



2.5.4 describe the laboratory preparation of a liquid organic compound, such as a halogenoalkane, from the corresponding alcohol;

Students record the method, including purification, for example, the production of1-bromobutane from butan-1-ol. Each step in the method is discussed and explained.

The reason for the equipment and the chemicals used is explained. Labelled diagrams for distillation and reflux are accurately drawn

Prepare a halogenoalkane using the techniques of refluxing, separating with a funnel, removing acidity, drying and distillation.Also, links to % yield

www.youtube.com/watch?v=ZTY2DdJY84U


http://www.youtube.com/watch?v=ZTY2DdJY84U

http://www.youtube.com/watch?v=ZTY2DdJY84U



Learning Outcomes


Practical Links


Suggested Time



2.5.5 describe the reaction of halogenoalkanes with aqueous alkali, ammonia and potassium cyanide;

Students discuss, write equations and name products for the reactions of halogenoalkanes with aqueous alkali, ammonia and potassium cyanide. Explain the use of the term ‘aqueous’ alkali compared to ethanolic alkali in 2.5.8

Prepare alcohols from halogenoalkanes using alkali

2.5.6 define the term nucleophile and outline the nucleophilic substitution mechanism involved in the

Students record the definition of the term nucleophile. Outline the two nucleophilic substitution

www.youtube.com/watch?v=DPA9yDmjuQo


http://www.youtube.com/watch?v=DPA9yDmjuQo

http://www.youtube.com/watch?v=DPA9yDmjuQo


reaction between primary and tertiary halogenoalkanes and aqueous alkali;

mechanisms involved in the reaction of aqueous alkali with primary and tertiary halogenoalkanes. The carbocations can be linked to 2.4.9




Learning Outcomes


Practical Links


Suggested Time



2.5.7 describe and explain, with reference to bond enthalpy, the relative rates of hydrolysis of primary halogenoalkanes with the same number of carbon atoms and different halogen atoms;

Discuss how the C-X bond strength (bond enthalpy) and the C-X bond polarity might affect the rate of the hydrolysis. Students record the conclusion. Revise covalent bonds and their relative strengths first. Students record the several methods of experiment and the reasons for the experimental methods

Investigate the relative rates of hydrolysis of halogenoalkanes

2.5.8 describe Introduce the Carry out the www.chemguide.co.uk/


http://www.chemguide.co.uk/mechanisms/elimmenu.html


elimination of hydrogen halides from symmetrical and unsymmetrical halogenoalkanes using ethanolic potassium hydroxide; and

term elimination reaction and the reasons why ethanolic alkali is used. Students record and name the products from symmetrical and unsymmetrical halogenoalkanes

elimination of hydrogen halides from halogenoalkanes using ethanolic potassium hydroxide

mechanisms/elimmenu.html






Learning Outcomes


Practical Links


Suggested Time



2.5.9 recall that chlorofluorocarbons (CFCs) are a major factor in reducing the ozone layer and allowing more harmful ultraviolet radiation to reach the Earth’s surface.

Discussion about CFCs forming radicals with ultraviolet light i.e. similar to photochemical chlorination. Ozone reacts with the radicals formed.

Discuss and record the problems associated with more ultraviolet light reaching the Earth’s surface

Students create presentation on 2.5.9

www.theozonehole.com/cfc.htm

www.ausetute.com.au/cfcozone.html


http://www.ausetute.com.au/cfcozone.html

http://www.ausetute.com.au/cfcozone.html

http://www.theozonehole.com/cfc.htm

http://www.theozonehole.com/cfc.htm



Learning Outcomes


Practical Links


Suggested Time

2.6 Alcohols Students should be able to:

2.6. Assessed using practice questions, past paper questions and end of unit test

7 hours

2.6.2

classify an alcohol as primary, secondary or tertiary;

Students should use the rules established in 2.5.3 for classifying alcohols as primary, secondary and tertiary. Students should draw and name a variety of alcohols using structural and skeletal formulae

Past paper AS 2June 2012 Q 1, 15 a–bJanuary 2011 Q 16 bJune 2011 Q 10

CCEA Factfile

www.masterorganicchemistry.com/2010/06/16/1%C2%B0-2%C2%B0-3%C2%B0-4%C2%B0/

2.6.1 refer to the effect Students revise www.chemguide.co.uk/



http://www.masterorganicchemistry.com/2010/06/16/1%C2%B0-2%C2%B0-3%C2%B0-4%C2%B0/




of hydrogen bonding on boiling point and solubility of alcohols with water;

factors affecting boiling points, including length of chain, branching and hydrogen bonding. Students record effects of hydrogen bonding on boiling point and solubility of alcohols

atoms/bonding/hbond.html


Learning Outcomes


Practical Links


Suggested Time

2.6 Alcohols(cont.)


2.6.3 recall the preparation of alcohols from halogenoalkanes;

Students apply similar practical techniques and follow a similar method as used in 2.5.4

www.chemguide.co.uk/mechanisms/nucsub/hydroxide.html#top

2.6.4 describe the complete and incomplete combustion of

Students compare section 2.3.4. The flame produced by a


http://www.chemguide.co.uk/mechanisms/nucsub/hydroxide.html#top






alcohols and their use as an alternative fuel;

burning alcohol is compared to alkanes and alkenes. The percentage of carbon, in all three molecules, is compared. Students record the use of alcohols as an alternative fuel in cars etc.

2.6.5 describe reaction of alcohols with sodium, hydrogen bromide and phosphorus pentachloride; and

Students record and write equations for the reaction of alcohols with sodium, hydrogen bromide and phosphorus pentachloride

Carry out test tube reactions of alcohols with sodium, hydrogen bromide (hydrobromic acid) and phosphorus pentachloride

www.youtube.com/watch?v=AryisNZQCss


http://www.youtube.com/watch?v=AryisNZQCss

http://www.youtube.com/watch?v=AryisNZQCss



Learning Outcomes


Practical Links


Suggested Time

2.6 Alcohols(cont.)


2.6.6 describe the oxidation of alcohols using acidified potassium dichromate(VI) with reference to formation of aldehydes and carboxylic acids from primary alcohols, formation of ketones from secondary alcohols and resistance to oxidation of tertiary alcohols.

Students write equations for the reactions of acidified potassium dichromate (VI) with primary and secondary alcohols, describe the colour changes and the changes in smell. Students determine primary, secondary and tertiary alcohols from unknown samples

Prepare aldehydes, carboxylic acids and ketones from alcohols using acidified potassium dichromate(VI)

www.chemguide.co.uk/organicprops/alcohols/oxidation.html


http://www.chemguide.co.uk/organicprops/alcohols/oxidation.html





Learning Outcomes


Practical Links


Suggested Time

2.7 Infrared spectroscopy



4 hours

2.7.1

explain that the absorption of infrared radiation arises from molecular vibrations;

Explain, using a spring model, the concept of molecular vibrations. Explain that infrared radiation is part of the electromagnetic spectrum i.e. similar to light. Students record that the absorption of infrared radiation arises from molecular vibrations

Past paper AS 2June 2012 Q 6January 2011 Q 16 a

CCEA Factfile

Useful to contact RSC for ‘spectroscopy in a suitcase’ demonstrations



2.7.2 understand that groups of atoms within a molecule absorb infrared radiation at characteristic frequencies; and

Explain infrared absorptions by referring to the data sheet which shows bonds and the frequencies at which they absorb

www.rsc.org/learn-chemistry/wiki/Introduction_to_IR_spectroscopy

www.chem.ucalgary.ca/courses/350/Carey5th/Ch13/ch13-ir-4.html


http://www.chem.ucalgary.ca/courses/350/Carey5th/Ch13/ch13-ir-4.html



http://www.rsc.org/learn-chemistry/wiki/Introduction_to_IR_spectroscopy





Learning Outcomes


Practical Links


Suggested Time

2.7 Infrared spectroscopy(cont.)


2.7.3 use infrared spectra to deduce functional groups present in organic compounds given wavenumber data.

Students use infrared spectra to deduce functional groups as indicated by the data sheet. Useful to compare alcohols, aldehydes and ketones noting the similarities and differences in the spectra. Note the reasons for the broad absorptions of the OH bond in alcohols. Discuss the fingerprint region of an infrared spectrum although the

If time allows, issue a ½ mock on Organic 2.1.1 – 2.7.3



whole spectrum can be used for identification purposes




Learning Outcomes


Practical Links


Suggested Time

2.8 Energetics



9 hours

2.8.1

define the terms exothermic and endothermic and understand that chemical reactions are usually accompanied by heat changes;

Students discuss and record the definitions of exothermic and endothermic reactions/processes

Past paper AS 2January 2012 Q 1, 5, 7, 8, 14June 2012 Q 3, 15 c, dJanuary 2011 Q 7, 8, 10, 14June 2011 Q 4, 14

CCEA Factfile

2.8.2 recall standard conditions as 100 kPa and 298 K;

Students discuss the need for using standard conditions when comparing the properties of substances.

www.chemguide.co.uk/physical/energetics/definitions.html


http://www.chemguide.co.uk/physical/energetics/definitions.html




Students record standard conditions as stated




Learning Outcomes


Practical Links


Suggested Time

2.8 Energetics(cont.)


2.8.3

2.8.4

define the term standard enthalpy change ΔHΘ;

construct a simple enthalpy level diagram;

Students discuss the concept of change in enthalpy and record the definition of standard enthalpy change, ΔHΘ

Students discuss exothermic and endothermic reactions, as in 2.8.1 and record them in enthalpy level diagrams. Values for the exothermic and endothermic reactions are obtained from the labelled diagram. Reactants and

www.bbc.co.uk/bitesize/higher/chemistry/calculations_1/potential_energy/revision/1/


http://www.bbc.co.uk/bitesize/higher/chemistry/calculations_1/potential_energy/revision/1/




products are shown with assigned enthalpy values




Learning Outcomes


Practical Links


Suggested Time



2.8.5 define the standard enthalpy of combustion, formation and neutralisation, namely ΔcHΘ, ΔfHΘ

and ΔnHΘ;

Students record the definitions of enthalpies of combustion, formation and neutralisation reactions. Construct equations which show the standard enthalpy of combustion, formation and neutralisation for a series of compounds together with the reverse and multiples of the equations

2.8.6 recall experimental

Students carry out a wide variety

Determine the enthalpy



methods to determine enthalpy changes;

of experiments understanding the reasons for the practical techniques and follow through worked examples

changes for combustion and neutralisation using simple apparatus




Learning Outcomes


Practical Links


Suggested Time



2.8.7 calculate enthalpy changes from experimental data using the equation

q = mcΔT;

Using data gathered in 2.8.6 students calculate enthalpy changes from experimental data using the equationq = mcΔT and convert into molar values. Experimental errors (actual and percentages) may be calculated

www.bbc.co.uk/bitesize/higher/chemistry/calculations_1/potential_energy/revision/4/

2.8.8 appreciate the principle of conservation of energy and define Hess’s Law;

Students discuss the law of conservation of energy and record the definition of

www.chemguide.co.uk/physical/energetics/sums.html


http://www.chemguide.co.uk/physical/energetics/sums.html







Hess’s Law

2.8.9 construct enthalpy cycles using Hess’s Law;

Students record several examples of constructing enthalpy cycles

2.8.10 calculate enthalpy changes indirectly using Hess’s Law;

Students calculate a variety of enthalpy changes using Hess’s Law




Learning Outcomes


Practical Links


Suggested Time



2.8.11 define the term average bond enthalpy and calculate the enthalpy change of a reaction using average bond enthalpies;

Students record the definition of average bond enthalpy. Students record worked examples and practice calculations

www.youtube.com/watch?v=o8HF-kdQuys

CCEA Factfile

2.8.12 calculate average bond enthalpies given enthalpy changes of reaction; and

Discuss the use of average bond enthalpies and how they can be calculated using enthalpy changes of reaction

2.8.13 explain why enthalpy changes of reaction calculated using average bond enthalpies differ from those

Students record the limitations of calculations based on ‘average’ bond enthalpy then explain why


http://www.youtube.com/watch?v=o8HF-kdQuys

http://www.youtube.com/watch?v=o8HF-kdQuys


determined using Hess’s Law.

enthalpy changes of reaction, calculated using average bond enthalpies, differ from those determined using Hess’s Law




Learning Outcomes


Practical Links


Suggested Time

2.9 Kinetics Students should be able to:


5 hours

2.9.1 recall how factors, including concentration, pressure, temperature and catalyst, affect the rate of a chemical reaction;

Students discuss the concept of rate of reaction, defined as change in concentration (reactant or product) per unit time e.g. mol dm-3

s-1. Look at graphs of concentration of products/reactants versus time. Students

Past paper AS 2January 2012 Q 11June 2012 Q 10, 12 aJanuary 2011 Q 11 (mixed), 15June 2011 Q 5, 15 c

CCEA Factfile

www.chemguide.co.uk/physical/basicratesmenu.html


http://www.chemguide.co.uk/physical/basicratesmenu.html




investigate the factors affecting rate in a variety of experiments




Learning Outcomes


Practical Links


Suggested Time

2.9 Kinetics(cont.)


2.9.2

2.9.3

use the collision theory and the concept of activation energy to qualitatively explain how these factors affect the reaction rate;

demonstrate a qualitative understanding of the Maxwell-Boltzmann distribution of molecular energies in gases and interpret curves for different temperatures and for catalysed and

Students record the definition of activation energy and include it in enthalpy diagrams of exothermic and endothermic reactions. Collision theory is used to explain the factors affecting reaction rate

Students record the Maxwell-Boltzmann distribution of molecular energies. Explain the meaning of the distribution curve and molecular



uncatalysed reactions; and

energies. Explain the changes in the curve and their significance for different temperatures and for catalysed and uncatalysed reactions

2.9.4




Learning Outcomes


Practical Links


Suggested Time

2.9 Kinetics(cont.)


2.9.4 relate the concept of activation energy to the Maxwell-Boltzmann distribution.

Students relate the changes in the Maxwell-Boltzmann distribution curve to the changes in the activation energy and how this relates to the rate of a reaction




Learning Outcomes


Practical Links


Suggested Time

2.10 Equilibrium



9 hours

2.10.1

appreciate that many chemical reactions are reversible and define the terms dynamic equilibrium, homogeneous and heterogeneous;

Students record the definitions of the terms dynamic equilibrium, homogeneous and heterogeneous. Discuss a variety of equilibrium changes e.g. changes in state as an example of a reversible physical change and many chemical reactions e.g.

Past paper AS 2January 2012 Q 16June 2012 Q 12 bJune 2011 Q 6, 15 a, b

CCEA Factfile

www.bbc.co.uk/bitesize/higher/chemistry/reactions/equilibrium/revision/1/

www.bbc.co.uk/bitesize/higher/chemistry/reactions/equilibrium/revision/2/


http://www.bbc.co.uk/bitesize/higher/chemistry/reactions/equilibrium/revision/2/







hydrated and anhydrous salts. Students record the meaning of a closed homogeneous system




Learning Outcomes


Practical Links


Suggested Time

2.10 Equilibrium(cont.)


2.10.2 deduce the qualitative effects of changes of temperature, pressure, concentration and catalysts on the position of equilibrium for a closed homogeneous system;

Students understand the terms LHS (left hand side) and RHS (right hand side). Explain that if an endothermic system moves to the LHS it is exothermic. Explain that more molecules on the RHS mean an increase in pressure. Catalysts have no effect on the position of equilibrium but they increase the rate of the forward and reverse reactions equally

www.chemguide.co.uk/physical/equilibria/lechatelier.html


http://www.chemguide.co.uk/physical/equilibria/lechatelier.html




2.10.3 deduce an expression for the equilibrium constant, KC and its units for a given homogeneous equilibrium system;

Introduce the idea of KC and the use of [ ] to denote concentrations, leading to an expression for the equilibrium constant, KC. Determine the units for a given equilibrium system

www.chemguide.co.uk/physical/equilibria/kc.html


Learning Outcomes


Practical Links


Suggested Time



2.10.4 relate the magnitude of KC to the position of equilibrium and extent of reaction;

Students discuss the equation for an equilibrium. A large value for KC

indicates the reaction is towards the RHS. A low value indicates the reaction is


http://www.chemguide.co.uk/physical/equilibria/kc.html




towards the LHS. A value of 1 means the equilibrium is neither to one side or the other

2.10.5 describe and explain the conditions used in industrial processes, for example the Haber process for the formation of ammonia and the Contact process for sulfuric acid; and

Students could research and describe and explain the conditions used in the industrial processes. Conditions considered should be temperature, pressure and catalyst and how they affect the equilibrium. The effects of costs re conditions should also be considered




Learning Outcomes


Practical Links


Suggested Time



7 hours

2.10.6 demonstrate an understanding of the importance of a compromise between equilibrium and reaction rate in the chemical industry.

Students record an explanation of the importance of getting a low yield quickly versus a high yield slowly and the costs associated with both methods




Learning Outcomes


Practical Links


Suggested Time

2.11 Group II elements and their compounds


This topic could be completed as a self-study unit with practicals being completed in class when the topic is being checked

Students create presentation on self-study elements

2.11 Assessed using practice questions, past paper questions and end of unit test. If time allows issue mock on full AS 2

Past paper AS 2January 2012 Q 12June 2012 Q 4, 14January 2011 Q 12June 2011 Q 13

CCEA Factfile

7 hours

2.11.1 explain why these are regarded as s-block elements;

Students record Group II electronic configurations and note the outer electronic structure

www.chemguide.co.uk/inorganic/group2menu.html

2.11.2 recall and explain the trends within

Students explain the trends within

https://en.wikibooks.org/wiki/A-


https://en.wikibooks.org/wiki/A-level_Chemistry/OCR/Group_2


http://www.chemguide.co.uk/inorganic/group2menu.html




the Group, limited to electronic configuration, atomic radius and first ionisation energy;

the Group, by completing a table listing electronic configuration, atomic radius and first ionisation energy

level_Chemistry/OCR/Group_2






Learning Outcomes


Practical Links


Suggested Time

2.11 Group II elements and their compounds(cont.)


2.11.3 investigate and describe the reactions of the elements with oxygen, water and dilute acids;

Students investigate and describe the reactions of the elements with oxygen, water and dilute acids and include methods, collection of gases and balanced equations including ionic equations. The elements’ reactivities are compared by these reactions

React Group II metals and other metals with oxygen, water and dilute acids, and determine the masses of solids and volumes of gases produced




Learning Outcomes


Practical Links


Suggested Time

2.11 Group II elements and their compounds (cont.)


2.11.4 describe the basic nature of the oxides and their reactions with water and dilute acids;

Students record the basic nature of the oxides and their reactions with water and dilute acids to include observations, balanced equations and ionic equations. The effect of the solubility of the sulfates on reaction speed is compared

www.youtube.com/watch?v=i-rFsFwdkTU

2.11.5 recall the use of magnesium oxide in indigestion remedies and the use of calcium

Students investigate the properties of a variety of indigestion


http://www.youtube.com/watch?v=i-rFsFwdkTU

http://www.youtube.com/watch?v=i-rFsFwdkTU


carbonate in toothpaste;

tablets and toothpastes with water and acids and their pHs. Rough titration with acids determines the amount of base present




Learning Outcomes


Practical Links


Suggested Time



2.11.6 state the trends in thermal stability of the carbonates and hydroxides and explain with reference to the charges and sizes of the cations;

Students record the definition of thermal decomposition and explain the trend in thermal stability of the carbonates and hydroxides using the ideas of polarisation of anions and charge density of cations

www.chemguide.co.uk/inorganic/group2/thermstab.html

2.11.7 recall the use of calcium carbonate to make calcium oxide (quick lime) and calcium hydroxide (slaked lime) and their

Students draw the limestone cycle, recording all the reactions, equations and observations. Students react

www.bbc.co.uk/schools/gcsebitesize/science/aqa/limestone/calciumcarbonaterev3.shtml


http://www.bbc.co.uk/schools/gcsebitesize/science/aqa/limestone/calciumcarbonaterev3.shtml



http://www.chemguide.co.uk/inorganic/group2/thermstab.html




use in producing cement and concrete;

cement and concrete with hydrochloric acid and record how calcium hydroxide is used to make them




Learning Outcomes


Practical Links


Suggested Time



2.11.8 recall the solubility trends of the sulfates and hydroxides; and

Students discuss the factors affecting solubility (strength of ionic bond, size of ion/ease of hydration) and use these factors to explain the trends in solubility of the sulfates and hydroxides

www.chemguide.co.uk/inorganic/group2/solubility.html

2.11.9 understand how solubility curves are drawn from experimental data.

Students plot solubility curves for Group II compounds such as nitrates and sulfates using experimental

www.docbrown.info/page03/AcidsBasesSalts08.htm

If time allows issue mock on full AS 2 and a practical mock


http://www.chemguide.co.uk/inorganic/group2/solubility.html




data. Students use solubility curves to predict how much solid is deposited when a solution is cooled

Students should have the opportunity to complete as many past papers as possible, including practical exams



Unit A2 1:Further Physical and

Organic Chemistry




Unit A2 1: Further Physical and Organic Chemistry

Specification Content

Learning Outcomes


Practical Links


Suggested Time

4.1 Lattice enthalpy


Calculations for A-Level Chemistry – E.N Ramsden

Past paper A2 1January 2014 Q11

6 hours

4.1.1 define and understand the term lattice enthalpy;

Students define lattice enthalpy and explain it in terms of a lattice of ions and their attractions and repulsions.The ions can be of various charges

CCEA Factfile

http://chemguide.co.uk/physical/energetics/lattice.html#top







Learning Outcomes


Practical Links


Suggested Time

4.1 Lattice enthalpy (cont.)


4.1.2 construct Born-Haber cycles and carry out associated calculations, such as the halides and oxides of Groups I and II; and

Hess’s law is applied to the Born-Haber cycle for sodium chloride i.e. the energy changes are the same no matter what route is taken in the cycle. Further cycles are constructed and a variety of calculations for various values in the cycle. Students define enthalpy of atomisation, formation, electron affinity, bond enthalpy and ionisation energy

Past paper A2 1June 2014 Q 15January 2014 Q 15




Learning Outcomes


Practical Links


Suggested Time

4.1 Lattice enthalpy (cont.)


4.1.3 define and understand the enthalpy changes associated with the dissolving of ionic compounds in water, and carry out associated calculations.

Students define the enthalpy of solvation/solution. Show an example of dissolving sodium chloride in water and the associated calculations for either enthalpy of solvation or lattice enthalpy (ΔHLE) from a constructed cycle. Students construct similar cycles for other compounds using data provided. Students draw diagrams for the solvation of the ions




Learning Outcomes


Practical Links


Suggested Time

4.2 Enthalpy, entropy and free energy


6 Hours

4.2.1 recall that enthalpy change is not sufficient to explain feasible change, for example the endothermic reaction between ammonium carbonate and ethanoic acid;

Explain that feasible change causes energy and/or matter to spread out regardless of what the enthalpy change is

CCEA Factfile

4.2.2 recall that the balance between entropy change and enthalpy change determines the feasibility of a reaction;

Endothermic reactions are not feasible, they shouldn’t happen, but they do in some cases. Use the example in 4.2.1 and e.g. certain salts dissolving in water




Learning Outcomes


Practical Links


Suggested Time

4.2 Enthalpy, entropy and free energy (cont.)


4.2.3 recall that entropy is a measure of disorder;

Entropy is a measure of the degree of disorder of a system, symbol ‘S’. Use a tidy and untidy room as an example. High entropy is untidy. Explain the idea that the more disordered a state is, resulting from reaction, the more likely it is to occur e.g. going from solid reactants to gaseous products. Two noble gases mix completely despite there



being no enthalpy changes

4.2.4 calculate the standard entropy change, ΔS , in a chemical reaction using standard entropy data;

Introduce the equation for calculating entropy change;∆S= Sproducts – Sreactants. Note that the units for entropy are in J mol-1 K-1




Learning Outcomes


Practical Links


Suggested Time



4.2.5 use the equationΔG = ΔH – TΔS to calculate standard free energy changes;

Introduce free energy, symbol ΔGθ, as the combination of the effect of entropy and enthalpy on a system.Students pick values for entropy and enthalpy where a reaction is likely to occur. Students deduce that ΔGθ for a feasible reaction is always negative

4.2.6 recall that processes are feasible when the

If ΔGθ is positive a reaction does not proceed i.e.



free energy change is negative; and

the reaction goes to the left rather than the right




Learning Outcomes


Practical Links


Suggested Time



4.2.7 recall that when the enthalpy change and the entropy change have the same sign the feasibility of the process depends on the temperature, and calculate the temperature at which these processes start/cease to be feasible.

Student perform calculations involving the equation forΔGθ = ΔHθ – TΔSθ to see if the value is negative

http://chemguide.co.uk/physical/basicratesmenu.html







Learning Outcomes


Practical Links


Suggested Time

4.3 Rates of Reaction


10 hours

4.3.1

use simple rate equations in the form: rate = k[A]X[B]Y (where x and y are 0, 1 or 2);

Revise kinetics from unit 2. Introduce the rate equation and define the terms. Explain that rate is proportional to certain concentrations and that the proportional sign is replaced by an equation

Carry out experiments to determine the rate of a reaction using a variety of methods to determine the concentration of reactants and/or products

Past paper A2 1June 2013 Q12 bJune 2011 Q11 aJune 2010 Q14 d

CCEA Factfile

http://chemguide.co.uk/physical/equilibmenu.html

4.3.2 understand the terms: rate of

reaction; order; and

Discuss the absence of pressure, surface area and temperature in






rate constant; the equation. Explain what terms must change in the equation if rate is to increase. Define and record the meanings of the terms order and rate constant




Learning Outcomes


Practical Links


Suggested Time

Rates of Reaction(cont.)


4.3.3 deduce simple rate equations from experimental data;

Show how a rate equation can be deduced from experimental data.Show how units can be deduced from rate equations.Students vary the concentration ofNa2S2O3 and time how long it takes for a sulfur precipitate to form

4.3.4 deduce, from a concentration-time or a rate-concentration graph, the rate of reaction and/or the order with respect to a

Plot graphs of concentration against time. They can be of reactants or products. Students calculate the rate



reactant; of reaction inmol dm-3. The graphs of rate v concentration for 0, 1st and 2nd order are shown




Learning Outcomes


Practical Links


Suggested Time



4.3.5 recall that there is a relationship between the rate equation and mechanism, for example SN1 and SN2 mechanisms for the alkaline hydrolysis of primary and tertiary halogenoalkanes;

Revise mechanisms from unit 2, i.e. the hydrolysis of halogenoalkanes in terms of kinetics and explain what is meant by SN1 and SN2 mechanisms

4.3.6 define and understand the term ‘rate determining step’;

Students record the definition of ‘rate determining step’. Introduce the idea of the slowest step of a reaction determining the rate i.e. the ‘rate determining step’




Learning Outcomes


Practical Links


Suggested Time



4.3.7 suggest experimental methods suitable for the study of the rate of a reaction, for example iodine titrations and colorimetry; and

Revise experimental methods which have changes that can be measured e.g. changes in pH, mass, colour, density, gas volumes. They can be changes in reactants or changes in products. Introduce the ideas of ‘quenching’ and ‘sampling’




Learning Outcomes


Practical Links


Suggested Time



4.3.8 explain, qualitatively, the effect of temperature and activation energy on rate constants.

Revise the Maxwell-Boltzmann distribution curve from 2.9 which shows how temperature affects rate in terms of particles and their energy. If the rate increases the rate constant must increase when the temperature is increased.Recall kinetic stability from unit 2. Explain the idea that high kinetic stability leads to high activation energy and a slow rate.



If the rate is low and the concentration is unchanged, the rate constant must be small




Learning Outcomes


Practical Links


Suggested Time

4.4 Equilibrium


10 hours

4.4.1

calculate equilibrium concentrations given suitable data; and

Revise equilibrium from unit 2 and how different factors affect equilibria together with unit 4. Calculate equilibrium concentrations when moles of products and/or reactants are given for reactions e.g. 4NH3 + 5O2→4NO + 6H2O using two lines for substances before

Past paper A2 1June 2014 Q 12January 2104 Q 17

CCEA Factfile

www.chemguide.co.uk/physical/acideqiamenu.html


http://www.chemguide.co.uk/physical/acideqiamenu.html




equilibrium and at equilibrium




Learning Outcomes


Practical Links


Suggested Time

4.4 Equilibrium (cont.)


4.4.2 calculate the numerical values, with units, for equilibrium constants, KC, given suitable data limited to homogeneous systems.

Revise equilibrium from unit 2 which states that an expression for the equilibrium constant can be deduced. Show worked example of calculating Kc with initial moles, moles reacting and moles at equilibrium. Students are encouraged to show working out. Show examples of calculation of units. Look at values of KC and link to the amount of



product formed




Learning Outcomes


Practical Links


Suggested Time

4.5 Acid-base equilibria


12 hours

4.5.1 use the Brønsted-Lowry theory of acids and bases to describe proton transfer in acid-base equilibria including the idea of conjugateacid-base pairs;

Revise that all acids produce H+ ions/protons. Introduce the idea of acids as proton donors and bases as proton acceptors. Introduce conjugate acid-base pairs using a variety of equations which students analyse. Revise what is meant by strong and weak acids and bases in terms of conjugate acid-base pairs

CCEA Factfile




Learning Outcomes


Practical Links


Suggested Time

4.5 Acid-base equilibria (cont.)


4.5.2 define and understand the terms KW, Ka, pH, pKW and pKa, and recall the associated units where appropriate;

Students record the definitions of the terms KW, Ka, pH, pKW and pKa, and the associated units. Introduce the ionisation of water and students write an equilibrium expression for this. The concentration of water being constant leads to KW.Students work out whether the dissociation of water is endo or exothermic




Learning Outcomes


Practical Links


Suggested Time



4.5.3 carry out calculations involving pH for strong acids, strong bases and weak acids;

Students are given the idea that very low values of H+ means that it is easier to look at logs i.e. pH. Introduce simple calculations of pH and the reverse to work out [H+] Introduce pOH. Look at the example of how acid concentration changes when alkali added

Introduce the expression for Ka and use it to



calculate pH. The use of pKa




Learning Outcomes


Practical Links


Suggested Time



4.5.4 define and understand the term buffer solution and give a qualitative explanation of how buffer solutions work;

Students record the definition of a buffer solution and its uses in biological systems. Write the buffer equation by taking logs of the quantities in the equationK[HX] = [H+][X-]. Show how buffers work when small amounts of acid/alkali are added

4.5.5 calculate the pH of a buffer solution made from a weak monobasic acid and sodium

Students carry out calculations using the buffer equation where [HX] = concentration of

Make buffer solutions from calculated quantities of salts and acids and determine



hydroxide; acid and [X-]= the concentration of salt

their pH values using universal indicator (UI) paper and a pH meter




Learning Outcomes


Practical Links


Suggested Time



4.5.6 recall how titration curves are determined by experiment;

Students use data loggers and pH probes to plot titration curves for SA-SB, WA-SB, SA-WB, WA-WB. Students also calculate some of the points on the curves.Introduce the term equivalence point

4.5.7 use titration curves to explain the choice of indicator; and

Students choose indicators based on changes in pH during the vertical section of a titration curve from a table of indicators. Students explain

Determine the shape of a titration curve by measuring the pH using specialised pH paper or a pH meter for the titration of an



the reasons for choosing appropriate indicators

acid with a base




Learning Outcomes


Practical Links


Suggested Time



4.5.8 predict whether a salt solution would be acidic, alkaline or neutral based on relative strengths of the parent acid and base.

Students look at salt hydrolysis e.g. SA-SB gives a salt which forms a neutral solution; WA-SB gives an alkaline solution; SA-WB gives an acidic solution. Students work through examples using a data book and test the pH of ammonium and aluminium salts

Determine the pH of a variety of salts using pH paper or a pH meter to illustrate the relative strength of acid and base




Learning Outcomes


Practical Links


Suggested Time

4.6 Isomerism


3 hours

4.6.1

recognise that structural isomerism can exist between molecules which belong to different families of compounds;

Students revise structural and E-Z isomers from unit 2. Students look at examples of structural isomers in terms of carboxylic acids and esters which have the same molecular formula but different arrangement of atoms

Past paper A2 1June 2014 Q 13January 21014 Q 12 cJune 2013 Q 15

CCEA Factfile

http://chemguide.co.uk/basicorg/isomermenu.html

http://chemguide.co.uk/basicorg/isomerism/optical.html#top

4.6.2 recall that asymmetric

Students record the definition of









(chiral) centres give rise to optical isomers which exist as non-superimposable mirror images;

an asymmetric/chiral centre and an optical isomer. Introduce optical isomerism as another example of stereoisomerism.




Learning Outcomes


Practical Links


Suggested Time

4.6 Isomerism (cont.)


4.6.3 draw 3D representations of optical isomers;

Students make models of optical isomers and show that they cannot be superimposed

Students draw 3D structures of optical isomers using dotted lines and wedges which are mirror images. Tetrahedrons are also drawn of the mirror images

4.6.4 recall that optical isomers rotate the plane of plane polarised light in opposite directions;

Students record the definition of optical isomers in terms of rotating the plane of polarised light. Students revise what is meant by



plane polarised light

4.6.5 define and understand the term optically active and explain why racemic mixtures are optically inactive; and

Students define the term racemic mixture and explain why they are optically inactive


Learning Outcomes


Practical Links


Suggested Time

4.6 Isomerism (cont.)


4.6.6 recognise that drug action may be determined by the stereochemistry of the drug and that receptor sites can be stereospecific.

Receptor sites are on molecules which are chiral. Only one of the optical isomers of the drug is recognised by the site. Hence only 50% of the stereoisomers work




Learning Outcomes


Practical Links


Suggested Time

4.7 Aldehydes and ketones


3 hours

4.7.1 recall the molecular and structural formulae of aldehydes and ketones, including branched structures, with up to six carbon atoms in the main chain;

Revise the term functional group from unit 2. Revise the naming of carbon chains and use with the endings –al and -one e.g. methanal and propanone. Students name compounds and draw structures of carbonyl compounds

Past paper A2 1January 2014 Q 17 b

CCEA Factfile

http://chemguide.co.uk/organicprops/carbonyls/background.html

4.7.2 explain the boiling points and solubility of aldehydes and ketones by

Discuss the polarity of the carbonyl group. Relate the solubility and






making reference to intermolecular forces;

boiling points to the dipolar forces. Compare to alcohols and alkanes of similar mass which have hydrogen bonding and van der Waals’ forces




Learning Outcomes


Practical Links


Suggested Time

4.7 Aldehydes and ketones (cont.)


4.7.3 recall that aldehydes and ketones can be prepared from the corresponding primary or secondary alcohol using a suitable oxidising agent;

Revise alcohols from unit 2 and the oxidation of 1o and 2o alcohols to aldehydes and ketones using potassium dichromate (VI) noting the colour change

4.7.4 recall the reaction of aldehydes and ketones with hydrogen cyanide;

Students discuss the reactions of carbonyl compounds with the cyanide ion due to the polarity of C=O

4.7.5 describe the mechanism for the nucleophilic addition reaction of hydrogen

Look at the mechanism for the reaction using curly arrows and



cyanide and propanone;

naming the product. Students write the mechanism with a variety of aldehydes and ketones




Learning Outcomes


Practical Links


Suggested Time



4.7.6 explain why racemic mixtures can be produced when hydrogen cyanide reacts with aldehydes and ketones;

The cyanide ion can attack the planar carbonyl compound from either side which leads to the two isomers in a 50/50 mixture. Students make models to show the two reactions

4.7.7 recall the reaction of aldehydes and ketones with 2,4-dinitrophenylhydrazine;

Students write equations for the reaction of 2,4-DNPH with carbonyl compounds paying particular attention to the – NH-N=C- structure.The reaction involves the lone



pair on the outer nitrogen atom attacking the carbonyl group




Learning Outcomes


Practical Links


Suggested Time



4.7.8 recall the preparation of 2,4-dinitrophenylhydrazones for identification purposes with reference to melting point determination;

Students form2,4-DNPHs from various carbonyl compounds and determine the m.pt. to identify the aldehyde or ketone. Students practice re-crystallisation techniques

Prepare, recrystallise and determine the melting point of 2,4-dinitro-phenylhydrazones

4.7.9 recall that aldehydes and ketones can be distinguished using acidified potassium dichromate(VI), Fehling's solution and Tollens’ reagent (with Fehling's solution and Tollens’

Students write ionic equations for Tollens’ and Fehling’s solutions giving positive tests and state the colour changes observed. Students are given an aldehyde and

Use Fehling’s solution and Tollens’ reagent to distinguish between aldehydes and ketones



reagent viewed as Cu2+ and Ag+ respectively); and

ketone, labelled X and Y and perform a series of tests to determine which is which




Learning Outcomes


Practical Links


Suggested Time



4.7.10 recall that aldehydes and ketones can be reduced using lithium tetrahydridoaluminate (III) (lithal).

Students write equations for the reduction of aldehydes and ketones to 1o and 2o alcohols with lithal using [H]. Students note that this is similar to the mechanism of cyanide ion addition except that hydride ions are used




Learning Outcomes


Practical Links


Suggested Time

4.8 Carboxylic acids


3 hours

4.8.1

recall the molecular and structural formulae of carboxylic acids, including branched structures, with up to six carbon atoms in the main chain;

Students revise the –COOH functional group from unit 2 and revise the naming, including branched structures. Students are given structures to name and a selection of molecular formulae where they have to draw structures

Past paper A2 1January 2013 Q 13June 2012 Q 5, 7, 12

CCEA Factfile

http://chemguide.co.uk/organicprops/acidmenu.html#top






4.8.2 explain the boiling points and solubility of carboxylic acids by making reference to intermolecular attractions;

Students explain solubility and boiling points by referring to hydrogen bonds formed from the–COOH group. Mention is made of the ability to form dimers




Learning Outcomes


Practical Links


Suggested Time

4.8 Carboxylic acids (cont.)


4.8.3 recall that carboxylic acids can be prepared from primary alcohols and aldehydes to include practical details;

Students revise the oxidation of primary alcohols and aldehydes to carboxylic acids using acidified potassium dichromate (VI). Students compare the methods of preparing a liquid carboxylic acid to those of preparing a chloroalkane in 2.5.4

Prepare a carboxylic acid from an alcohol




Learning Outcomes


Practical Links


Suggested Time



4.8.4 recall that carboxylic acids, or their salts, can also be formed by acid or base catalysed hydrolysis of esters and nitriles;

Students revise the reversible reaction of alcohols and carboxylic acids to form esters. Idea that esters can be converted back to acid and alcohol using acid or base hydrolysis. Students using examples of nitriles and esters predict the structure of the carboxylic acids formed from their hydrolysis




Learning Outcomes


Practical Links


Suggested Time



4.8.5 recall that carboxylic acids form salts with sodium carbonate, sodium hydroxide and ammonia; and

Students revise carboxylic acids as weak acids which undergo partial dissociation in solution. Students write equations for the reactions of carboxylic acids with sodium carbonate, sodium hydroxide and ammonia. Explain the exothermic nature of sodium hydroxide with hydrochloric compared to ammonia with ethanoic acid




Learning Outcomes


Practical Links


Suggested Time



4.8.6 recall the reaction of carboxylic acids with alcohols, phosphorus pentachloride and lithium tetrahydridoaluminate (III) (lithal).

Students write equations for the reactions of carboxylic acids with alcohols, phosphorus pentachloride and lithal and state the observations to be made with phosphorus pentachloride

Carry out test tube reactions of a carboxylic acid with sodium carbonate, sodium hydroxide and aqueous ammonia, and measure the pH changes




Learning Outcomes


Practical Links


Suggested Time

4.9 Derivatives of carboxylic acids


3 hours

4.9.1 recall that derivatives of carboxylic acids include esters and acyl chlorides;

Students revise the ester functional group from unit 2 and learn the acyl functional group

Past paper A2 1June 2014 Q 12, 16June 2013 Q 11June 2012 Q 15

CCEA Factfile

www.chemguide.co.uk/organicprops/estermenu.html#top

www.chemguide.co.uk/organicprops/acylclmenu.html#top

4.9.2 recall the molecular and structural formulae of monoesters and of acyl chlorides;

Students name a variety of esters and acyl chlorides and draw their structures

4.9.3 explain the boiling points and solubility of

Students look at the structure of the esters and


http://www.chemguide.co.uk/organicprops/acylclmenu.html#top



http://www.chemguide.co.uk/organicprops/estermenu.html#top




monoesters by making reference to intermolecular forces;

relate it to boiling points and solubility. They compare the boiling points to alkanes, alcohols and carboxylic acids etc. of similar mass




Learning Outcomes


Practical Links


Suggested Time

4.9 Derivatives of carboxylic acids (cont.)


4.9.4 recall that esters can be formed from alcohols using carboxylic acids or acyl chlorides;

Revise from unit 2 the formation of esters from alcohols and carboxylic acids. Students compare the alternative method from acyl chlorides which go to completion and the ester is the only liquid product

4.9.5 recall the laboratory preparation of a liquid ester from a carboxylic acid and an alcohol;

Students compare the methods of preparing a liquid ester to those of preparing a chloroalkane in



2.5.4

4.9.6 recall the structure of fats and oils as esters of propane-1,2,3-triol (glycerol) and fatty acids;

Students draw the structures of fats produced from given structures of fatty acids and realise the possibility of optical isomers

Prepare a liquid ester from a carboxylic acid and an alcohol




Learning Outcomes


Practical Links


Suggested Time



4.9.7 recall the transesterification reactions of esters with alcohols and carboxylic acids and the use of these reactions to produce biodiesel and margarines;

Students record the definition of transesterification and write equations for transesterifications. They explain why transesterification is carried out

4.9.8 recall that margarines/spreads are produced from the hardening of oils by catalytic hydrogenation using finely divided nickel; and

Students compare samples of oils to samples of margarines using bromine water and a solvent. Outline the differences in properties between margarine and unsaturated oils.



Link the hydrogenation of unsaturated oils to alkenes in section 2.4.5




Learning Outcomes


Practical Links


Suggested Time



4.9.9 recall the reactions of acyl chlorides with water and alcohols.

Students write equations, draw structures and name the organic products




Learning Outcomes


Practical Links


Suggested Time

4.10 Aromatic chemistry


4 hours

4.10.1 explain the structure of the benzene molecule with reference to delocalised π electrons;

Students revise p orbitals and π bonds from 2.4.3 and delocalised electrons from 1.3.10. The structure of the hexagonal benzene ring between delocalised rings of electrons is drawn. Compare Kekulé structure with modern day

CEA Factfile

http://chemguide.co.uk/organicprops/arenesmenu.html#top

http://chemguide.co.uk/mechanisms/elsubmenu.html#top

4.10.2 explain the reactivities of benzene and alkenes related to the relative stabilities of the π electron systems, for example the

Students predict the reactions of the Kekulé structure with bromine and revise alkenes with bromine from 2.4.6 and









resistance of benzene to addition of bromine compared with an alkene;

2.4.8. Benzene doesn’t undergo addition reactions because of the stability of the delocalised π electrons, it undergoes electrophilic substitution reactions


Learning Outcomes


Practical Links


Suggested Time

4.10 Aromatic chemistry (cont.)


4.10.3 explain the mechanisms of the monobromination, mononitration, monoalkylation and monoacylation of benzene including equations for the

Explain that all the mechanisms involve electrophiles i.e. Br+, NO2

+, R+ and RCO+. Students record equations for the formation of the electrophiles and



formation of the electrophile;

realise that their formation uses “catalysts”; draw mechanisms for all of the reactions

4.10.4 recall the names of the electrophiles for bromination and nitration of benzene; and

Explain that the electrophile for bromination is the δ+ end of a polarised bromine molecule and it is the nitronium ion for nitration




Learning Outcomes


Practical Links


Suggested Time

4.10 Aromatic chemistry (cont.)


4.10.5 prepare methyl3-nitrobenzoate from methyl benzoate to illustrate nitration of the benzene ring.

Explain the preparation of the solid methyl 3-nitrobenzoate. Either the preparation is demonstrated or students carry it out. Explain why the temperature is kept under 60oC. Together with the preparation in 4.7.8 revise the purification of organic solids



Unit A2 2:Analytical, Transition

Metals, Electrochemistry and Organic Nitrogen




Unit A2 2: Analytical, Transition Metals, Electrochemistry and Organic Nitrogen Chemistry


Learning Outcomes


Practical Links


Suggested Time

5.1 Mass spectrometry


2 hours

5.1.1 recall the meaning of and identify base peak, molecular ion peak, M+1 peak and fragmentation ions in a mass spectrum;

Students revise mass spectrometry from unit 1 and record the definitions of base peak, molecular ion peak, M+1 peak and fragmentation ions. Students look at a molecular spectrum e.g. ethanol and its features. They

Past paper A2 2June 2014 Q14 b

CCEA Factfile

http://chemguide.co.uk/analysis/masspecmenu.html

www.youtube.com/watch?v=J-wao0O0_qM


http://www.youtube.com/watch?v=J-wao0O0_qM

http://www.youtube.com/watch?v=J-wao0O0_qM





explain the peak at 47, name the base peak, M, M+1, and how M-1 occurs




Learning Outcomes


Practical Links


Suggested Time

5.1 Mass spectrometry (cont.)


5.1.2 suggest formulae for the fragment ions in a given mass spectrum; and

Students suggest the masses of all the fragmentation ions in the ethanol mass spectrum. They look at simple mass spectra of unknown substances and deduce the formulae of the ions that are responsible for peaks

5.1.3 distinguish between molecules of the same RMM/mass using high resolution mass spectrometry.

Students use very accurate RAMs, to several decimal places, to determine RMMs e.g. N2, CO and C2H4.



Students record the definition of high resolution




Learning Outcomes


Practical Links


Suggested Time

5.2 Nuclear magnetic resonance spectroscopy

5.2.1


12 ours

understand the difference between low resolution and high resolution NMR spectra;

Students look at the low resolution and high resolution spectra of ethanol. The hydrogen atoms/ protons absorb energy in the radio wave area of the electromagnetic spectrum. Explain briefly how it works

Past paper A2 2June 2014 Q14June 2013 Q13June 2012 Q13

CCEA Factfile

www.youtube.com/watch?v=uNM801B9Y84

www.nmrdb.org/

http://sdbs.db.aist.go.jp/sdbs/cgi-bin/cre_index.cgi

http://chemguide.co.uk/analysis/nmrmenu.html






http://www.nmrdb.org/

http://www.youtube.com/watch?v=uNM801B9Y84

http://www.youtube.com/watch?v=uNM801B9Y84


comparing the spin of nuclei with the spin of electrons




Learning Outcomes


Practical Links


Suggested Time

5.2 Nuclear magnetic resonance spectroscopy (cont.)

5.2.2

5.2.3


understand the reasons for the use of TMS (tetramethylsilane) as a standard;

Students record the reasons for the use of TMS as a standard i.e. absorbs as a single peak, easily identifiable and well away from proton absorption. Students consider toxicity and volatility as possible reasons

Explain what ‘chemically equivalent hydrogens’ mean. Give simple examples e.g. methane and propanone etc.

recognise chemically equivalent hydrogen atoms;



Students are given structures to analyse equivalent hydrogen atoms




Learning Outcomes


Practical Links


Suggested Time



5.2.4

5.2.5

understand that chemical shifts depend on the chemical environment of hydrogen atoms;

use integration curves to determine the relative number of hydrogen atoms in different chemical environments;

Show students the low resolution NMR spectrum for ethanol explaining the location of the OH proton relative to the TMS peak. Introduce chemical shifts based on the difference of environment i.e. the electronegativity of the adjacent atom the proton is attached to

Students examine



integrated NMR spectra and measure the integrated curves to find the ratio. They then relate the ratio of the hydrogens in the structural formula to the peak integration




Learning Outcomes


Practical Links


Suggested Time



5.2.6

5.2.7

apply the n+1 rule to analyse spin-spin splitting, limited to doublets, triplets and quartets where n is the number of hydrogen atoms on an adjacent carbon atom; and

deduce a molecular structure from an NMR spectrum, limited to simple splitting patterns.

Explain spin-spin splitting, the effect of neighbouring protons on other protons. Introduce the n+1 rule and show examples e.g. the ethyl group CH3CH2- with a quartet and triplet. The n+1 rule is linked to the series 1, 1 1, 1 2 1, 1 3 3 1

Students use the NMR data on the data sheet and splitting patterns to work out the



structures of unknown compounds. This can be done with or without infrared data, mass spectrometry data and chemical data




Learning Outcomes


Practical Links


Suggested Time

5.3 Volumetric analysis


10 hours

5.3.1 titrate iodine with sodium thiosulfate using starch as an indicator and estimate oxidising agents, for example hydrogen peroxide and iodate(V) ions by their reactions with excess acidified potassium iodide;

Students revise good titration technique from unit 1. When to add the starch solution is discussed, together with the colour change. Students react H2O2 and excess KI to liberate I2 which is titrated against standard Na2S2O3 solution. Students calculate the concentration of

Titrate iodine with sodium thiosulfate using starch and hence estimate oxidising agents by their reaction with excess acidified potassium iodide

Titrate acidified potassium manganite(VII)

Past paper A2 2June 2014 Q 13 dJune 2013 Q 12June 2012 Q 14

CCEA Factfile

www.chemguide.co.uk/inorganic/transition/manganese.html


http://www.chemguide.co.uk/inorganic/transition/manganese.html




H2O2. Students carry out titrations using other oxidising agents instead of H2O2

with reducing agents




Learning Outcomes


Practical Links


Suggested Time

5.3 Volumetric analysis (cont.)


5.3.2 titrate acidified potassium manganate(VII) with iron(II) and other reducing agents;

Discuss the colour change of potassium manganate(VII) which is self-indicating. Students titrate iron(II) against potassium manganate(VII). Knowing the concentration of both reactants the titration values obtained are used to work out the molar ratio in the reaction. Students use other reducing agents and calculate the

Determine the purity of a Group II metal oxide or carbonate by back titration



molar ratio of reactants

5.3.3 deduce titration equations given the half-equations for the oxidant and the reductant; and

Students combine the unbalanced ionic half-equations for the reaction of Fe2+ with MnO4

-. Students deduce equations for a range of oxidants and reductants




Learning Outcomes


Practical Links


Suggested Time

5.3 Volumetric analysis (cont.)


5.3.4 understand the method of back titration, for example determine the purity of a Group II metal oxide or carbonate.

Students carry out a back titration to analyse the calcium carbonate in eggshells and the purity of a sample of magnesium oxide. Students carry out calculations for various back titrations




Learning Outcomes


Practical Links


Suggested Time

5.4 Chromatography


3 hours

5.4.1 describe and explain how paper (one-way and two way), thin-layer (TLC) and gas-liquid chromatography (GLC) is carried out qualitatively and quantitatively;

Students revise paper chromatography and its use with coloured compounds/mixtures. Explain the use of 2D paper chromatography with amino acids. Explain how to isolate substances from paper chromatography and TLC using a solvent. Explain how to see

Carry out paper and thin layer chromatography and measure the Rf values of the components and interpret the chromatograms

Past paper A2 2June 2011 Q 11 eJune 2010 Q 16

CCEA Factfile

http://chemguide.co.uk/analysis/chromatogrmenu.html






substances using iodine and UV light




Learning Outcomes


Practical Links


Suggested Time

5.4 Chromatography (cont.)


5.4.2 interpret GLC data in terms of the percentage composition of a mixture; and

Show students examples of output data and that the area under the peaks is proportional to the amount of compound present. Explain that the percentage composition is calculated by the area of one peak divided by the total area of the other peaks




Learning Outcomes


Practical Links


Suggested Time

5.4 Chromatography (cont.)


5.4.3 interpret one-way and two-way paper and TLC chromatograms including calculations of Rf values.

Students use two-way chromatography to give better separation allowing substances with similar Rf values to be distinguished.Students are given a mixture of amino acids, run a chromatogram and compare Rf values to data tables to identify individual amino acids. Students use ninhydrin to develop the location of amino acids




Learning Outcomes


Practical Links


Suggested Time

5.5 Transition metals


12 hours

5.5.1 recall that transition metals or their ions have an incomplete d-shell, variable oxidation states, catalytic activity, and form coloured complexes;

Students define transition metals according to the d-shell. Show students the range of oxidation numbers of transition metals in a table. Students discuss and list the uses of transition metals and their compounds e.g. Haber Process, Contact Process,

Past paper A2 2June 2014 Q 11, Q 13June 2013 Q 12June 2012 Q 14, Q 16

CCEA Factfile

http://chemguide.co.uk/inorganic/transitionmenu.html






Hydrogenation of vegetable fats etc.




Learning Outcomes


Practical Links


Suggested Time

5.5 Transition metals (cont.)


5.5.2

5.5.3

deduce the electronic configuration of transition metals and their ions and explain their stabilities based on the filling of the sub-shells;

understand that complexes consist of a central metal atom or ion surrounded by a number of ligands, defined as anions or molecules possessing lone pairs of electrons;

Students write the electronic configurations of the atoms and ions from scandium to zinc and explain their stabilities

Introduce the term “complex ion” – define as central metal ion surrounded by anions or molecules (ligands). Students record the definition of a complex ion

5.5.4 explain that Students record



ligands are molecules or atoms that contain a lone pair which can be donated to a transition metal atom or ion;

the definition of a ligand and record examples of ligands showing their lone pairs e.g. H2O, NH3, Cl-, CN- etc.




Learning Outcomes


Practical Links


Suggested Time



5.5.5

5.5.6

explain the meaning of and deduce coordination numbers in complexes;

deduce the oxidation number of transition metals in complexes and use them to explain redox and disproportionation reactions;

Students are given examples of the formulae of complexes, work out the coordination number as the number of coordinate bonds from ligands to the central metal ions and name the complexes

Students work out the oxidation numbers of transition metals in complexes in a series of reactions which include redox and disproportionatio



n

5.5.7 understand the distinction between monodentate, for example Cl-, H2O and NH3, and bidentate, for example:

Students discuss examples of ligands and the different structures formed by a variety of monodentate ligands e.g. MA4B2 and the possibility of E/Z isomerism.




Learning Outcomes


Practical Links


Suggested Time



5.5.8 NH2CH2CH2NH2 (represented by en), and polydentate ligands (edta);

explain the relative strengths of ligands in terms of the availability of lone pairs;

Students make models of structures formed with ligands especially bidentate ligands to show the possibility of chiral and non-chiral complexes. Draw the structure of an edta complex showing the six bonding points

Students carry out a series of experiments by adding different ligands to [Cu(H2O)6]2+(aq) and observe the colour changes

Use ethylene diamine, phenylamine and aqueous ammonia to demonstrate ligand replacement based on lone pair availability



5.5.9 understand the ligand replacement reactions of hexaaquacopper (II) ions with concentrated hydrochloric acid and ammonia solution to include colours and shapes of the complexes;

Students look at equations for the ligand replacement of H2O by Cl- and record colours and shapes of the complexes


Learning Outcomes


Practical Links


Suggested Time



5.5.10 explain ligand replacement in terms of positive entropy changes, for example a bidentate ligand displacing two monodentate ligands;

Students record examples of ligand replacement reactions. The reaction is feasible if the products result in more disorder

Demonstrate the relative strengths of ligands using hydrated copper(II) ions and hydrochloric acid



than the reactants e.g. replacing water as a ligand with edta

5.5.11 recall the colours of the aqueous complexes of Cr3+, Cr(VI), Mn2+, Fe2+, Fe3+, Co2+, Ni2+, Cu2+, V2+, V3+, V(IV), V(V);

Students record the colours of the aqueous complexes from the acceptable colours booklet

Carry out qualitative detection tests for the formation of transition metal hydroxides with sodium hydroxide and aqueous ammonia




Learning Outcomes


Practical Links


Suggested Time



5.5.12 use as qualitative detection tests the formation of precipitates of the hydroxides of Cr3+, Mn2+, Fe2+, Fe3+, Co2+, Ni2+, Cu2+ with NaOH(aq) and NH3(aq) and, where appropriate, their subsequent dissolution;

Students perform test tube reactions of the transition metal ions with NH3(aq) and NaOH(aq), note the colour changes and write equations for the reactions

5.5.13 recall the reduction of VO2+ (acidified ammonium metavanadate), by zinc to form VO2+, V3+ and V2+; and

Students look at the colour changes of VO2

+ when reduced using zinc in concentrated hydrochloric acid. The colour

Carry out the reduction of acidified ammonium metavanadate with zinc and observe the sequence of



changes are recorded and equations are written for all the reactions taking placeTo be covered in electrode potentials, 5.6

colours




Learning Outcomes


Practical Links


Suggested Time



5.5.14 deduce, given appropriate emf values, reagents for the interconversion of vanadium between its oxidation state and to combine half-cells to give an overall equation for a reaction.

Students construct a table of electrode potentials which include standard emf values for vanadium ions and reducing agents. Using the “anticlockwise rule” reactions are predicted and equations are written




Learning Outcomes


Practical Links


Suggested Time

5.6 Electrode potentials


8 hours

5.6.1

5.6.2

define standard electrode potential, Eo, and explain the construction and significance of the hydrogen electrode and understand the importance of conditions when measuring electrode potentials;

use standard electrode potentials to predict feasibility

Students record the definition of standard electrode potential and the standard conditions for the performance of a hydrogen electrode

Students record that positive emf’s lead to

Determine the electrode potentials of a series of cells and predict their values using standard electrode potentials

Past paper A2 2June 2014 Q 15June 2013 Q 1June 2012 Q 12

CCEA Factfile

http://alevelchem.com/aqa_a_level_chemistry/unit3.5/s353/04.htm






and direction of reactions, to calculate the emf and understand the limitations of such predictions in terms of concentrations and kinetics;

forward reactions and negative emf’s lead to reverse reaction. emf’s vary depending on concentrations and the presence of acids and alkalis.




Learning Outcomes


Practical Links


Suggested Time

5.6 Electrode potentials (cont.)


Students measure and record the variation of emf with concentration. Although a positive emf suggests that a reaction will occur it might not in terms of kinetics

5.6.3

5.6.4

use conventional representations for cells;

classify cells into non-rechargeable,

Use a similar system to the following and be able to explain all the symbols. Zn(s)|Zn2+(aq) || Cu2+(aq)|Cu(s)

Students record the differences



rechargeable and fuel cells, and state examples of each;

between the different types of cells and describe them, give examples of them explaining how they work




Learning Outcomes


Practical Links


Suggested Time



5.6.5 understand the electrode reactions in a lithium cell;

Students record the following reactions: anode:Li+ + CoO2 + e- → Li+[CoO2]-; cathode: Li → Li+

+ e-




Learning Outcomes


Practical Links


Suggested Time



5.6.6 understand that a fuel cell uses the energy from the reaction of a fuel with oxygen to generate a voltage;

Students write the reactions for a fuel cell at the anode and the cathode e.g.2H2 + 2O2- → 2H2O + 4e- and O2 + 4e- → 2O2-. They compare the calculation of the energy obtained from bond energies to that obtained from a fuel cell. Other fuels used in a fuel cell are examined

5.6.7 recall the electrode reactions that occur in an alkaline hydrogen-

Students record the reactions at the anode and cathode of an alkaline fuel cell;



oxygen fuel cell; and

anode:O2 + 2H2O + 4e- →4OH-, cathode: 2H2O + 4OH- →4H2O + 4e-




Learning Outcomes


Practical Links


Suggested Time



5.6.8 recall the environmental issues associated with cells.

The danger of various heavy metals to human metabolism are recorded by students. The issues of recycling are raised. The dangers of reactive metals such as lithium are recalled




Learning Outcomes


Practical Links


Suggested Time

5.7 Amines Students should be able to:

9 hours

5.7.1 recall the molecular and structural formulae of amines with up to six carbon atoms, and refer to primary, secondary and tertiary amines;

Students revise the amine functional group from unit 2. Show examples of amines from nature e.g. phenylalanine, adrenaline and putrescine etc. Students revise the explanation of primary, secondary and tertiary structures with regard to amines and name a variety of

Past paper A2 2June 2013 Q 11 b, cJune 2012 Q 11

CCEA Factfile

www.chemguide.co.uk/organicprops/aminemenu.html

www.docbrown.info/page07/equilibria7d.htm


http://www.docbrown.info/page07/equilibria7d.htm

http://www.docbrown.info/page07/equilibria7d.htm

http://www.chemguide.co.uk/organicprops/aminemenu.html




structures




Learning Outcomes


Practical Links


Suggested Time

5.7 Amines (cont.)


5.7.2 refer to the effect of hydrogen bonding on boiling point and solubility with water;

Students look at the amine functional group and record the possibilities of hydrogen bonding using the N and H atoms. They apply hydrogen bonding to explain boiling point and solubility and compare methylamine with methanol in terms of boiling points

5.7.3 recall the formation of primary aliphatic amines by reduction of

Students revise the formation of amines from unit 2 i.e. halogenoalkanes



nitriles using lithium tetrahydridoaluminate (III) (lithal) and by the reaction of ammonia with alkyl halides;

with ammonia and note the formation of primary, secondary and tertiary amines. They revise the nitrile functional group from unit 4 and write equations for the reduction with lithal using [H]




Learning Outcomes


Practical Links


Suggested Time

5.7 Amines (cont.)


5.7.4 explain the formation of phenylamine by reduction of nitrobenzene using tin and concentrated hydrochloric acid, to form the phenylammonium salt, followed by liberation of the free amine by addition of alkali;

Students are shown the structure of phenyl amine and the method for its formation using the reduction of nitrobenzene. Equations are writtenfor the reaction of tin with concentrated hydrochloric acid, the reduction of nitrobenzene with [H], the production of the phenylammonium salt and its reaction with alkali to release the amine



5.7.5 recall the formation of salts by the reaction of amines with mineral acids and the liberation of amines from their salts using alkali;

Relate amines to ammonia reacting as a base due to the lone pair on nitrogen. Students write equations for the reaction of bases with acids and the reverse reaction with alkali




Learning Outcomes


Practical Links


Suggested Time

5.7 Amines (cont.)


5.7.6 explain the relative basic strengths of ammonia, primary, secondary, tertiary aliphatic amines and phenylamine using the availability of the lone pair on the nitrogen atom;

Students explore what is meant by a base according to the Lewis theory of electron pair donors and how the availability of electron pairs governs how strong a base is. Explain how alkyl groups increase the electron density of the lone pair and how the electron withdrawing power of the aromatic ring reduces it. Students write equations for the reactions




Learning Outcomes


Practical Links


Suggested Time

5.7 Amines (cont.)


5.7.7 recall the reaction of amines with ethanoyl chloride and use this reaction to identify unknown amines;

Students react ethanoyl chloride with a variety of amines and recrystallize the organic product. They write the equations and determine the melting points of the derivatives. Students compare this method to the similarity with 2, 4-DNPH and the identification of carbonyl compounds




Learning Outcomes


Practical Links


Suggested Time

5.7 Amines (cont.)


5.7.8 explain the formation of benzenediazonium chloride from phenylamine and recall the coupling of diazonium ions with phenol and phenylamine; and

Students look at the formation of nitrous acid by the reaction of sodium nitrite with hydrochloric acid at room temperature and at low temperature. Equations are written for the reaction and the decomposition of nitrous acid. Observations are recorded

Students compare the reaction of butylamine with nitrous acid. Students write equations for the



coupling reactions with phenol and phenylamine and try reactions with substituted phenols and amines




Learning Outcomes


Practical Links


Suggested Time

5.7 Amines (cont.)


5.7.9 explain the colour of compounds such as dyestuffs and indicators based on the extent of delocalisation of electrons leading to the closer proximity of electronic energy levels.

Students examine a variety of structures to see if they are delocalised. Students realise the greater the extent of delocalisation the more the colour moves through the electromagnetic spectrum i.e. from red to blue. They draw the structures of indicators e.g. methyl orange in acidic and alkaline solutions and use delocalisation to explain the change




Learning Outcomes


Practical Links


Suggested Time

5.8 Amides Students should be able to:

3 hours

5.8.1 recall the preparation of amides via the reaction of carboxylic acids with ammonia and the reaction of amines with acyl chlorides;

Introduce the amide functional group and link it back to amine derivatives with ethanoyl chloride.Students name amides from structures

Past paper A2 2June 2014 Q 14 dJune 2013 Q 13 a, dJune 2010 Q 17 b (iv)June 2011 Q 12 f

CCEA Factfile

www.chemguide.co.uk/organicprops/amidemenu.html

5.8.2 recall the hydrolysis of amides with acids and alkalis;

Students write equations for the hydrolysis of amides with acids and alkalis and explain how to


http://www.chemguide.co.uk/organicprops/amidemenu.html




obtain the resulting acid used




Learning Outcomes


Practical Links


Suggested Time

5.8 Amides (cont.)


5.8.3

5.8.4

recall the dehydration of amides with phosphorus pentoxide to form nitriles; and

explain the basicity of amides relative to amines by referring to the delocalisation of the lone pair on the nitrogen atom.

Students write the equation for the formation of a nitrile from the dehydration of an amide.P4O10 is used as a dehydrating agent, the reaction is heated and phosphoric acid is formed

Explain why amides are different from amines in basic character i.e. the C=O π bond overlaps with the lone pair of electrons on nitrogen and is delocalised




Learning Outcomes


Practical Links


Suggested Time

5.9 Amino acids


8 hours

5.9.1

5.9.2

recall the formulae of glycine and alanine;

explain the optical activity of amino acids;

Students look at the structures of the acids and are told they are alpha amino acids i.e. the –NH2 and the–COOH groups are on the same carbon atom. Compare other amino acids to see if they are alpha

Students revise optical activity i.e. four different groups around a carbon atom. All

Past paper A2 2June 2011 Q 11June 2010 Q 17

CCEA Factfile

www.chemguide.co.uk/organicprops/aminoacidmenu.html


http://www.chemguide.co.uk/organicprops/aminoacidmenu.html




alpha amino acids, except glycine, are optically active




Learning Outcomes


Practical Links


Suggested Time

5.9 Amino acids(cont.)


5.9.3

5.9.4

explain the formation of dipolar ions (zwitterions) from amino acid molecules;

explain the solubility of amino acids in water and their relatively high melting point;

Students record the definition of dipolar ions and zwitterions leading to electrostatic attraction between amino acid molecules. which explains why they are solid at room temperatures

Students test glycine’s solubility in water and explain it in terms of the –NH2 and–COOH groups. The very high melting point (compare data) is explained in



terms of the zwitterion structure which is similar to ionic structures

5.9.5 recall the reactions of amino acids with sodium carbonate, copper(II) sulfate and nitrous acid;

Students write equations for the reactions of amino acids with CuSO4, Na2CO3 and HNO2 and make observations




Learning Outcomes


Practical Links


Suggested Time



5.9.6 recall the primary secondary and tertiary structures of a protein and the formation of peptide links;

Show the primary structure, the joining of amino acids by a peptide link eliminating water molecules. Students use the idea of condensation polymerisation, reversible l by hydrolysis with acid or alkali. Students record what is meant by secondary structure and tertiary structure

5.9.7 define enzymes as biological catalysts and use the concept of

Students recap the catalyst definition from unit 2 and apply



‘induced fit’ to explain enzyme action;

it to enzymes. Explain the concept of ‘induced fit’ which is similar to the active sites on enzymes specific to substrate molecules e.g. amylase for starch etc.




Learning Outcomes


Practical Links


Suggested Time



5.9.8 recall the use of enzymes in washing powders and their economic advantage in operating at lower temperatures;

Students recall that enzymes remove “biological stains” which consist of organic molecules. They operate best at room temperature/ 37ºC.

Lower temperatures mean less energy and less cost

5.9.9 explain that an enzyme is a protein and provides a path of lower activation energy; and

Students record that enzymes are special structures of proteins that contain “active sites” into which substrate



molecules fit and revise kinetics regarding the lowering of activation energies by enzymes




Learning Outcomes


Practical Links


Suggested Time



5.9.10 explain the effect of pH and temperature on enzyme activity.

Discuss the effect of pH changes and temperature on enzyme action and explain denaturing i.e. active site shape is changed. Students record graphs of pH and temperature against enzyme activity




Learning Outcomes


Practical Links


Suggested Time

5.10 Polymer chemistry


6 hours

5.10.1 understand that condensation polymers are formed from molecules containing COOH, OH and NH2 groups and be able to draw polymer structures from monomers and vice versa;

Students revise polymer formation from unit 1 i.e. alkenes. They research the possibilities for the formation of polymers from molecules containing–COOH, –OH and–NH2 using a variety of structures. They write equations for their formation

Past paper A2 2June 2014 Q 14 cJune 2013 Q 5, 13 eJune 2012 Q 17

CCEA Factfile

www.chemguide.co.uk/organicprops/alkenes/polymerisation.html#top

www.chemguide.co.uk/organicprops/esters/polyesters.html


http://www.chemguide.co.uk/organicprops/esters/polyesters.html



http://www.chemguide.co.uk/organicprops/alkenes/polymerisation.html#top





Learning Outcomes


Practical Links


Suggested Time

5.10 Polymer chemistry (cont.)


5.10.2 understand the formation, structure and uses of the polyester polyethylene terephthalate (PET);

Students revise ester formation and compare it to polyester formation. They consider the possibility of polymers formed from the isomers of terephthalic acid. Students record the uses and advantages of PET and relate them to structure

5.10.3 understand the formation, structure and uses of the polyamide, nylon; and

Students record the structures of nylons such as nylon 66 and the monomers from which they are formed.



They compare polyamides with polypeptides. They record the many uses of nylon based on its strength and durability




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5.10 Polymer chemistry (cont.)


5.10.4 recall that polyesters and polyamides can be hydrolysed and are therefore biodegradable.

Students look at other condensation polymers such as Kevlar which has an amide bond. Relate holes in nylon/polyester when acid is spilled on thin cloth. Hydrolysis of polymers in the laboratory with acid and alkali and biodegradability with bacteria




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5.11 Chemistry in medicine


6 hours

5.11.1

5.11.2

recall the use of digestion remedies, for example hydroxides and carbonates, to cure excess acid in the stomach (link with Section 1.9);

use a back titration to determine the percentage of an active ingredient in an indigestion remedy (link with Section 5.3.4);

Students titrate acid with insoluble alkalis and compare with soluble alkalis using methyl orange and phenolphthalein. Add indigestion tablets to acid and measure the pH with a pH meter

Students practise the back titration of a metal carbonate with hydrochloric acid and standard sodium hydroxide and then use indigestion

Determine the amount of a carbonate, for example calcium carbonate or magnesium carbonate in an indigestion tablet

An introduction to Medicinal Chemistry, Graham L. Patrick



remedies




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5.11 Chemistry in medicine (cont.)


5.11.3

5.11.4

recall methods to deal with excessive pH values of skin and explain the use of corrosive chemicals in removing warts (link with Sections 1.9 and 2.11.5);

recall and explain the use of silver nitrate in the treatment of eye diseases;

Explain the action of remedies such as pHisoderm to control the pH of skin. Explain why skin is acidic and name the compounds produced by sweat. Students record the chemicals used in treating warts and why others are not used

Students record the caustic nature of silver nitrate. Explain why a solution of silver nitrate is



used to treat eyes but the solid is used for other medical purposes. Students note the developments of nano-technology in medicine with silver




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5.11.5

5.11.6

explain the action of anticancer drugs, for example cisplatin in preventing DNA replication in cancer cells and how varying the structure of cisplatin affects the efficiency of anticancer activity;

use volumetric analysis to determine the concentration of aspirin in solution (link with Section

Explain cis/trans is now E/Z and how cisplatin relates to this system. Draw the structures of cisplatin and transplatin and explain why cisplatin prevents DNA replication. Studentslook at the variety of structural variations of cisplatin and why their activity variesStudents titrate solutions of aspirin tablets and relate the

Prepare aspirin using salicylic acid and ethanoic anhydride

Use chromatography to compare the purity of

www.youtube.com/watch?v=PTH7xOKY_Co

teachers.sduhsd.net/lcale/Aspirin%20Titration%20Lab.doc

CCEA practical video


http://www.teachers.sduhsd.net/

http://www.youtube.com/watch?v=PTH7xOKY_Co

http://www.youtube.com/watch?v=PTH7xOKY_Co


5.3); results to the composition of the tablets. The equation for the reaction is recorded and the choice of indicator explained. The effect of any other chemicals in the tablet is noted. The cost per mg/g of aspirin is compared

laboratory-made aspirin with commercial tablets




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5.11.7 recall the synthesis of aspirin from salicylic acid using ethanoic anhydride and reasons for its use as a sodium salt;

Students research the numerous articles on the synthesis of aspirin and compare them explaining the reasons for the various techniques and reaction times. Students carry out solubility tests on a variety of aspirin tablets and “soluble” aspirin

5.11.8 use TLC and GLC MS to identify drugs and their purity (link with

Students revise TLC, GLC and MS. They use TLC to analyse



Section 5.4); their aspirin preparations and compare them with commercial samples. Explain that GLC uses specific columns for drug analysis




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5.11.9 explain the role of iron(II) in haemoglobin in the transportation of oxygen in blood and the poisonous nature of carbon monoxide (link with Section 2.3.5); and

Explain that iron(II) is surrounded by a porphyrin ring which is conjugated leading to the red colour of haemoglobin and hence blood. The iron (II) combines with oxygen but more strongly with carbon monoxide. Both are equilibrium reactions. Treatment for carbon monoxide poisoning is to use excess oxygen and move the equilibrium



backwards




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5.11.10 explain the role of edta in sequestering calcium ions and thus preventing the clotting of blood (link with Section 5.5.7).

Students record the meaning of the term sequester. They record that calcium ions are used in the clotting of blood, hence removing Ca2+ with edta prevents clotting.

Students draw the structure of the calcium/edta complex. They note that edta sequesters most divalent cations

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