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Chemistry

Chemistry YCH11 — specification explorer

All 319 specification statements from the official Pearson Edexcel specification, as printed — filter by paper, unit or tier exactly as Pearson Edexcel assesses them.

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319 spec points

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Assessment homes come from the parsed Pearson specification (content summaries & assessment overviews). Hover a chip for the printed rule. Filters appear only where the qualification has that structure — a linear IGCSE has no unit scope, and only Maths A splits Foundation / Higher tiers.

1. Formulae, Equations and Amount of Substance

1.1know the terms ‘atom', 'element', 'ion', 'molecule', 'compound', 'empirical formula' and 'molecular formula’Unit 1
1.2know that the mole (mol) is the unit for the amount of a substance and be able to perform calculations using the Avogadro constant L (6.02 x 1023 mol-1)Unit 1
1.3write balanced full and ionic equations, including state symbols, for chemical reactionsUnit 1
1.4understand the terms: i ‘relative atomic mass’ based on the 12 C scale ii ‘relative molecular mass’ and ‘relative formula mass’, including calculating these values from relative atomic masses The term ‘relative formula mass’ should be used for compounds with giant structures. iii ‘molar mass’ as the mass per mole of a substance in g mol -1 iv parts per million (ppm), including gases in the atmosphereUnit 1
1.5calculate the concentration of a solution in mol dm −3 and g dm −3 Titration calculations are not required at this stage.Unit 1
1.6be able to use experimental data to calculate empirical and molecular formulaeUnit 1
1.7be able to use chemical equations to calculate reacting masses and vice versa, using the concepts of amount of substance and molar massUnit 1
1.8be able to use chemical equations to calculate volumes of gases and vice versa, using: i the concepts of amount of substance ii the molar volume of gases iii the expression pV = n RT for gases and volatile liquidsUnit 1
1.9be able to calculate percentage yields and percentage atom economies (by mass) in laboratory and industrial processes, using chemical equations and experimental results Atom economy = molar mass of the desired product × 100% sum of the molar masses of all productsUnit 1
1.10be able to determine a formula or confirm an equation by experiment, including evaluation of the dataUnit 1
1.11CORE PRACTICAL 1 Measurement of the molar volume of a gas.Unit 1
1.12be able to relate ionic and full equations, with state symbols, to observations from simple test-tube experiments, to include: i displacement reactions ii typical reactions of acids iii precipitation reactionsUnit 1

2. Atomic Structure and the Periodic Table

1. Atomic Structure and the Periodic Table

2.1know the structure of an atom in terms of electrons, protons and neutronsUnit 1
2.2know the relative mass and charge of protons, neutrons and electronsUnit 1
2.3know what is meant by the terms ‘atomic (proton) number’ and ‘mass number’Unit 1
2.4be able to use the atomic number and the mass number to determine the number of each type of subatomic particle in an atom or ionUnit 1
2.5understand the term ‘isotope’Unit 1
2.6understand the basic principles of a mass spectrometer and be able to analyse and interpret mass spectra to: i deduce the isotopic composition of a sample of an element ii calculate the relative atomic mass of an element from relative abundances of isotopes and vice versa iii determine the relative molecular mass of a molecule, and hence identify molecules in a sample iv understand that ions in a mass spectrometer may have a 2+ chargeUnit 1
2.7be able to predict mass spectra, including relative peak heights, for diatomic molecules, including chlorine, given the isotopic abundancesUnit 1
2.8be able to define first, second and third ionisation energies and understand that all ionisation energies are endothermicUnit 1
2.9know that an orbital is a region within an atom that can hold up to two electrons with opposite spinsUnit 1
2.10understand how ionisation energies are influenced by the number of protons in the nucleus, the electron shielding and the sub-shell from which the electron is removedUnit 1
2.11know that ideas about electronic structure developed from: i an understanding that successive ionisation energies provide evidence for the existence of quantum shells and the group to which the element belongs ii an understanding that the first ionisation energy of successive elements provides evidence for electron sub-shellsUnit 1
2.12be able to describe the shapes of s and p orbitalsUnit 1
2.13know that orbitals in sub-shells: i each take a single electron before pairing up ii pair up with two electrons of opposite spinUnit 1
2.14be able to predict the electronic configuration of atoms of the elements from hydrogen to krypton inclusive and their ions, using s, p, d notation and electron-in- boxes notationUnit 1
2.15understand that electronic configuration determines the chemical properties of an elementUnit 1
2.16know that the Periodic Table is divided into blocks, such as s, p and d, and know the number of electrons that can occupy s, p and d sub-shells in the first four quantum shellsUnit 1
2.17be able to represent data, in a graphical form (including the use of logarithms of first ionisation energies on a graph) for elements 1 to 36 and hence explain the meaning of the term ‘periodic property’Unit 1
2.18be able to explain: i the trends in melting and boiling temperatures of the elements of Periods 2 and 3 of the Periodic Table in terms of the structure of the element and the bonding between its atoms or molecules ii the general increase and the specific trends in ionisation energy of the elements across Periods 2 and 3 of the Periodic Table iii the decrease in first ionisation energy down a groupUnit 1

3. Bonding and Structure

1. Bonding and Structure

3.1know and be able to interpret evidence for the existence of ions, limited to physical properties of ionic compounds, electron density maps and the migration of ionsUnit 1
3.2be able to describe the formation of ions in terms of loss or gain of electronsUnit 1
3.3be able to draw dot-and-cross diagrams to show electrons in cations and anionsUnit 1
3.4be able to describe ionic crystals as giant lattices of ionsUnit 1
3.5know that ionic bonding is the result of strong net electrostatic attraction between ionsUnit 1
3.6understand the effects of ionic radius and ionic charge on the strength of ionic bondingUnit 1
3.7understand reasons for the trends in ionic radii down a group in the Periodic Table, and for a set of isoelectronic ions, including N 3− to Al 3+Unit 1
3.8understand the meaning of the term ‘polarisation’ as applied to ionsUnit 1
3.9understand that the polarising power of a cation depends on its radius and charge, and the polarisability of an anion also depends on its radius and charge Further suggested practical The migration of ions in a U-tube using copper(II) chromate solution or on a microscope slide using potassium manganate(VII) crystalsUnit 1
3.10understand that covalent bonding is the strong electrostatic attraction between two nuclei and the shared pair of electrons between them, based on the evidence: i the physical properties of giant atomic structures ii electron density maps for simple moleculesUnit 1
3.11be able to draw dot-and-cross diagrams to show electrons in covalent substances, including: i molecules with single, double and triple bonds ii species with dative covalent (coordinate) bonds, including Al 2 Cl 6 and the ammonium ionUnit 1
3.12be able to describe the different structures formed by giant lattices of carbon atoms, including graphite, diamond and graphene, and discuss the applications of eachUnit 1
3.13understand the meaning of the term ‘electronegativity’ as applied to atoms in a covalent bondUnit 1
3.14know that ionic and covalent bonding are the extremes of a continuum of bonding type and be able to explain this in terms of electronegativity differences, leading to bond polarity in bonds and molecules, and to ionic bonding if the electronegativity is large enoughUnit 1
3.15be able to distinguish between polar bonds and polar molecules and predict whether or not a given molecule is likely to be polar Further suggested practical Determine the effect of an electrostatic force on jets of liquids (water, ethanol and cyclohexane) and use the results to determine whether the molecules are polar or non-polarUnit 1
3.16understand the principles of the electron-pair repulsion theory, used to interpret and predict the shapes of simple molecules and ionsUnit 1
3.17understand the terms ‘bond length’ and ‘bond angle’Unit 1
3.18know and be able to explain the shapes of, and bond angles in, BeCl 2 , BCl 3 , CH 4 , NH 3 , NH + , H O, CO , gaseous PCl , SF and C H 4 2 2 5 6 2 4Unit 1
3.19be able to apply the electron-pair repulsion theory to predict the shapes of, and bond angles in, molecules and ions analogous to those in 3.18Unit 1
3.20understand that metals consist of giant lattices of metal ions in a sea of delocalised electronsUnit 1
3.21know that metallic bonding is the strong electrostatic attraction between metal ions and the delocalised electronsUnit 1
3.22be able to use the models in 3.20 and 3.21 to interpret simple properties of metals, including electrical conductivity and high melting temperatureUnit 1

4. Introductory Organic Chemistry and Alkanes

1. Introductory Organic Chemistry and Alkanes

4.1understand the difference between hazard and riskUnit 1
4.2understand the hazards associated with organic compounds and why it is necessary to carry out risk assessments when dealing with potentially hazardous materialsUnit 1
4.3be able to suggest ways in which risks can be reduced and reactions carried out safely, for example: i working on a smaller scale ii taking precautions specific to the hazard iii using an alternative method that involves less hazardous substancesUnit 1
4.4understand the concepts of homologous series and functional groupUnit 1
4.5be able to apply the rules of International Union of Pure and Applied Chemistry (IUPAC) nomenclature to: i name compounds relevant to this specification ii draw these compounds, as they are encountered in the specification, using structural, displayed and skeletal formulae Students will be expected to know prefixes for compounds up to C 10Unit 1
4.6be able to classify reactions as addition, substitution, oxidation, reduction or polymerisationUnit 1
4.7understand that bond breaking can be: i homolytic, to produce free radicals ii heterolytic, to produce ionsUnit 1
4.8know definitions of the terms ‘free radical’ and ‘electrophile’Unit 1
4.9know the general formula of alkanes and cycloalkanes, and understand that they are hydrocarbons (compounds of carbon and hydrogen only) which are saturated (contain single bonds only)Unit 1
4.10understand the term ‘structural isomerism’ and be able to draw the structural isomers of organic molecules, given their molecular formulaUnit 1
4.11be able to draw and name the structural isomers of alkanes and cycloalkanes with up to six carbon atomsUnit 1
4.12know that alkanes are used as fuels and obtained from the fractional distillation, cracking and reforming of crude oil, and be able to write equations for these reactionsUnit 1
4.13know that pollutants, including carbon monoxide, oxides of nitrogen and sulfur, carbon particulates and unburned hydrocarbons, are emitted during the combustion of alkane fuelsUnit 1
4.14understand the problems arising from pollutants from the combustion of alkane fuels, limited to the toxicity of carbon monoxide and why it is toxic, and the acidity of oxides of nitrogen and sulfurUnit 1
4.15be able to discuss the reasons for developing alternative fuels in terms of sustainability and reducing emissions, including the emission of CO 2 and its relationship to climate changeUnit 1
4.16be able to apply the concept of carbon neutrality to different fuels, such as petrol, bioethanol and hydrogenUnit 1
4.17understand the reactions of alkanes with: i oxygen in the air (combustion) ii halogensUnit 1
4.18understand the mechanism of the free radical substitution reaction between an alkane and a halogen: i using free radicals, which are species with an unpaired electron, represented by a single dot ii showing the initiation step of the mechanism, with curly half-arrows for free radical formation iii showing the propagation and termination steps of the mechanism iv having limited use in synthesis because of further substitution reactions Further suggested practical Cracking alkanes by thermal decomposition, including liquid paraffin using aluminium oxide as a catalystUnit 1

5. Alkenes

1. Alkenes

5.1know the general formula of alkenes and understand that alkenes and cycloalkenes are hydrocarbons which are unsaturated (have a carbon-carbon double bond which consists of a σ bond and a π bond)Unit 1
5.2be able to explain geometric isomerism in terms of restricted rotation around a C = C double bond and the nature of the substituents on the carbon atomsUnit 1
5.3understand the E – Z naming system for geometric isomers and why it is necessary to use this when the cis- and trans - naming system breaks downUnit 1
5.4be able to describe the reactions of alkenes, limited to: i the addition of hydrogen, using a nickel catalyst, to form an alkane ii the addition of halogens to produce a di-substituted halogenoalkane iii the addition of hydrogen halides to produce mono-substituted halogenoalkanes iv the addition of steam, in the presence of an acid catalyst, to produce alcohols v oxidation of the double bond by acidified potassium manganate(VII) to produce a diolUnit 1
5.5know the qualitative test for a C=C double bond using bromine or bromine waterUnit 1
5.6be able to describe the mechanism (including diagrams), giving evidence where possible, of: i the electrophilic addition of bromine and hydrogen bromide to ethene ii the electrophilic addition of hydrogen bromide to propene Use of the curly arrow notation is expected – the curly arrows should start from either a bond or from a lone pair of electrons. Knowledge of the relative stability of primary, secondary and tertiary carbocation intermediates is expected.Unit 1
5.7be able to describe the addition polymerisation of alkenes and draw the repeat unit given the monomer, and vice versaUnit 1
5.8understand how chemists limit the problems caused by polymer disposal by: i developing biodegradable polymers ii removing toxic waste gases produced by the incineration of polymers Further suggested practicals: i investigating the difference in reactivity of alkanes and alkenes, including combustion, reaction with bromine water, reaction with acidified potassium manganate(VII) ii preparation of cyclohexene from cyclohexanol iii preparation of limonene from orange peel by steam distillation iv preparation of Perspex ® from methyl 2-methylpropenoateUnit 1

6. Energetics

1. Energetics

6.1know that the enthalpy change, ∆ H , is the heat energy change measured at constant pressure and that standard conditions are 100 kPa and a specified temperature, usually 298 KUnit 2
6.2know that, by convention, exothermic reactions have a negative enthalpy change and endothermic reactions have a positive enthalpy changeUnit 2
6.3be able to construct and interpret enthalpy level diagrams, showing exothermic and endothermic enthalpy changesUnit 2
6.4know the definition of standard enthalpy change of: i reaction, ∆ r H ii formation, ∆ f H iii combustion, ∆ c H iv neutralisation, ∆ neut H v atomisation, ∆ at HUnit 2
6.5be able to use experimental data to calculate: i energy transferred in a reaction recalling and using the expression: energy transferred (J) = mass (g) × specific heat capacity (J g-1 °C-1) × temperature change (°C) ii enthalpy change of the reaction in kJ mol-1 This will be limited to experiments where substances are mixed in an insulated container and combustion experiments using a suitable calorimeter.Unit 2
6.6know Hess’s Law and be able to apply it to: i constructing enthalpy cycles ii calculating enthalpy changes of reaction using data provided, or data selected from a table or obtained from experimentsUnit 2
6.7CORE PRACTICAL 2 Determination of the enthalpy change of a reaction using Hess’s Law.Unit 2
6.8be able to evaluate the results obtained from experiments and comment on sources of error and uncertainty and any assumptions made in the experiments Students will need to consider experiments where substances are mixed in an insulated container and combustion experiments using, for example, a spirit burner and be able to draw suitable graphs and use cooling curve corrections.Unit 2
6.9understand the terms ‘bond enthalpy’ and ‘mean bond enthalpy’, and be able to use bond enthalpies to calculate enthalpy changes, understanding the limitations of this methodUnit 2
6.10be able to calculate mean bond enthalpies from enthalpy changes of reactionUnit 2
6.11understand that bond enthalpy data gives some indication about which bond will break first in a reaction, how easy or difficult it is and therefore how rapidly a reaction will take place at room temperatureUnit 2

7. Intermolecular Forces

1. Intermolecular Forces

7.1understand the nature of the following intermolecular forces: i London forces (instantaneous dipole-induced dipole) ii permanent dipole-permanent dipole interactions iii hydrogen bondsUnit 2
7.2understand the interactions in molecules, such as H 2 O, liquid NH 3 and liquid HF, which give rise to hydrogen bondingUnit 2
7.3understand the following anomalous properties of water resulting from hydrogen bonding: i its high melting and boiling temperature when compared with similar molecules ii the density of ice compared to that of waterUnit 2
7.4be able to predict the presence of hydrogen bonding in molecules analogous to those mentioned in 7.2Unit 2
7.5understand, in terms of intermolecular forces, physical properties shown by substances, including: i the trends in boiling temperatures of alkanes with increasing chain length ii the effect of branching in the carbon chain on the boiling temperatures of alkanes iii the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons iv the trends in boiling temperatures of the hydrogen halides HF to HIUnit 2
7.6understand factors that influence the choice of solvents, including: i water, to dissolve some ionic compounds, in terms of the hydration of the ions ii water, to dissolve simple alcohols, in terms of hydrogen bonding iii water, as a poor solvent for compounds (to include polar molecules such as halogenoalkane), in terms of inability to form hydrogen bonds iv non-aqueous solvents, for compounds that have similar intermolecular forces to those in the solvent Further suggested practicals: i the solubility of simple molecules in different solvents ii measuring the enthalpy change of vaporisation of water iii measuring temperature changes when substances dissolveUnit 2

8. Redox Chemistry and Groups 1, 2 and

1. Redox Chemistry and Groups 1, 2 and

8.1know what is meant by the term ‘oxidation number’ and understand the rules for assigning oxidation numbersUnit 2
8.2be able to calculate the oxidation number of elements in compounds and ions, including in peroxides and metal hydridesUnit 2
8.3be able to indicate the oxidation number of an element in a compound or an ion, using a Roman numeralUnit 2
8.4be able to write formulae given oxidation numbersUnit 2
8.5understand oxidation and reduction in terms of electron transfer and changes in oxidation number, and the application of these ideas to reactions of s- block and p- block elementsUnit 2
8.6know that oxidising agents gain electrons and reducing agents lose electronsUnit 2
8.7understand that a disproportionation reaction involves an element in a single species being simultaneously oxidised and reducedUnit 2
8.8know that oxidation number is a useful concept in terms of the classification of reactions as redox and as disproportionationUnit 2
8.9understand that metals, in general, form positive ions by loss of electrons with an increase in oxidation number whereas non-metals, in general, form negative ions by gain of electrons with a decrease in oxidation numberUnit 2
8.10be able to write ionic half-equations and use them to construct full ionic equationsUnit 2
8.11understand reasons for the trend in ionisation energy down Groups 1 and 2Unit 2
8.12understand reasons for the trend in reactivity of the elements down Group 1 (Li to K) and Group 2 (Mg to Ba)Unit 2
8.13know the reactions of the elements of Group 1 (Li to K) and Group 2 (Mg to Ba) with oxygen, chlorine and waterUnit 2
8.14know the reactions of: i oxides of Group 1 and 2 elements with water and dilute acid ii hydroxides of Group 1 and 2 elements with dilute acidUnit 2
8.15know the trends in solubility of the hydroxides and sulfates of Group 2 elementsUnit 2
8.16understand the reasons for the trends in thermal stability of the nitrates and the carbonates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involvedUnit 2
8.17understand the formation of characteristic flame colours by Group 1 and 2 compounds in terms of electron transitions Students will be expected to know the flame colours for Group 1 and 2 compounds.Unit 2
8.18know experimental procedures to show: i patterns in the thermal decomposition of Group 1 and 2 nitrates and carbonates Students will be expected to know tests for carbon dioxide and oxygen; and to recognise nitrogen dioxide by its colour and acidic pH. ii flame colours in compounds of Group 1 and 2 elementsUnit 2
8.19know reactions, including ionic equations where appropriate, for identifying: i carbonate ions, CO 3 2- , and hydrogencarbonate ions, HCO - 3 , using an aqueous acid to form carbon dioxide (and testing the gas with limewater) ii sulfate ions, SO 2- 4 , using acidified barium chloride solution iii ammonium ions, NH + 4 , using sodium hydroxide solution and warming to form ammonia (and testing with litmus and HCl fumes)Unit 2
8.20be able to calculate solution concentrations, in mol dm -3 and g dm -3 , including simple acid-base titrations using the indicators methyl orange and phenolphthaleinUnit 2
8.21CORE PRACTICAL 3 Finding the concentration of a solution of hydrochloric acid.Unit 2
8.22understand how to minimise the sources of measurement uncertainty in volumetric analysis and estimate the overall uncertainty in the calculated resultUnit 2
8.23CORE PRACTICAL 4 Preparation of a standard solution from a solid acid and use it to find the concentration of a solution of sodium hydroxide. Further suggested practicals: i experiments to study the thermal decomposition of Group 1 and 2 nitrates and carbonates ii flame tests on compounds of Group 1 and 2 iii simple acid-base titrations using the indicators methyl orange and phenolphthalein to calculate solution concentrations in g dm -3 and mol dm -3 iv the solubility of calcium hydroxide by titration v determination of moles of water of crystallisation by titrationUnit 2
8.24understand reasons for the trends for Group 7 elements in: i melting and boiling temperatures and physical state at room temperature ii electronegativity iii reactivity down the groupUnit 2
8.25understand the trend in reactivity of Group 7 elements in terms of the redox reactions of Cl 2 , Br 2 and I 2 with halide ions in aqueous solution Students are expected to know the colours of the elements in standard conditions, in aqueous solution and in a non-polar organic solvent.Unit 2
8.26understand, in terms of changes in oxidation number, the following reactions of the halogens: i oxidation reactions with Group 1 and 2 metals ii the disproportionation reaction of chlorine with water and the use of chlorine in water treatment iii the disproportionation reaction of chlorine with cold, dilute aqueous sodium hydroxide to form bleach iv the disproportionation reaction of chlorine with hot alkali v reactions analogous to those specified aboveUnit 2
8.27understand the following reactions: i solid Group 1 halides with concentrated sulfuric acid, to illustrate the trend in reducing ability of the hydrogen halides ii precipitation reactions of the aqueous anions Cl - , Br - and I - with aqueous silver nitrate solution and nitric acid, and the solubility of the precipitates in aqueous ammonia solution iii hydrogen halides with ammonia gas (to produce ammonium halides) and with water (to produce acids)Unit 2
8.28be able to make predictions about fluorine and astatine and their compounds, in terms of knowledge of trends in halogen chemistry Further suggested practicals: i reaction of solid potassium halides with concentrated sulfuric acid ii precipitation reaction for halides and other anionsUnit 2

9. Introduction to Kinetics and Equilibria

1. Introduction to Kinetics and Equilibria

9.1understand, in terms of the collision theory, the effect of changes in concentration, temperature, pressure and surface area on the rate of a chemical reactionUnit 2
9.2understand that reactions take place only when collisions have sufficient energy, known as the activation energyUnit 2
9.3be able to calculate the rate of a reaction from: i the time taken for a reaction, using rate = 1/time ii the gradient of suitable graph, by drawing a tangent, either for initial rate, or at a time, tUnit 2
9.4understand qualitatively, in terms of the Maxwell-Boltzmann distribution of molecular energies, how changes in temperature affect the rate of a reactionUnit 2
9.5understand the role of catalysts in providing alternative reaction routes of lower activation energyUnit 2
9.6be able to draw the reaction profiles for uncatalysed and catalysed reactions, including the energy level of the intermediate formed with the catalystUnit 2
9.7understand the use of catalysts in industry to make processes more sustainable by using less energy and/or higher atom economyUnit 2
9.8be able to interpret the action of a catalyst in terms of a qualitative understanding of the Maxwell-Boltzmann distribution of molecular energies Further suggested practical Experiments to demonstrate the factors that influence the rate of chemical reactions, including the decomposition of hydrogen peroxide, reaction of marble chips with acid, reaction of thiosulfate ions with acidUnit 2
9.9know that many reactions are readily reversible and that they can reach a state of dynamic equilibrium in which: i the rate of the forward reaction is equal to the rate of the backward reaction ii the concentrations of the reactants and the products remain constantUnit 2
9.10be able to predict and justify the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium in a homogeneous systemUnit 2
9.11evaluate data to explain the necessity, for many industrial processes, to reach a compromise between the yield and the rate of reactionUnit 2

10. Organic Chemistry: Halogenoalkanes, Alcohols

1. Organic Chemistry: Halogenoalkanes, Alcohols

10.1be able to classify reactions (including those in Unit 1) as addition, elimination, substitution, oxidation, reduction, hydrolysis or polymerisationUnit 2
10.2understand the concept of a reaction mechanismUnit 2
10.3understand that heterolytic bond breaking results in species that are electrophiles or nucleophilesUnit 2
10.4know the definition of the term ‘nucleophile’Unit 2
10.5understand the link between bond polarity and the type of reaction mechanism a compound will undergoUnit 2
10.6understand the nomenclature of halogenoalkanes and be able to draw their structural, displayed and skeletal formulaeUnit 2
10.7understand the distinction between primary, secondary and tertiary halogenoalkanesUnit 2
10.8understand the reactions of halogenoalkanes with: i aqueous alkali, including KOH(aq) to produce alcohols (where the hydroxide ion acts as a nucleophile) ii ethanolic potassium hydroxide to produce alkenes by an elimination reaction (where the hydroxide ion acts as a base) iii aqueous silver nitrate in ethanol (where water acts as a nucleophile) iv alcoholic ammonia under pressure to produce amines (where the ammonia acts as a nucleophile) v alcoholic potassium cyanide to produce nitriles (where the cyanide ion acts as a nucleophile) Students should know this is an example of increasing the length of the carbon chain .Unit 2
10.9understand the mechanisms of the nucleophilic substitution reactions between primary halogenoalkanes and: i aqueous potassium hydroxide ii ammonia S N 1 and S N 2 substitution mechanisms will be tested in Unit 4.Unit 2
10.10understand that experimental observations and data can be used to compare the relative rates of hydrolysis of: i primary, secondary and tertiary structural isomers of a halogenoalkane ii primary chloro-, bromo- and iodoalkanes using aqueous silver nitrate in ethanolUnit 2
10.11CORE PRACTICAL 5 Investigation of the rates of hydrolysis of some halogenoalkanes.Unit 2
10.12know the trend in reactivity of primary, secondary and tertiary halogenoalkanesUnit 2
10.13understand, in terms of bond enthalpy, the trend in reactivity of chloro-, bromo- and iodoalkanesUnit 2
10.14CORE PRACTICAL 6 Chlorination of 2-methylpropan-2-ol with concentrated hydrochloric acid. Further suggested practicals: i the use of silver nitrate solution to identify the halogen present in halogenoalkanes ii preparation of 1-bromobutane from butan-1-ol, potassium bromide and sulfuric acidUnit 2
10.15understand the nomenclature of alcohols and be able to draw their structural, displayed and skeletal formulaeUnit 2
10.16understand the distinction between primary, secondary and tertiary alcoholsUnit 2
10.17
understand the reactions of alcohols with: i oxygen in air (combustion) ii halogenating agents:
- PCl 5 to produce chloroalkanes (including its use as a qualitative test for the presence of the –OH group)
- 50% concentrated sulfuric acid and potassium bromide to produce bromoalkanes
- red phosphorus and iodine to produce iodoalkanes iii concentrated phosphoric acid to form alkenes by elimination Descriptions of the mechanisms of these reactions are not required.
Unit 2
10.18understand that potassium dichromate(VI) in dilute sulfuric acid can oxidise: i primary alcohols to produce aldehydes (which give a positive result with Benedict’s or Fehling’s solution) if the product is distilled as it forms ii primary alcohols to produce carboxylic acids (which give a positive result with sodium carbonate or sodium hydrogencarbonate) if the reagents are heated under reflux iii secondary alcohols to produce ketones In equations, the oxidising agent can be represented by [O].Unit 2
10.19understand, the following techniques in the preparation and purification of a liquid organic compound: i heating under reflux ii extraction with a solvent using a separating funnel iii distillation iv drying with an anhydrous salt v boiling temperature determinationUnit 2
10.20CORE PRACTICAL 7 The oxidation of propan-1-ol to produce propanal and propanoic acid. Further suggested practical: Investigation of reactions of primary and secondary alcohols, including propan-1-ol and propan-2-olUnit 2
10.21be able to interpret data from mass spectra to suggest possible structures of simple organic compounds using the m/z of the molecular ion and fragmentation patternsUnit 2
10.22be able to use infrared spectra, or data from infrared spectra, to deduce functional groups present in organic compounds, and predict infrared absorptions, given wavenumber data, due to familiar functional groups including: i C–H stretching absorptions in alkanes, alkenes and aldehydes ii C=C stretching absorption in alkenes iii O–H stretching absorptions in alcohols and carboxylic acids iv C=O stretching absorptions in aldehydes, ketones and carboxylic acids v C–X stretching absorption in halogenoalkanes vi N-H stretching absorption in aminesUnit 2
10.23CORE PRACTICAL 8 Analysis of some inorganic and organic unknowns.Unit 2

11. Kinetics

1. Kinetics

11.1understand the terms: i rate of reaction ii rate equation, rate= k [A] m [B] n where m and n are 0, 1 or 2 iii order with respect to a substance in a rate equation iv overall order of a reaction v rate constant vi half-life vii rate-determining step viii activation energy ix heterogeneous and homogeneous catalystUnit 4
11.2be able to calculate the half-life of a reaction, using data from a suitable graph, and identify a reaction with a constant half-life as being first orderUnit 4
11.3be able to select and justify a suitable experimental technique to obtain rate data for a given reaction, including: i titration ii colorimetry iii mass change iv volume of gas evolved v other suitable technique(s) for a given reactionUnit 4
11.4understand experiments that can be used to investigate reaction rates by: i an initial-rate method, carrying out separate experiments where different initial concentrations of one reagent are used A ‘clock reaction’ is an acceptable approximation of this method. ii a continuous monitoring method to generate data to enable concentration-time or volume-time graphs to be plottedUnit 4
11.5be able to deduce the order (0, 1 or 2) with respect to a substance in a rate equation, using data from: i a concentration-time graph ii a rate-concentration graph iii an initial-rate methodUnit 4
11.6understand how to: i obtain data to calculate the order with respect to the reactants (and the hydrogen ion) in the acid-catalysed iodination of propanone ii use these data to make predictions about species involved in the rate-determining step iii deduce a possible mechanism for the reactionUnit 4
11.7be able to deduce the rate-determining step from a rate equation and vice versaUnit 4
11.8be able to deduce a reaction mechanism, using knowledge of the rate equation and the stoichiometric equation for a reactionUnit 4
11.9understand that knowledge of the rate equations for the hydrolysis of halogenoalkanes can be used to provide evidence for S N 1 and S N 2 mechanisms for tertiary and primary halogenoalkane hydrolysisUnit 4
11.10be able to use calculations and graphical methods to find the activation energy for a reaction from experimental data The Arrhenius equation will be given if needed.Unit 4
11.11understand the use of a solid (heterogeneous) catalyst for industrial reactions, in the gas phase, in terms of providing a surface for the reactionUnit 4
11.12CORE PRACTICALS 9a and 9b Following the rate of the iodine-propanone reaction by a titrimetric method and investigating a ‘clock reaction’ (Harcourt-Esson, iodine clock).Unit 4
11.13CORE PRACTICAL 10 Finding the activation energy of a reaction. Further suggested practicals: i the reaction between marble chips and hydrochloric acid (change of mass or change in volume of gas) ii the reaction between magnesium and hydrochloric acid to determine the activation energy iii following the rate of the iodine-propanone reaction by a colorimetric method iv the catalysis by a cobalt(II) salt of potassium sodium tartrate and hydrogen peroxide v the action of the enzyme urease on urea and thioureaUnit 4

12. Entropy and Energetics

1. Entropy and Energetics

12.1understand that, since endothermic reactions can occur spontaneously at room temperature, enthalpy changes alone do not control whether reactions occurUnit 4
12.2understand entropy as a measure of disorder of a system in terms of the random dispersal of molecules and of energy quanta between moleculesUnit 4
12.3understand that the entropy of a substance increases with temperature, that entropy increases as solid → liquid → gas and that perfect crystals at zero kelvin have zero entropyUnit 4
12.4be able to interpret the natural direction of change as being in the direction of increasing total entropy (positive entropy change), including gases spread spontaneously through a roomUnit 4
12.5understand why entropy changes occur during: i changes of state ii dissolving of a solid ionic lattice iii reactions in which there is a change in the number of moles from reactants to productsUnit 4
12.6understand that the total entropy change of any reaction is the sum of the entropy change of the system and the entropy change of the surroundings, summarised by the expression: ∆ S total = ∆ S system + ∆ S surroundingsUnit 4
12.7be able to calculate the entropy change of the system for a reaction, ∆ S system , given the entropies of the reactants and productsUnit 4
12.8be able to calculate the entropy change in the surroundings, and hence ∆ S total , using the expression ∆ S surroundings = − ∆ H TUnit 4
12.9understand that the feasibility of a reaction depends on: i the balance between ∆ S system and ∆ S surroundings , so that even endothermic reactions can occur spontaneously at room temperature ii temperature, as higher temperatures decrease the magnitude of ∆ S surroundings so its contribution to ∆ S total is less Students should be able to calculate the temperature at which a reaction is feasible . Students may also use ∆ G = ∆ H - T∆ S system in answers, although this approach is not a requirement of the specification.Unit 4
12.10understand that reactions can occur as long as ∆ S total is positive even if one of the other entropy changes is negativeUnit 4
12.11understand and distinguish between the concepts of thermodynamic stability and kinetic stabilityUnit 4
12.12be able to define the terms: i standard enthalpy change of atomisation, ∆ at H ii electron affinity iii lattice energy (as the exothermic process for the formation of one mole of an ionic solid from its gaseous ions)Unit 4
12.13be able to construct Born-Haber cycles and carry out related calculationsUnit 4
12.14understand that a comparison of the experimental lattice energy value (from a Born-Haber cycle) with the theoretical value (obtained from electrostatic theory) in a particular compound indicates the degree of covalent bondingUnit 4
12.15understand that polarisation of anions by cations leads to some covalency in an ionic bond, based on evidence from the Born-Haber cycleUnit 4
12.16be able to define the terms ‘enthalpy change of solution, ∆ sol H ’ and ‘ enthalpy change of hydration , ∆ hyd H of an ion’Unit 4
12.17be able to use energy cycles and energy level diagrams to calculate the enthalpy change of solution of an ionic compound, using enthalpy change of hydration and lattice energyUnit 4
12.18understand the effect of ionic charge and ionic radius on the values of enthalpy change of hydration and the lattice energy of an ionic compoundUnit 4
12.19be able to use entropy and enthalpy changes of solution values to predict the solubility of ionic compounds and discuss trends in the solubility of ionic compounds covered in Unit 2 Further suggested practical Calculate the enthalpy change when a variety of ionic solids are dissolved in waterUnit 4

13. Chemical Equilibria

1. Chemical Equilibria

13.1be able to deduce an expression for K c , for homogeneous and heterogeneous systems, in terms of equilibrium concentrationsUnit 4
13.2be able to deduce an expression for K p for homogeneous and heterogeneous systems, in terms of equilibrium partial pressures in atmUnit 4
13.3be able to calculate a value, with units where appropriate, for the equilibrium constants (Kc and Kp) for homogeneous and heterogeneous reactions, from experimental dataUnit 4
13.4understand how, if at all, a change in temperature, pressure or the presence of a catalyst affects the equilibrium composition in a homogeneous or heterogeneous systemUnit 4
13.5understand that the value of the equilibrium constant is not affected by changes in concentration or pressure or by the addition of a catalystUnit 4
13.6know the effect of changing the temperature on the equilibrium constant (Kc and Kp) for both exothermic and endothermic reactionsUnit 4
13.7understand that the effect of temperature on the position of equilibrium is explained using a change in the value of the equilibrium constantUnit 4
13.8understand the effect of a change in temperature on: i the value of ∆ S total ii the magnitude of the equilibrium constant, since ∆ S total = R ln KUnit 4
13.9be able to apply knowledge of the value of equilibrium constants to predict the extent to which a reaction takes place Further suggested practicals: i the reaction of ethanol and ethanoic acid (this can be used as an example of the use of ICT to present and analyse data) ii the equilibrium Fe 2+ (aq) + Ag + (aq) ⇌ Fe 3+ (aq) + Ag(s) iii the distribution of ammonia or iodine between two immiscible solvents iv the thermal decomposition of ammonium chloride v the effect of temperature and pressure changes in the system 2NO 2 ⇌ N 2 O 4Unit 4

14. Acid-base Equilibria

1. Acid-base Equilibria

14.1understand that a Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton acceptor and that acid-base reactions involve proton transferUnit 4
14.2be able to identify Brønsted–Lowry conjugate acid-base pairsUnit 4
14.3be able to define the term ‘pH’Unit 4
14.4be able to calculate pH from hydrogen ion concentrationUnit 4
14.5be able to calculate the concentration of hydrogen ions in a solution, in mol dm -3 , from its pH, using the expression [H + ] = 10 -pHUnit 4
14.6understand the difference between a strong acid and a weak acid in terms of the degree of dissociationUnit 4
14.7be able to calculate the pH of a strong acidUnit 4
14.8be able to deduce the expression for the acid dissociation constant, K a , for a weak acidUnit 4
14.9be able to calculate the pH of a weak acid from K a or p K a values, making relevant assumptions Students will not be expected to solve quadratic equations.Unit 4
14.10be able to define the ionic product of water, K wUnit 4
14.11be able to calculate the pH of a strong base from its concentration, using K w or p K wUnit 4
14.12be able to define the terms ‘p K a ’ and ‘p K w ’Unit 4
14.13be able to analyse data from the following experiments: i measuring the pH of a variety of substances, including equimolar solutions of strong and weak acids, strong and weak bases, and salts ii comparing the pH of a strong and weak acid after dilution 10, 100 and 1000 timesUnit 4
14.14be able to calculate K a for a weak acid from experimental data given the pH of a solution containing a known mass of acidUnit 4
14.15be able to draw and interpret titration curves, using all combinations of strong and weak monoprotic and diprotic acids with bases, and apply these principles to diprotic acids and basesUnit 4
14.16be able to select a suitable indicator for a titration, using a titration curve and appropriate dataUnit 4
14.17know what is meant by the term ‘buffer solution’Unit 4
14.18understand the action of a buffer solutionUnit 4
14.19be able to calculate the pH of a buffer solution given appropriate dataUnit 4
14.20be able to calculate the concentrations of solutions required to prepare a buffer solution of a given pHUnit 4
14.21understand how to use a weak acid-strong base or strong acid-weak base titration curve to: i demonstrate buffer action ii determine K a from the pH at the point where half the acid is neutralised/ equivalence pointUnit 4
14.22understand the importance of buffer solutions in biological environments: i buffers in cells and in blood (H 2 CO 3 /HC O - 3 ) ii in foods to prevent deterioration due to pH change (caused by bacterial or fungal activity)Unit 4
14.23CORE PRACTICAL 11 Finding the K a value for a weak acid. Further suggested practicals: i carry out the measuring of the pH of solutions mentioned in 14.13 ii obtain data to draw titration curves mentioned in 14.15 (which gives an opportunity to use data loggers) iii analysis of vinegarUnit 4

15. Organic Chemistry: Carbonyls, Carboxylic Acids

1. Organic Chemistry: Carbonyls, Carboxylic Acids

15.1know that optical isomerism is a result of chirality in molecules with a single chiral centreUnit 4
15.2understand that optical isomerism results from chiral centre(s) in a molecule with asymmetric carbon atom(s) and that optical isomers (enantiomers) are object and non-superimposable mirror images and be able to draw 3D diagrams of these optical isomersUnit 4
15.3know that optical activity is the ability of a single optical isomer to rotate the plane of polarisation of plane-polarised monochromatic light in molecules containing a single chiral centreUnit 4
15.4know what is meant by the term ‘racemic mixture’Unit 4
15.5be able to use data on optical activity of reactants and products as evidence for S N 1 and S N 2 mechanisms and addition to carbonyl compoundsUnit 4
15.6understand the nomenclature of aldehydes and ketones and be able to draw their structural, displayed and skeletal formulaeUnit 4
15.7understand that aldehydes and ketones: i do not form intermolecular hydrogen bonds and this affects their physical properties ii can form hydrogen bonds with water and this affects their solubilityUnit 4
15.8understand the reactions of carbonyl compounds with: i Fehling’s or Benedict’s solution, Tollens’ reagent and acidified dichromate(VI) ions In equations, the oxidising agent can be represented as [O] . ii lithium tetrahydridoaluminate(III) (lithium aluminium hydride) in dry ether (ethoxyethane) In equations, the reducing agent can be represented by [H] . iii HCN, in the presence of KCN, as a nucleophilic addition reaction, using curly arrows, relevant lone pairs, dipoles and evidence of optical activity to show the mechanism iv 2,4-dinitrophenylhydrazine (2,4-DNPH), as a qualitative test for the presence of a carbonyl group and to identify a carbonyl compound given data of the melting temperatures of derivatives The equation for this reaction is not required. v iodine in the presence of alkali (the iodoform test) Further suggested practical: Reactions of aldehydes and ketones given in 15.8 i, iv and vUnit 4
15.9understand the nomenclature of carboxylic acids and be able to draw their structural, displayed and skeletal formulaeUnit 4
15.10understand that hydrogen bonding affects the physical properties of carboxylic acids, in relation to their boiling temperatures and solubilityUnit 4
15.11understand that carboxylic acids can be prepared by the oxidation of alcohols or aldehydes and the hydrolysis of nitrilesUnit 4
15.12understand the reactions of carboxylic acids with: i lithium tetrahydridoaluminate (III) (lithium aluminium hydride) in dry ether (ethoxyethane) ii bases to produce salts iii phosphorus(V) chloride (phosphorus pentachloride) iv alcohols in the presence of an acid catalyst Further suggested practicals : i solubility of a range of carboxylic acids, aldehydes and ketones ii preparation of carboxylic acids by the oxidation of alcohols and aldehydes iii reactions of carboxylic acids given in 15.12 ii, iii and ivUnit 4
15.13understand the nomenclature of acyl chlorides and esters and be able to draw their structural, displayed and skeletal formulaeUnit 4
15.14understand the reactions of acyl chlorides with: i water ii alcohols iii concentrated ammonia iv aminesUnit 4
15.15understand the hydrolysis reactions of esters, in acidic and alkaline solutionUnit 4
15.16understand how polyesters, such as terylene, are formed by condensation polymerisation reactions. Further suggested practicals: i demonstration of the reactions of ethanoyl chloride given in 15.14 i, ii and iii ii the preparation of esters such as ethyl ethanoate as a solvent or a pineapple flavouring iii hydrolysis of an esterUnit 4
15.17be able to use data from mass spectra to: i suggest possible structures of a simple organic compound given accurate relative molecular masses ii calculate the accurate relative molecular mass of a compound, given accurate relative atomic masses to four decimal placesUnit 4
15.18understand that carbon-13, (13C) NMR spectroscopy provides information about the positions of 13C atoms in a moleculeUnit 4
15.19be able to use data from 13 C NMR spectroscopy to: i predict the different environments for carbon atoms present in a molecule, given values of chemical shift, δ ii justify the number of peaks present in a 13 C NMR spectrum in terms of the number of carbon atoms in different environmentsUnit 4
15.20be able to use both low and high resolution proton NMR spectroscopy to: i predict the different types of proton present in a molecule, given values of chemical shift, δ ii relate relative peak areas, or ratio number of protons, to the relative numbers of 1 H atoms in different environments iii deduce the splitting patterns of adjacent, non-equivalent protons using the (n+1) rule and hence suggest the possible structures for a molecule iv predict the chemical shifts and splitting patterns of the 1 H atoms in a given moleculeUnit 4
15.21know that chromatography separates components of a mixture using a mobile phase and a stationary phaseUnit 4
15.22be able to calculate R f values from one-way chromatograms in paper and thin-layer chromatography (TLC) and understand reasons for differences in R f valuesUnit 4
15.23know that high-performance liquid chromatography, HPLC, and gas chromatography, GC, are types of column chromatography that separate substances because of different retention times in the column and may be used in conjunction with mass spectrometry, in applications such as forensics or drug testing in sportUnit 4

16. Redox Equilibria

1. Redox Equilibria

16.1understand the terms ‘oxidation’ and ‘reduction’ in terms of electron transfer and changes in oxidation number, applied to s- , p- and d -block elementsUnit 5
16.2know what is meant by the term ‘standard electrode potential’, E oUnit 5
16.3know that the standard electrode potential, E o , is measured in conditions of: i 298 K temperature ii 100 kPa pressure of gases iii 1.00 mol dm -3 concentration of ionsUnit 5
16.4know the features of the standard hydrogen electrode and understand why a reference electrode is necessaryUnit 5
16.5understand that different methods are used to measure standard electrode potentials of: i metals or non-metals in contact with their ions in aqueous solution ii ions of the same element with different oxidation numbersUnit 5
16.6CORE PRACTICAL 12 Investigating some electrochemical cells.Unit 5
16.7be able to calculate a standard emf, E o cell , by combining two standard electrode potentialsUnit 5
16.8be able to write cell diagrams using the conventional representation of half-cellsUnit 5
16.9understand the importance of the conditions when measuring an electrode potential, EUnit 5
16.10be able to use standard electrode potentials to predict the thermodynamic feasibility of a reactionUnit 5
16.11understand that E o cell is directly proportional to the total entropy change and to ln K for a reactionUnit 5
16.12understand the limitations of predictions made using standard electrode potentials, in terms of kinetic stability of systems and departure from standard conditionsUnit 5
16.13know that standard electrode potentials are sometimes referred to as standard reduction potentials and can be listed as an electrochemical seriesUnit 5
16.14understand how standard electrode potentials can be used to predict the thermodynamic feasibility of disproportionation reactionsUnit 5
16.15be able to carry out both structured and unstructured titration calculations involving redox reactions, including iron(II) ions and potassium manganate(VII) and sodium thiosulfate and iodineUnit 5
16.16be able to discuss the uncertainty of measurements and their implications for the validity of the final resultsUnit 5
16.17CORE PRACTICALS 13a and 13b Carry out redox titrations with both: i iron(II) ions and potassium manganate(VII) ii sodium thiosulfate and iodineUnit 5
16.18understand that fuel cells use the energy released on the reaction of a fuel with oxygen to generate a voltage Knowledge that methanol and other hydrogen-rich fuels are used in fuel cells is expected.Unit 5
16.19know the electrode reactions that occur in a hydrogen-oxygen fuel cell Knowledge of hydrogen-oxygen fuel cells with both acidic and alkaline electrolyte is expected. Further suggested practicals: i investigate the percentage of copper in brass, using iodine-thiosulfate titration ii investigate the percentage of iron in iron tablets, using potassium manganate(VII) titration iii prepare crystals of potassium iodate(VII) and measure their purityUnit 5

17. Transition Metals and their Chemistry

1. Transition Metals and their Chemistry

17.1know that transition metals are d- block elements that form one or more stable ions with incompletely-filled d- orbitalsUnit 5
17.2be able to deduce the electronic configurations of atoms and ions of the d -block elements of Period 4 (Sc-Zn) given their atomic number and charge (if any)Unit 5
17.3understand why transition metals show variable oxidation numberUnit 5
17.4know what is meant by the term ‘ligand’Unit 5
17.5understand that dative (coordinate) covalent bonding is involved in the formation of complex ionsUnit 5
17.6know that a complex ion is a central metal ion surrounded by ligandsUnit 5
17.7know that aqueous solutions of transition metal ions are usually colouredUnit 5
17.8understand that the colour of aqueous ions, and other complex ions, is a consequence of the splitting of the energy levels of the d -orbitals by ligandsUnit 5
17.9understand why there is a lack of colour in some aqueous ions and other complex ionsUnit 5
17.10understand the meaning of the term ‘coordination number’Unit 5
17.11understand that colour changes in transition metal ions may arise as a result of changes in: i oxidation number of the ion ii ligand iii coordination number of the complexUnit 5
17.12understand that H 2 O, OH - and NH 3 act as monodentate ligandsUnit 5
17.13understand why complexes with six-fold coordination have an octahedral shape, such as those formed by metal ions with H 2 O, OH - and NH 3 as ligandsUnit 5
17.14know that transition metal ions may form tetrahedral complexes with relatively large ions such as Cl -Unit 5
17.15know that square planar complexes are also formed by transition metal ions and that cis- platin is an example of such a complex which is used in cancer treatment where it is supplied as a single isomer and not in a mixture with the trans formUnit 5
17.16understand the terms ‘bidentate’ and ‘hexadentate’ in relation to ligands, and be able to identify examples such as NH 2 CH 2 CH 2 NH 2 and EDTA 4-Unit 5
17.17know that haemoglobin is an iron(II) complex containing a polydentate ligand and that ligand exchange occurs when an oxygen molecule bound to haemoglobin is replaced by a carbon monoxide molecule The structure of the haem group will not be assessed.Unit 5
17.18know the colours of the oxidation states of vanadium (+5, +4, +3 and +2) in its compoundsUnit 5
17.19understand redox reactions for the interconversion of the oxidation states of vanadium (+5, +4, +3 and +2), in terms of the relevant E o valuesUnit 5
17.20understand, in terms of the relevant E values, that the dichromate(VI) ion, Cr 2 O 2- 7 i can be reduced to Cr 3+ and Cr 2+ ions using zinc in acidic conditions ii can be produced by the oxidation of Cr 3+ ions using hydrogen peroxide in alkaline conditions (followed by acidification)Unit 5
17.21know that the dichromate(VI) ion, Cr 2 O 7 2- can be converted into chromate(VI) ions as a result of the equilibrium Cr 2 O 2- 7 + H 2 O ⇌ 2CrO 2- 4 +2H +Unit 5
17.22be able to record observations and write suitable equations for the reactions of Cr 3+ (aq), Mn 2+ (aq), Fe 2+ (aq), Fe 3+ (aq), Co 2+ (aq), Ni 2+ (aq), Cu 2+ (aq) and Zn 2+ (aq) with aqueous sodium hydroxide and aqueous ammonia, including in excessUnit 5
17.23be able to write ionic equations to show the meaning of amphoteric behaviour, deprotonation and ligand exchange in the reactions in 17.22Unit 5
17.24understand that ligand exchange, and an accompanying colour change, occurs in the formation of: i [Cu(NH3)4(H2O)2]2+ from [Cu(H2O)6]2+ via Cu(OH)2(H2O)4 ii [CuCl4]2- from [Cu(H2O)6]2+ iii [CoCl4]2- from [Co(H2O)6]2+Unit 5
17.25understand, in terms of the positive increase in ∆ S system , that the substitution of a monodentate ligand by a bidentate or hexadentate ligand leads to a more stable complex ionUnit 5
17.26know that transition metals and their compounds can act as heterogeneous and homogeneous catalystsUnit 5
17.27know that a heterogeneous catalyst is in a different phase from the reactants and that the reaction occurs at the surface of the catalystUnit 5
17.28understand, in terms of oxidation number, how V 2 O 5 acts as a catalyst in the contact processUnit 5
17.29understand how a catalytic converter decreases carbon monoxide and nitrogen monoxide emissions from internal combustion engines by: i adsorption of CO and NO molecules onto the surface of the catalyst, resulting in the weakening of bonds and chemical reaction ii desorption of CO 2 and N 2 product molecules from the surface of the catalystUnit 5
17.30know that a homogeneous catalyst is in the same phase as the reactants and appreciate that the catalysed reaction will proceed via an intermediate speciesUnit 5
17.31understand the role of Fe 2+ ions in catalysing the reaction between I - and S 2 O 8 2- ionsUnit 5
17.32know the role of Mn 2+ ions in autocatalysing the reaction between MnO - 4 and C 2 O 4 2- ionsUnit 5
17.33CORE PRACTICAL 14 The preparation of a transition metal complex.Unit 5

18. Organic Chemistry – Arenes

1. Organic Chemistry – Arenes

18.1be able to use thermochemical, X-ray diffraction and infrared data as evidence for the structure and stability of the benzene ring Students may represent the structure of benzene as or as appropriate in equations and mechanisms.Unit 5
18.2understand that the delocalised model for the structure of benzene involves overlap of p -orbitals to form π -bondsUnit 5
18.3understand why benzene is resistant to bromination, compared to alkenes, in terms of delocalisation of π -bonds in benzene compared to the localised electron density of the π -bond in alkenesUnit 5
18.4know the following reactions of benzene, limited to: i oxygen in air (combustion to form a smoky flame) ii bromine, in the presence of a catalyst iii a mixture of concentrated nitric and sulfuric acids iv fuming sulfuric acid v halogenoalkanes and acyl chlorides with aluminium chloride as catalyst (Friedel-Crafts reaction)Unit 5
18.5understand the mechanism of the electrophilic substitution reactions of benzene in halogenation, nitration and Friedel-Crafts reactions, including the generation of the electrophileUnit 5
18.6understand the reaction of phenol with bromine water and the reasons for the relative ease of this reaction compared to benzene Further suggested practical: Carry out the reactions in 18.4, and 18.6 where appropriate (using methylbenzene or methoxybenzene)Unit 5

19. Organic Nitrogen Compounds: Amines, Amides,

1. Organic Nitrogen Compounds: Amines, Amides,

19.1understand the nomenclature of amides, amines and amino acids and be able to draw their structural, displayed and skeletal formulaeUnit 5
19.2understand the reactions of primary aliphatic amines (using butylamine as an example) and aromatic amines (using phenylamine as an example) with: i water to form an alkaline solution ii acids to form salts iii halogenoalkanes iv ethanoyl chloride v copper(II) ions to form a complex ionUnit 5
19.3understand that amines are miscible with water as a result of hydrogen bonding, and the reasons for the difference in basicity between ammonia, primary aliphatic amines and primary aromatic aminesUnit 5
19.4understand, in terms of reagents and general reaction conditions, the preparation of primary aliphatic amines: i from halogenoalkanes ii by the reduction of nitrilesUnit 5
19.5know the preparation of aromatic amines by the reduction of aromatic nitro- compounds using tin and concentrated hydrochloric acidUnit 5
19.6be able to describe the reaction of aromatic amines with nitrous acid to form benzenediazonium ions, followed by a coupling reaction with phenol to form a dyeUnit 5
19.7understand that amides can be prepared from acyl chloridesUnit 5
19.8be able to describe: i condensation polymerisation for the formation of polyamides such as nylon and proteins ii addition polymerisation, including poly(propenamide) and poly(ethenol)Unit 5
19.9be able to draw the structural formulae of the repeat units of the polymers in 19.8Unit 5
19.10be able to comment on the physical properties of polyamides and the solubility in water of the addition polymer poly(ethenol) in terms of hydrogen bonding, including soluble laundry bags or liquid-detergent capsules (liquitabs)Unit 5
19.11be able to describe experiments to investigate the characteristic behaviour of amino acids limited to: i acidity and basicity and the formation of zwitterions ii effect of aqueous solutions on plane-polarised monochromatic light iii formation of peptide bonds by condensation polymerisationUnit 5
19.12CORE PRACTICAL 15 Analysis of some inorganic and organic unknowns.Unit 5

20. Organic Synthesis

1. Organic Synthesis

20.1be able to deduce the empirical formulae, molecular formulae and structural formulae from data drawn from combustion analysis, element percentage composition, characteristic reactions of functional groups, infrared spectra, mass spectra and NMR spectra (both 13 C and proton)Unit 5
20.2understand methods of increasing the length of the carbon chain in a molecule by the use of magnesium to form Grignard reagents and the reactions of the latter with carbon dioxide and with carbonyl compounds in dry etherUnit 5
20.3be able to use knowledge of organic chemistry contained given in this specification to solve problems such as: i predicting the properties of unfamiliar compounds containing one or more of the functional groups included in the specification and explain these predictions ii planning reaction schemes of up to four steps, recalling familiar reactions and using unfamiliar reactions given sufficient information iii selecting suitable practical procedures for carrying out reactions involving compounds with functional groups included in this specification iv identifying appropriate control measures to reduce risk based on data of hazardsUnit 5
20.4CORE PRACTICAL 16 The preparation of aspirin.Unit 5
20.5understand the following techniques used in the preparation and purification of organic compounds: i refluxing ii purification by washing, including with water and sodium carbonate solution iii solvent extraction iv recrystallisation v drying vi distillation vii steam distillation viii melting temperature determination ix boiling temperature determination Further suggested practicals: i carry out the preparation of an organic compound, including cholesteryl benzoate (a liquid crystal) or methyl 3-nitrobenzoate ii preparation of oil of wintergreenUnit 5