Lessons

Subatomic Particles and Isotopes

know the structure of an atom in terms of electrons, protons and neutrons · know the relative mass and charge of protons, neutrons and electrons · know what is meant by the terms 'atomic (proton) number' and 'mass number' · be able to use the atomic number and the mass number to determine the number of each type of subatomic particle in an atom or ion · understand the term 'isotope'

Electronic Configuration and Ionisation Energies

be able to define first, second and third ionisation energies and understand that all ionisation energies are endothermic · know that an orbital is a region within an atom that can hold up to two electrons with opposite spins · understand 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 removed · know 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-shells · be able to describe the shapes of s and p orbitals · know that orbitals in sub-shells: (i) each take a single electron before pairing up; (ii) pair up with two electrons of opposite spin · be 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 notation · understand that electronic configuration determines the chemical properties of an element · know 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 shells

Formulae, Equations and the Mole

know the terms 'atom', 'element', 'ion', 'molecule', 'compound', 'empirical formula' and 'molecular formula' · know 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 10^23 mol^-1) · write balanced full and ionic equations, including state symbols, for chemical reactions · understand the terms: (i) 'relative atomic mass' based on the 12C 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 atmosphere · calculate the concentration of a solution in mol dm^-3 and g dm^-3 (titration calculations are not required at this stage)

Stoichiometric Calculations: Masses, Gas Volumes and Limiting Reagents

be able to use experimental data to calculate empirical and molecular formulae · be able to use chemical equations to calculate reacting masses and vice versa, using the concepts of amount of substance and molar mass · be 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 = nRT for gases and volatile liquids · be 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 / sum of the molar masses of all products x 100%) · be able to determine a formula or confirm an equation by experiment, including evaluation of the data · CORE PRACTICAL 1: Measurement of the molar volume of a gas

Ionic Equations and Reaction Observations

be 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 reactions

Mass Spectrometry and Relative Masses

understand 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+ charge · be able to predict mass spectra, including relative peak heights, for diatomic molecules, including chlorine, given the isotopic abundances

Volumetric Analysis and Titration Calculations

be able to calculate solution concentrations, in mol dm^-3 and g dm^-3, including simple acid-base titrations using the indicators methyl orange and phenolphthalein · CORE PRACTICAL 3: Finding the concentration of a solution of hydrochloric acid · understand how to minimise the sources of measurement uncertainty in volumetric analysis and estimate the overall uncertainty in the calculated result · CORE PRACTICAL 4: Preparation of a standard solution from a solid acid and use it to find the concentration of a solution of sodium hydroxide

Ionic Bonding and Ion Polarisation

know 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 ions · be able to describe the formation of ions in terms of loss or gain of electrons · be able to draw dot-and-cross diagrams to show electrons in cations and anions · be able to describe ionic crystals as giant lattices of ions · know that ionic bonding is the result of strong net electrostatic attraction between ions · understand the effects of ionic radius and ionic charge on the strength of ionic bonding · understand reasons for the trends in ionic radii down a group in the Periodic Table, and for a set of isoelectronic ions, including N3- to Al3+ · understand the meaning of the term 'polarisation' as applied to ions · understand 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

Covalent and Dative Bonding, Electronegativity and Polarity

understand 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 molecules · be 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 Al2Cl6 and the ammonium ion · be able to describe the different structures formed by giant lattices of carbon atoms, including graphite, diamond and graphene, and discuss the applications of each · understand the meaning of the term 'electronegativity' as applied to atoms in a covalent bond · know 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 enough · be able to distinguish between polar bonds and polar molecules and predict whether or not a given molecule is likely to be polar

Hybridisation and Sigma/Pi Bonding

use sp, sp2 and sp3 hybridisation to describe bonding and geometry · describe sigma bonds (end-on overlap) and pi bonds (side-on overlap) · explain restricted rotation about C=C in terms of pi overlap · relate hybridisation to shape and bond angles (180, 120, 109.5 degrees) · apply hybridisation across carbon allotropes (diamond sp3, graphite/graphene sp2)

VSEPR: Shapes of Molecules and Ions

understand the principles of the electron-pair repulsion theory, used to interpret and predict the shapes of simple molecules and ions · understand the terms 'bond length' and 'bond angle' · know and be able to explain the shapes of, and bond angles in, BeCl2, BCl3, CH4, NH3, NH4+, H2O, CO2, gaseous PCl5, SF6 and C2H4 · be able to apply the electron-pair repulsion theory to predict the shapes of, and bond angles in, molecules and ions analogous to those in 3.18

Intermolecular Forces: van der Waals and Hydrogen Bonding

understand the nature of the following intermolecular forces: (i) London forces (instantaneous dipole-induced dipole); (ii) permanent dipole-permanent dipole interactions; (iii) hydrogen bonds · understand the interactions in molecules, such as H2O, liquid NH3 and liquid HF, which give rise to hydrogen bonding · understand 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 water · be able to predict the presence of hydrogen bonding in molecules analogous to those mentioned in 7.2

Boiling-Point Trends and Solvent Choice

understand, 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 HI · understand 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

Gases and the Ideal Gas Equation

explain the origin of gas pressure from molecular collisions with the walls · state the assumptions of the ideal gas (kinetic) model · recall and use the ideal gas equation pV = nRT · explain the limitations of the ideal gas model and conditions for real-gas deviation

Metallic Bonding and Lattice Types

understand that metals consist of giant lattices of metal ions in a sea of delocalised electrons · know that metallic bonding is the strong electrostatic attraction between metal ions and the delocalised electrons · be able to use the models in 3.20 and 3.21 to interpret simple properties of metals, including electrical conductivity and high melting temperature

Enthalpy Changes and Standard Definitions

know that the enthalpy change, delta H, is the heat energy change measured at constant pressure and that standard conditions are 100 kPa and a specified temperature, usually 298 K · know that, by convention, exothermic reactions have a negative enthalpy change and endothermic reactions have a positive enthalpy change · be able to construct and interpret enthalpy level diagrams, showing exothermic and endothermic enthalpy changes · know the definition of standard enthalpy change of: (i) reaction, delta_rH; (ii) formation, delta_fH; (iii) combustion, delta_cH; (iv) neutralisation, delta_neutH; (v) atomisation, delta_atH

Calorimetry and Hess's Law

be able to use experimental data to calculate: (i) energy transferred in a reaction recalling and using the expression: energy transferred (J) = mass (g) x specific heat capacity (J g^-1 °C^-1) x 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. · know 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 experiments · CORE PRACTICAL 2: Determination of the enthalpy change of a reaction using Hess's Law · be 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.

Bond Energies and Enthalpy Calculations

understand the terms 'bond enthalpy' and 'mean bond enthalpy', and be able to use bond enthalpies to calculate enthalpy changes, understanding the limitations of this method · be able to calculate mean bond enthalpies from enthalpy changes of reaction · understand 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 temperature

Lattice Energy, Born-Haber Cycles and Enthalpy of Solution

be able to define the terms: (i) standard enthalpy change of atomisation, delta_atH; (ii) electron affinity; (iii) lattice energy (as the exothermic process for the formation of one mole of an ionic solid from its gaseous ions) · be able to construct Born-Haber cycles and carry out related calculations · understand 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 bonding · understand that polarisation of anions by cations leads to some covalency in an ionic bond, based on evidence from the Born-Haber cycle · be able to define the terms 'enthalpy change of solution, delta_solH' and 'enthalpy change of hydration, delta_hydH of an ion' · be 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 energy · understand the effect of ionic charge and ionic radius on the values of enthalpy change of hydration and the lattice energy of an ionic compound · be 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

Entropy and Gibbs Free Energy

understand that, since endothermic reactions can occur spontaneously at room temperature, enthalpy changes alone do not control whether reactions occur · understand entropy as a measure of disorder of a system in terms of the random dispersal of molecules and of energy quanta between molecules · understand 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 entropy · be 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 room · understand 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 products · understand 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: delta S_total = delta S_system + delta S_surroundings · be able to calculate the entropy change of the system for a reaction, delta S_system, given the entropies of the reactants and products · be able to calculate the entropy change in the surroundings, and hence delta S_total, using the expression delta S_surroundings = -delta H / T · understand that the feasibility of a reaction depends on: (i) the balance between delta S_system and delta S_surroundings, so that even endothermic reactions can occur spontaneously at room temperature; (ii) temperature, as higher temperatures decrease the magnitude of delta S_surroundings so its contribution to delta S_total is less. Students should be able to calculate the temperature at which a reaction is feasible. Students may also use delta G = delta H - T delta S_system in answers, although this approach is not a requirement of the specification. · understand that reactions can occur as long as delta S_total is positive even if one of the other entropy changes is negative · understand and distinguish between the concepts of thermodynamic stability and kinetic stability

Oxidation Numbers and Redox

know what is meant by the term 'oxidation number' and understand the rules for assigning oxidation numbers · be able to calculate the oxidation number of elements in compounds and ions, including in peroxides and metal hydrides · be able to indicate the oxidation number of an element in a compound or an ion, using a Roman numeral · be able to write formulae given oxidation numbers · understand 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 elements · know that oxidising agents gain electrons and reducing agents lose electrons · understand that a disproportionation reaction involves an element in a single species being simultaneously oxidised and reduced · know that oxidation number is a useful concept in terms of the classification of reactions as redox and as disproportionation · understand 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 number · be able to write ionic half-equations and use them to construct full ionic equations

Standard Electrode Potentials and the Hydrogen Electrode

understand the terms 'oxidation' and 'reduction' in terms of electron transfer and changes in oxidation number, applied to s-, p- and d-block elements · know what is meant by the term 'standard electrode potential', E° · know that the standard electrode potential, E°, is measured in conditions of: (i) 298 K temperature; (ii) 100 kPa pressure of gases; (iii) 1.00 mol dm^-3 concentration of ions · know the features of the standard hydrogen electrode and understand why a reference electrode is necessary · understand 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 numbers · CORE PRACTICAL 12: Investigating some electrochemical cells

Cell Potentials, Feasibility and the Nernst Equation

be able to calculate a standard emf, E°_cell, by combining two standard electrode potentials · be able to write cell diagrams using the conventional representation of half-cells · understand the importance of the conditions when measuring an electrode potential, E · be able to use standard electrode potentials to predict the thermodynamic feasibility of a reaction · understand that E°_cell is directly proportional to the total entropy change and to ln K for a reaction · understand the limitations of predictions made using standard electrode potentials, in terms of kinetic stability of systems and departure from standard conditions · know that standard electrode potentials are sometimes referred to as standard reduction potentials and can be listed as an electrochemical series · understand how standard electrode potentials can be used to predict the thermodynamic feasibility of disproportionation reactions

Redox Titrations and Electrochemical Cells

be 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 iodine · be able to discuss the uncertainty of measurements and their implications for the validity of the final results · CORE PRACTICALS 13a and 13b: Carry out redox titrations with both: (i) iron(II) ions and potassium manganate(VII); (ii) sodium thiosulfate and iodine · understand 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) · know 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)

Electrolysis and Faraday Calculations

predict the products of electrolysis of molten and aqueous electrolytes · recall and use F = Le (the Faraday constant) and Q = It · calculate masses and volumes of products from charge passed (n = Q/zF) · explain selective discharge of ions at the electrodes

Dynamic Equilibrium and Le Chatelier's Principle

know 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 constant · be able to predict and justify the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium in a homogeneous system · evaluate data to explain the necessity, for many industrial processes, to reach a compromise between the yield and the rate of reaction

The Equilibrium Constant Kc

Writing the Kc expression for any homogeneous equilibrium · Setting up ICE tables (Initial / Change / Equilibrium) to compute Kc from experimental data · Deriving the units of Kc from Δn and the invariance of Kc under concentration / pressure / catalyst changes

Kp for Gas-Phase Equilibria

Mole fractions and partial pressures: p_i = x_i × P_total · Building the Kp expression, calculating Kp from equilibrium moles + total pressure, and deriving the units (atm)^Δn

Temperature Effect on K and Extent of Reaction

K depends only on T: concentration, pressure, volume, and catalyst all leave K unchanged · Temperature effect on K for exothermic vs endothermic reactions, linked to ΔS_total = R ln K · Using the magnitude of K (K >> 1, K ≈ 1, K << 1) to predict the extent of reaction

The Haber and Contact Processes

describe and explain the conditions used in the Haber process · describe and explain the conditions used in the Contact process · apply Le Chatelier's principle to the yield/rate compromise in industrial equilibria

Bronsted-Lowry Theory, Strong/Weak Acids and pH Basics

understand 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 transfer · be able to identify Brønsted-Lowry conjugate acid-base pairs · be able to define the term 'pH' · be able to calculate pH from hydrogen ion concentration · be able to calculate the concentration of hydrogen ions in a solution, in mol dm^-3, from its pH, using the expression [H+] = 10^-pH

pH Titration Curves and Indicator Selection

be 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 bases · be able to select a suitable indicator for a titration, using a titration curve and appropriate data

Conjugate Pairs, Ka, pKa, Kw and pH Calculations

understand the difference between a strong acid and a weak acid in terms of the degree of dissociation · be able to calculate the pH of a strong acid · be able to deduce the expression for the acid dissociation constant, Ka, for a weak acid · be able to calculate the pH of a weak acid from Ka or pKa values, making relevant assumptions (students will not be expected to solve quadratic equations) · be able to define the ionic product of water, Kw · be able to calculate the pH of a strong base from its concentration, using Kw or pKw · be able to define the terms 'pKa' and 'pKw' · be 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 times · be able to calculate Ka for a weak acid from experimental data given the pH of a solution containing a known mass of acid

Buffer Solutions, Ksp and Partition Coefficients

know what is meant by the term 'buffer solution' · understand the action of a buffer solution · be able to calculate the pH of a buffer solution given appropriate data · be able to calculate the concentrations of solutions required to prepare a buffer solution of a given pH · understand how to use a weak acid-strong base or strong acid-weak base titration curve to: (i) demonstrate buffer action; (ii) determine Ka from the pH at the point where half the acid is neutralised/equivalence point · understand the importance of buffer solutions in biological environments: (i) buffers in cells and in blood (H2CO3/HCO3-); (ii) in foods to prevent deterioration due to pH change (caused by bacterial or fungal activity) · CORE PRACTICAL 11: Finding the Ka value for a weak acid

Collision Theory, Boltzmann Distribution and Catalysis

understand, in terms of the collision theory, the effect of changes in concentration, temperature, pressure and surface area on the rate of a chemical reaction · understand that reactions take place only when collisions have sufficient energy, known as the activation energy · be able to calculate the rate of a reaction from: (i) the time taken for a reaction, using rate = 1/time; (ii) the gradient of a suitable graph, by drawing a tangent, either for initial rate, or at a time, t · understand qualitatively, in terms of the Maxwell-Boltzmann distribution of molecular energies, how changes in temperature affect the rate of a reaction · understand the role of catalysts in providing alternative reaction routes of lower activation energy · be able to draw the reaction profiles for uncatalysed and catalysed reactions, including the energy level of the intermediate formed with the catalyst · understand the use of catalysts in industry to make processes more sustainable by using less energy and/or higher atom economy · be able to interpret the action of a catalyst in terms of a qualitative understanding of the Maxwell-Boltzmann distribution of molecular energies

Rate Equations, Order and Half-Life

understand 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 catalyst · be 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 order

Determining Order from Graphs and Data

be 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 reaction · understand 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 plotted · be 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 method

Reaction Mechanisms, Rate-Determining Step and Catalysis

understand 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 reaction · be able to deduce the rate-determining step from a rate equation and vice versa · be able to deduce a reaction mechanism, using knowledge of the rate equation and the stoichiometric equation for a reaction · understand that knowledge of the rate equations for the hydrolysis of halogenoalkanes can be used to provide evidence for SN1 and SN2 mechanisms for tertiary and primary halogenoalkane hydrolysis · be 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) · understand the use of a solid (heterogeneous) catalyst for industrial reactions, in the gas phase, in terms of providing a surface for the reaction · CORE 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) · CORE PRACTICAL 10: Finding the activation energy of a reaction

Period 3: Periodicity of Physical Properties

be 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' · be 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 group

Period 3: Reactions, Oxides, Hydroxides and Chlorides

describe the reactions of Period 3 elements with oxygen, chlorine and water · describe the acid-base behaviour of Period 3 oxides and hydroxides, including amphoterism · describe the reactions of Period 3 chlorides with water and the resulting pH

Group 2: Reactions, Solubility and Thermal Stability

understand reasons for the trend in ionisation energy down Groups 1 and 2 · understand reasons for the trend in reactivity of the elements down Group 1 (Li to K) and Group 2 (Mg to Ba) · know the reactions of the elements of Group 1 (Li to K) and Group 2 (Mg to Ba) with oxygen, chlorine and water · know 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 acid · know the trends in solubility of the hydroxides and sulfates of Group 2 elements · understand 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 involved · understand 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) · know 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 elements · know reactions, including ionic equations where appropriate, for identifying: (i) carbonate ions, CO3^2-, and hydrogencarbonate ions, HCO3^-, using an aqueous acid to form carbon dioxide (and testing the gas with limewater); (ii) sulfate ions, SO4^2-, using acidified barium chloride solution; (iii) ammonium ions, NH4+, using sodium hydroxide solution and warming to form ammonia (and testing with litmus and HCl fumes)

Group 17: Halogens and Halides

understand 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 group · understand the trend in reactivity of Group 7 elements in terms of the redox reactions of Cl2, Br2 and I2 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) · understand, 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 above · understand 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) · be able to make predictions about fluorine and astatine and their compounds, in terms of knowledge of trends in halogen chemistry

Nitrogen and Sulfur

explain the lack of reactivity of nitrogen via its triple bond and non-polarity · describe the basicity of ammonia and formation of the ammonium ion · describe the oxides of nitrogen and their catalytic removal from exhaust gases · explain the role of NOx in photochemical smog and acid rain · describe the formation and uses of sulfuric acid

Transition Elements: Characteristics, Ligands and Complexes

know that transition metals are d-block elements that form one or more stable ions with incompletely-filled d-orbitals · be 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) · understand why transition metals show variable oxidation number · know what is meant by the term 'ligand' · understand that dative (coordinate) covalent bonding is involved in the formation of complex ions · know that a complex ion is a central metal ion surrounded by ligands

Colour, Coordination Number and Complex Geometry

know that aqueous solutions of transition metal ions are usually coloured · understand 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 ligands · understand why there is a lack of colour in some aqueous ions and other complex ions · understand the meaning of the term 'coordination number' · understand 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 complex · understand that H2O, OH- and NH3 act as monodentate ligands · understand why complexes with six-fold coordination have an octahedral shape, such as those formed by metal ions with H2O, OH- and NH3 as ligands · know that transition metal ions may form tetrahedral complexes with relatively large ions such as Cl- · know 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 form · understand the terms 'bidentate' and 'hexadentate' in relation to ligands, and be able to identify examples such as NH2CH2CH2NH2 and EDTA^4- · know 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)

Variable Oxidation States: Vanadium and Chromium

know the colours of the oxidation states of vanadium (+5, +4, +3 and +2) in its compounds · understand redox reactions for the interconversion of the oxidation states of vanadium (+5, +4, +3 and +2), in terms of the relevant E° values · understand, in terms of the relevant E° values, that the dichromate(VI) ion, Cr2O7^2-, (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) · know that the dichromate(VI) ion, Cr2O7^2-, can be converted into chromate(VI) ions as a result of the equilibrium Cr2O7^2- + H2O <-> 2CrO4^2- + 2H+

Ligand Exchange, Stability Constants and the Chelate Effect

be 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 excess · be able to write ionic equations to show the meaning of amphoteric behaviour, deprotonation and ligand exchange in the reactions in 17.22 · understand 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+ · understand, in terms of the positive increase in delta S_system, that the substitution of a monodentate ligand by a bidentate or hexadentate ligand leads to a more stable complex ion

Transition Metals as Catalysts

know that transition metals and their compounds can act as heterogeneous and homogeneous catalysts · know that a heterogeneous catalyst is in a different phase from the reactants and that the reaction occurs at the surface of the catalyst · understand, in terms of oxidation number, how V2O5 acts as a catalyst in the contact process · understand 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 CO2 and N2 product molecules from the surface of the catalyst · know that a homogeneous catalyst is in the same phase as the reactants and appreciate that the catalysed reaction will proceed via an intermediate species · understand the role of Fe^2+ ions in catalysing the reaction between I- and S2O8^2- ions · know the role of Mn^2+ ions in autocatalysing the reaction between MnO4^- and C2O4^2- ions · CORE PRACTICAL 14: The preparation of a transition metal complex

Introduction to Organic Chemistry: Formulae, Nomenclature and Reaction Types

understand the difference between hazard and risk · understand the hazards associated with organic compounds and why it is necessary to carry out risk assessments when dealing with potentially hazardous materials · be 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 substances · understand the concepts of homologous series and functional group · be 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 C10) · be able to classify reactions as addition, substitution, oxidation, reduction or polymerisation · understand that bond breaking can be: (i) homolytic, to produce free radicals; (ii) heterolytic, to produce ions · know definitions of the terms 'free radical' and 'electrophile'

Reaction Mechanisms, Nucleophiles and Electrophiles

be able to classify reactions (including those in Unit 1) as addition, elimination, substitution, oxidation, reduction, hydrolysis or polymerisation · understand the concept of a reaction mechanism · understand that heterolytic bond breaking results in species that are electrophiles or nucleophiles · know the definition of the term 'nucleophile' · understand the link between bond polarity and the type of reaction mechanism a compound will undergo

Alkenes: Structure, Sigma/Pi Bonding and E-Z Isomerism

know 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 sigma bond and a pi bond) · be 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 atoms · understand the E-Z naming system for geometric isomers and why it is necessary to use this when the cis- and trans- naming system breaks down

Optical Isomerism, Enantiomers and Chirality in Synthesis

know that optical isomerism is a result of chirality in molecules with a single chiral centre · understand 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 isomers · know 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 centre · know what is meant by the term 'racemic mixture' · be able to use data on optical activity of reactants and products as evidence for SN1 and SN2 mechanisms and addition to carbonyl compounds

Alkanes: Free-Radical Substitution and Combustion

know 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) · understand the term 'structural isomerism' and be able to draw the structural isomers of organic molecules, given their molecular formula · be able to draw and name the structural isomers of alkanes and cycloalkanes with up to six carbon atoms · know 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 reactions · know that pollutants, including carbon monoxide, oxides of nitrogen and sulfur, carbon particulates and unburned hydrocarbons, are emitted during the combustion of alkane fuels · understand 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 sulfur · be able to discuss the reasons for developing alternative fuels in terms of sustainability and reducing emissions, including the emission of CO2 and its relationship to climate change · be able to apply the concept of carbon neutrality to different fuels, such as petrol, bioethanol and hydrogen · understand the reactions of alkanes with: (i) oxygen in the air (combustion); (ii) halogens · understand 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

Alkenes: Electrophilic Addition and Oxidation

be 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 diol · know the qualitative test for a C=C double bond using bromine or bromine water · be 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.

Arenes: Benzene Structure, Reactions and Side-Chain vs Ring Halogenation

be 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 the Kekulé or the delocalised form as appropriate in equations and mechanisms) · understand that the delocalised model for the structure of benzene involves overlap of p-orbitals to form pi-bonds · understand why benzene is resistant to bromination, compared to alkenes, in terms of delocalisation of pi-bonds in benzene compared to the localised electron density of the pi-bond in alkenes · know 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)

Halogenoalkanes: Nucleophilic Substitution and Elimination

understand the nomenclature of halogenoalkanes and be able to draw their structural, displayed and skeletal formulae · understand the distinction between primary, secondary and tertiary halogenoalkanes · understand 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. · understand the mechanisms of the nucleophilic substitution reactions between primary halogenoalkanes and: (i) aqueous potassium hydroxide; (ii) ammonia. SN1 and SN2 substitution mechanisms will be tested in Unit 4. · understand 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 ethanol · CORE PRACTICAL 5: Investigation of the rates of hydrolysis of some halogenoalkanes · know the trend in reactivity of primary, secondary and tertiary halogenoalkanes · understand, in terms of bond enthalpy, the trend in reactivity of chloro-, bromo- and iodoalkanes · CORE PRACTICAL 6: Chlorination of 2-methylpropan-2-ol with concentrated hydrochloric acid

Halogenoarenes: Formation and Reactivity vs Halogenoalkanes

recall the production of halogenoarenes by catalysed halogenation of arenes · explain the difference in reactivity between halogenoalkanes and halogenoarenes · explain halogenoarene inertness via C-X bond strength and lone-pair delocalisation

Alcohols: Classification, Oxidation and Reactions

understand the nomenclature of alcohols and be able to draw their structural, displayed and skeletal formulae · understand the distinction between primary, secondary and tertiary alcohols · understand the reactions of alcohols with: (i) oxygen in air (combustion); (ii) halogenating agents: PCl5 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. · understand 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]. · understand 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 determination · CORE PRACTICAL 7: The oxidation of propan-1-ol to produce propanal and propanoic acid

Electrophilic Substitution in Arenes and the Chemistry of Phenol

understand the mechanism of the electrophilic substitution reactions of benzene in halogenation, nitration and Friedel-Crafts reactions, including the generation of the electrophile · understand the reaction of phenol with bromine water and the reasons for the relative ease of this reaction compared to benzene

Carbonyl Compounds: Aldehydes and Ketones

understand the nomenclature of aldehydes and ketones and be able to draw their structural, displayed and skeletal formulae · understand 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 solubility · understand 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)

Carboxylic Acids: Preparation, Acidity and Reactions

understand the nomenclature of carboxylic acids and be able to draw their structural, displayed and skeletal formulae · understand that hydrogen bonding affects the physical properties of carboxylic acids, in relation to their boiling temperatures and solubility · understand that carboxylic acids can be prepared by the oxidation of alcohols or aldehydes and the hydrolysis of nitriles · understand 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

Acyl Chlorides and Esters: Addition-Elimination and Hydrolysis

understand the nomenclature of acyl chlorides and esters and be able to draw their structural, displayed and skeletal formulae · understand the reactions of acyl chlorides with: (i) water; (ii) alcohols; (iii) concentrated ammonia; (iv) amines · understand the hydrolysis reactions of esters, in acidic and alkaline solution · understand how polyesters, such as terylene, are formed by condensation polymerisation reactions

Amines: Basicity, Preparation and Azo Dyes

understand the nomenclature of amides, amines and amino acids and be able to draw their structural, displayed and skeletal formulae · understand 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 ion · understand 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 amines · understand, in terms of reagents and general reaction conditions, the preparation of primary aliphatic amines: (i) from halogenoalkanes; (ii) by the reduction of nitriles · know the preparation of aromatic amines by the reduction of aromatic nitro-compounds using tin and concentrated hydrochloric acid · be 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 dye

Amino Acids, Zwitterions, Peptides and Electrophoresis

be 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 polymerisation · CORE PRACTICAL 15: Analysis of some inorganic and organic unknowns

Addition Polymerisation of Alkenes

be able to describe the addition polymerisation of alkenes and draw the repeat unit given the monomer, and vice versa · understand 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

Amides and Condensation Polymers (Polyamides and Polyesters)

understand that amides can be prepared from acyl chlorides · be 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) · be able to draw the structural formulae of the repeat units of the polymers in 19.8 · be 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)

Organic Synthesis: Multi-Step Routes and Structure Determination

be 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 13C and proton) · understand 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 ether · be able to use knowledge of organic chemistry contained 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 hazards

Mass Spectrometry and Infrared Spectroscopy

be 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 patterns · be 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 amines · CORE PRACTICAL 8: Analysis of some inorganic and organic unknowns

NMR and Chromatography

be 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 places · understand that carbon-13, (13C) NMR spectroscopy provides information about the positions of 13C atoms in a molecule · be able to use data from 13C NMR spectroscopy to: (i) predict the different environments for carbon atoms present in a molecule, given values of chemical shift, delta; (ii) justify the number of peaks present in a 13C NMR spectrum in terms of the number of carbon atoms in different environments · be 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, delta; (ii) relate relative peak areas, or ratio number of protons, to the relative numbers of 1H 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 1H atoms in a given molecule · know that chromatography separates components of a mixture using a mobile phase and a stationary phase · be able to calculate Rf values from one-way chromatograms in paper and thin-layer chromatography (TLC) and understand reasons for differences in Rf values · know 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 sport

Practical Skills: Measurement, Titration & Uncertainties