3.1 Physical chemistry
3.1.1 Atomic structure
The chemical properties of elements depend on their atomic structure and in particular on the arrangement of electrons around the nucleus. The arrangement of electrons in orbitals is linked to the way in which elements are organised in the Periodic Table. Chemists can measure the mass of atoms and molecules to a high degree of accuracy in a mass spectrometer. The principles of operation of a modern mass spectrometer are studied.
3.1.1.1 Fundamental particles
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Appreciate that knowledge and understanding of atomic structure has evolved over time. Protons, neutrons and electrons: relative charge and relative mass. An atom consists of a nucleus containing protons and neutrons surrounded by electrons. 
3.1.1.2 Mass number and isotopes
Updated
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Mass number (A) and atomic (proton) number (Z). Students should be able to:
The principles of a simple time of flight (TOF) mass spectrometer, limited to ionisation, acceleration to give all ions constant kinetic energy, ion drift, ion detection, data analysis. The mass spectrometer gives accurate information about relative isotopic mass and also about the relative abundance of isotopes. Mass spectrometry can be used to identify elements. Mass spectrometry can be used to determine relative molecular mass. Students should be able to:

MS 1.1 Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. MS 1.2 Students calculate weighted means, eg calculation of an atomic mass based on supplied isotopic abundances. MS 3.1 Students interpret and analyse spectra. 
3.1.1.3 Electron configuration
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Electron configurations of atoms and ions up to Z = 36 in terms of shells and subshells (orbitals) s, p and d. Ionisation energies. Students should be able to:

3.1.2 Amount of substance
When chemists measure out an amount of a substance, they use an amount in moles. The mole is a useful quantity because one mole of a substance always contains the same number of entities of the substance. An amount in moles can be measured out by mass in grams, by volume in dm^{3} of a solution of known concentration and by volume in dm^{3} of a gas.
3.1.2.1 Relative atomic mass and relative molecular mass
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Relative atomic mass and relative molecular mass in terms of ^{12}C. The term relative formula mass will be used for ionic compounds. Students should be able to:

3.1.2.2 The mole and the Avogadro constant
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The Avogadro constant as the number of particles in a mole. The mole as applied to electrons, atoms, molecules, ions, formulas and equations. The concentration of a substance in solution, measured in mol dm^{–3}. Students should be able to carry out calculations:
Students will not be expected to recall the value of the Avogadro constant. 
MS 0.1 Students carry out calculations using numbers in standard and ordinary form, eg using the Avogadro constant. MS 0.4 Students carry out calculations using the Avogadro constant. MS 1.1 Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. Students understand that calculated results can only be reported to the limits of the least accurate measurement. 
3.1.2.3 The ideal gas equation
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The ideal gas equation pV = nRT with the variables in SI units. Students should be able to:
Students will not be expected to recall the value of the gas constant, R. 
AT a, b and k PS 3.2 Students could be asked to find the M_{r} of a volatile liquid. MS 0.0 Students understand that the correct units need to be in pV = nRT. MS 2.2, 2.3 and 2.4 Students carry out calculations with the ideal gas equation, including rearranging the ideal gas equation to find unknown quantities. 
3.1.2.4 Empirical and molecular formula
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Empirical formula is the simplest whole number ratio of atoms of each element in a compound. Molecular formula is the actual number of atoms of each element in a compound. The relationship between empirical formula and molecular formula. Students should be able to:

AT a and k PS 2.3 and 3.3 Students could be asked to find the empirical formula of a metal oxide. 
3.1.2.5 Balanced equations and associated calculations
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Equations (full and ionic). Percentage atom economy is:
Economic, ethical and environmental advantages for society and for industry of developing chemical processes with a high atom economy. Students should be able to:
Students should be able to use balanced equations to calculate:

AT a, d, e, f and k PS 4.1Students could be asked to find:
AT a and k Students could be asked to find the percentage conversion of a Group 2 carbonate to its oxide by heat. AT d, e, f and k Students could be asked to determine the number of moles of water of crystallisation in a hydrated salt by titration. MS 0.2 Students construct and/or balance equations using ratios. Students calculate percentage yields and atom economies of reactions. MS 1.2 and 1.3 Students select appropriate titration data (ie identify outliers) in order to calculate mean titres. Students determine uncertainty when two burette readings are used to calculate a titre value. 
Required practical 1 Make up a volumetric solution and carry out a simple acid–base titration. 
3.1.3 Bonding
The physical and chemical properties of compounds depend on the ways in which the compounds are held together by chemical bonds and by intermolecular forces. Theories of bonding explain how atoms or ions are held together in these structures. Materials scientists use knowledge of structure and bonding to engineer new materials with desirable properties. These new materials may offer new applications in a range of different modern technologies.
3.1.3.1 Ionic bonding
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Ionic bonding involves electrostatic attraction between oppositely charged ions in a lattice. The formulas of compound ions, eg sulfate, hydroxide, nitrate, carbonate and ammonium. Students should be able to:

3.1.3.2 Nature of covalent and dative covalent bonds
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A single covalent bond contains a shared pair of electrons. Multiple bonds contain multiple pairs of electrons. A coordinate (dative covalent) bond contains a shared pair of electrons with both electrons supplied by one atom. Students should be able to represent:

3.1.3.3 Metallic bonding
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Metallic bonding involves attraction between delocalised electrons and positive ions arranged in a lattice. 
3.1.3.4 Bonding and physical properties
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The four types of crystal structure:
The structures of the following crystals as examples of these four types of crystal structure:
Students should be able to:

AT a, b, h and k PS 1.1 Students could be asked to find the type of structure of unknowns by experiment (eg to test solubility, conductivity and ease of melting). 
3.1.3.5 Shapes of simple molecules and ions
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Bonding pairs and lone (nonbonding) pairs of electrons as charge clouds that repel each other. Pairs of electrons in the outer shell of atoms arrange themselves as far apart as possible to minimise repulsion. Lone pair–lone pair repulsion is greater than lone pair–bond pair repulsion, which is greater than bond pair–bond pair repulsion. The effect of electron pair repulsion on bond angles. Students should be able to:

MS 0.3 and 4.1 Students could be given familiar and unfamiliar examples of species and asked to deduce the shape according to valence shell electron pair repulsion (VSEPR) principles. 
3.1.3.6 Bond polarity
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Electronegativity as the power of an atom to attract the pair of electrons in a covalent bond. The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical. This produces a polar covalent bond, and may cause a molecule to have a permanent dipole. Students should be able to:

3.1.3.7 Forces between molecules
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Forces between molecules:
The melting and boiling points of molecular substances are influenced by the strength of these intermolecular forces. The importance of hydrogen bonding in the low density of ice and the anomalous boiling points of compounds. Students should be able to:

AT d and k PS 1.2 Students could try to deflect jets of various liquids from burettes to investigate the presence of different types and relative size of intermolecular forces. 
3.1.4 Energetics
The enthalpy change in a chemical reaction can be measured accurately. It is important to know this value for chemical reactions that are used as a source of heat energy in applications such as domestic boilers and internal combustion engines.
3.1.4.1 Enthalpy change
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Reactions can be endothermic or exothermic. Enthalpy change (∆H) is the heat energy change measured under conditions of constant pressure. Standard enthalpy changes refer to standard conditions, ie 100 kPa and a stated temperature (eg ∆H_{298}^{Ɵ}). Students should be able to:

3.1.4.2 Calorimetry
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The heat change, q, in a reaction is given by the equation q = mc∆T where m is the mass of the substance that has a temperature change ∆T and a specific heat capacity c. Students should be able to:
Students will not be expected to recall the value of the specific heat capacity, c, of a substance. 
MS 0.0 and 1.1 Students understand that the correct units need to be used in q = mc∆T Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. Students understand that calculated results can only be reported to the limits of the least accurate measurement. 
Required practical 2 Measurement of an enthalpy change. 
AT a and k PS 2.4, 3.1, 3.2, 3.3 and 4.1Students could be asked to find ∆H for a reaction by calorimetry. Examples of reactions could include:

3.1.4.3 Applications of Hessâ€™s law
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Hess’s law. Students should be able to:

MS 2.4 Students carry out Hess's law calculations. AT a and k PS 2.4, 3.2 and 4.1Students could be asked to find ∆H for a reaction using Hess’s law and calorimetry, then present data in appropriate ways. Examples of reactions could include:

3.1.4.4 Bond enthalpies
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Mean bond enthalpy. Students should be able to:

MS 1.2 Students understand that bond enthalpies are mean values across a range of compounds containing that bond. 
3.1.5 Kinetics
The study of kinetics enables chemists to determine how a change in conditions affects the speed of a chemical reaction. Whilst the reactivity of chemicals is a significant factor in how fast chemical reactions proceed, there are variables that can be manipulated in order to speed them up or slow them down.
3.1.5.1 Collision theory
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Reactions can only occur when collisions take place between particles having sufficient energy. This energy is called the activation energy. Students should be able to:

3.1.5.2 Maxwellâ€“Boltzmann distribution
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Maxwell–Boltzmann distribution of molecular energies in gases. Students should be able to:

3.1.5.3 Effect of temperature on reaction rate
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Meaning of the term rate of reaction. The qualitative effect of temperature changes on the rate of reaction. Students should be able to:

AT a, b, k and l PS 2.4 and 3.1 Students could investigate the effect of temperature on the rate of reaction of sodium thiosulfate and hydrochloric acid by an initial rate method. Research opportunity Students could investigate how knowledge and understanding of the factors that affect the rate of chemical reaction have changed methods of storage and cooking of food. 
Required practical 3 Investigation of how the rate of a reaction changes with temperature. 
3.1.5.4 Effect of concentration and pressure
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The qualitative effect of changes in concentration on collision frequency. The qualitative effect of a change in the pressure of a gas on collision frequency. Students should be able to:

AT a, e, k and i Students could investigate the effect of changing the concentration of acid on the rate of a reaction of calcium carbonate and hydrochloric acid by a continuous monitoring method. 
3.1.5.5 Catalysts
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A catalyst is a substance that increases the rate of a chemical reaction without being changed in chemical composition or amount. Catalysts work by providing an alternative reaction route of lower activation energy. Students should be able to:

3.1.6 Chemical equilibria, Le Chatelierâ€™s principle and K_{c}
In contrast with kinetics, which is a study of how quickly reactions occur, a study of equilibria indicates how far reactions will go. Le Chatelier’s principle can be used to predict the effects of changes in temperature, pressure and concentration on the yield of a reversible reaction. This has important consequences for many industrial processes. The further study of the equilibrium constant, K_{c}, considers how the mathematical expression for the equilibrium constant enables us to calculate how an equilibrium yield will be influenced by the concentration of reactants and products.
3.1.6.1 Chemical equilibria and Le Chatelier's principle
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Many chemical reactions are reversible. In a reversible reaction at equilibrium:
Le Chatelier’s principle. Le Chatelier's principle can be used to predict the effects of changes in temperature, pressure and concentration on the position of equilibrium in homogeneous reactions. A catalyst does not affect the position of equilibrium. Students should be able to:

PS 1.1 Students could carry out testtube equilibrium shifts to show the effect of concentration and temperature (eg Cu(H_{2}O)_{6}^{2+} with concentrated HCl). 
3.1.6.2 Equilibrium constant K_{c} for homogeneous systems
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The equilibrium constant K_{c} is deduced from the equation for a reversible reaction. The concentration, in mol dm^{–3}, of a species X involved in the expression for K_{c} is represented by [X] The value of the equilibrium constant is not affected either by changes in concentration or addition of a catalyst. Students should be able to:

MS 0.3 Students estimate the effect of changing experimental parameters on a measurable value, eg how the value of K_{c} would change with temperature, given different specified conditions. MS 1.1 Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. Students understand that calculated results can only be reported to the limits of the least accurate measurement. MS 2.2 and 2.3 Students calculate the concentration of a reagent at equilibrium. Students calculate the value of an equilibrium constant K_{c} PS 2.3 Students could determine the equilibrium constant, K_{c}, for the reaction of ethanol with ethanoic acid in the presence of a strong acid catalyst to ethyl ethanoate. 
3.1.7 Oxidation, reduction and redox equations
Redox reactions involve a transfer of electrons from the reducing agent to the oxidising agent. The change in the oxidation state of an element in a compound or ion is used to identify the element that has been oxidised or reduced in a given reaction. Separate halfequations are written for the oxidation or reduction processes. These halfequations can then be combined to give an overall equation for any redox reaction.
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Oxidation is the process of electron loss and oxidising agents are electron acceptors. Reduction is the process of electron gain and reducing agents are electron donors. The rules for assigning oxidation states. Students should be able to:

3.1.8 Thermodynamics (Alevel only)
The further study of thermodynamics builds on the Energetics section and is important in understanding the stability of compounds and why chemical reactions occur. Enthalpy change is linked with entropy change enabling the freeenergy change to be calculated.
3.1.8.1 Bornâ€“Haber cycles (Alevel only)
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Lattice enthalpy can be defined as either enthalpy of lattice dissociation or enthalpy of lattice formation. Born–Haber cycles are used to calculate lattice enthalpies using the following data:
Students should be able to:
Cycles are used to calculate enthalpies of solution for ionic compounds from lattice enthalpies and enthalpies of hydration. Students should be able to:

3.1.8.2 Gibbs freeenergy change, âˆ† G, and entropy change, âˆ† S (Alevel only)
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∆H, whilst important, is not sufficient to explain feasible change. The concept of increasing disorder (entropy change, ∆S). ∆S accounts for the above deficiency, illustrated by physical changes and chemical changes.The balance between entropy and enthalpy determines the feasibility of a reaction given by the relationship: ∆G = ∆H – T∆S (derivation not required).For a reaction to be feasible, the value of ∆G must be zero or negative. Students should be able to:

AT a, b and k PS 3.2 Students could be asked to find ∆S for vaporization of water using a kettle. MS 2.2, 2.3 and 2.4 Students rearrange the equation ∆G = ∆H – T∆S to find unknown values. MS 3.3 Students determine ∆S and ∆H from a graph of ∆G versus T. 
3.1.9 Rate equations (Alevel only)
In rate equations, the mathematical relationship between rate of reaction and concentration gives information about the mechanism of a reaction that may occur in several steps.
3.1.9.1 Rate equations (Alevel only)
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The rate of a chemical reaction is related to the concentration of reactants by a rate equation of the form: Rate = k[A]^{m} [B]^{n} where m and n are the orders of reaction with respect to reactants A and B and k is the rate constant. The orders m and n are restricted to the values 0, 1, and 2. The rate constant k varies with temperature as shown by the equation: k = Ae^{–}^{Ea/RT} where A is a constant, known as the Arrhenius constant, E_{a} is the activation energy and T is the temperature in K. Students should be able to:
These equations and the gas constant, R, will be given when required. 
MS 0.0 and 2.4 Students use given rate data and deduce a rate equation, then use some of the data to calculate the rate constant including units. Rate equations could be given and students asked to calculate rate constant or rate. MS 3.3 and 3.4 Students use a graph of concentration–time and calculate the rate constant of a zeroorder reaction by determination of the gradient. 
3.1.9.2 Determination of rate equation (Alevel only)
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The rate equation is an experimentally determined relationship. The orders with respect to reactants can provide information about the mechanism of a reaction. Students should be able to:

AT a, b, k and l PS 2.4 and 3.1 Students could determine the order of reaction for a reactant in the iodine clock reaction. MS 3.1 Students could be given data to plot and interpret in terms of order with respect to a reactant. Alternatively, students could just be given appropriate graphs and asked to derive order(s). MS 3.3 and 3.4 Students calculate the rate constant of a zeroorder reaction by determining the gradient of a concentration–time graph. MS 3.5 Students plot concentration–time graphs from collected or supplied data and draw an appropriate bestfit curve. Students draw tangents to such curves to deduce rates at different times. 
Required practical 7Measuring the
rate of reaction:

3.1.10 Equilibrium constant K_{p} for homogeneous systems (Alevel only)
The further study of equilibria considers how the mathematical expression for the equilibrium constant K_{p} enables us to calculate how an equilibrium yield will be influenced by the partial pressures of reactants and products. This has important consequences for many industrial processes.
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The equilibrium constant K_{p} is deduced from the equation for a reversible reaction occurring in the gas phase. K_{p} is the equilibrium constant calculated from partial pressures for a system at constant temperature. Students should be able to:

MS 1.1
Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. Students understand that calculated results can only be reported to the limits of the least accurate measurement. MS 2.2 and 2.3 Students calculate the partial pressures of reactants and products at equilibrium. Students calculate the value of an equilibrium constant K_{p} 
3.1.11 Electrode potentials and electrochemical cells (Alevel only)
Redox reactions take place in electrochemical cells where electrons are transferred from the reducing agent to the oxidising agent indirectly via an external circuit. A potential difference is created that can drive an electric current to do work. Electrochemical cells have very important commercial applications as a portable supply of electricity to power electronic devices such as mobile phones, tablets and laptops. On a larger scale, they can provide energy to power a vehicle.
3.1.11.1 Electrode potentials and cells (Alevel only)
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IUPAC convention for writing halfequations for electrode reactions. The conventional representation of cells. Cells are used to measure electrode potentials by reference to the standard hydrogen electrode. The importance of the conditions when measuring the electrode potential, E (Nernst equation not required). Standard electrode potential, E^{Ɵ}, refers to conditions of 298 K, 100 kPa and 1.00 mol dm^{−3} solution of ions. Standard electrode potentials can be listed as an electrochemical series. Students should be able to:

AT j and k PS 1.1 Students could make simple cells and use them to measure unknown electrode potentials. AT a, b, j and k PS 2.1 and 2.4 Students could be asked to plan and carry out an experiment to investigate the effect of changing conditions, such as concentration or temperature, in a voltaic cell such as ZnZn^{2+}Cu^{2+}Cu AT j and k PS 2.2 Students could use E^{Ɵ} values to predict the direction of simple redox reactions, then test these predictions by simple testtube reactions. 
Required practical 8 Measuring the EMF of an electrochemical cell. 
3.1.11.2 Commercial applications of electrochemical cells (Alevel only)
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Electrochemical cells can be used as a commercial source of electrical energy. The simplified electrode reactions in a lithium cell: Positive electrode: Li^{+} + CoO_{2} + e^{–} → Li^{+}[CoO_{2}]^{–} Negative electrode: Li → Li^{+} + e^{–} Cells can be nonrechargeable (irreversible), rechargeable or fuel cells. Fuel cells are used to generate an electric current and do not need to be electrically recharged. The electrode reactions in an alkaline hydrogen–oxygen fuel cell. The benefits and risks to society associated with using these cells. Students should be able to:

Research opportunity Students could investigate how knowledge and understanding of electrochemical cells has evolved from the first voltaic battery. 
3.1.12 Acids and bases (Alevel only)
Acids and bases are important in domestic, environmental and industrial contexts. Acidity in aqueous solutions is caused by hydrogen ions and a logarithmic scale, pH, has been devised to measure acidity. Buffer solutions, which can be made from partially neutralised weak acids, resist changes in pH and find many important industrial and biological applications.
3.1.12.1 BrÃ¸nstedâ€“Lowry acidâ€“base equilibria in aqueous solution (Alevel only)
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An acid is a proton donor. A base is a proton acceptor. Acid–base equilibria involve the transfer of protons. 
3.1.12.2 Definition and determination of pH (Alevel only)
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The concentration of hydrogen ions in aqueous solution covers a very wide range. Therefore, a logarithmic scale, the pH scale, is used as a measure of hydrogen ion concentration. pH = –log_{10}[H^{+}] Students should be able to:

MS 0.4 Students carry out pH calculations. MS 2.5 Students could be given concentration values and asked to calculate pH or vice versa. 
3.1.12.3 The ionic product of water, K_{w} (Alevel only)
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Water is slightly dissociated. K_{w} is derived from the equilibrium constant for this dissociation. K_{w} = [H^{+}][OH^{–}] The value of K_{w} varies with temperature. Students should be able to:

MS 0.1 Students use an appropriate number of decimal places in pH calculations. Students understand standard form when applied to areas such as (but not limited to) K_{w} MS 2.2 Students use K_{w} = [H^{+}][OH^{–}] to find the pH of strong bases. 
3.1.12.4 Weak acids and bases K_{a} for weak acids (Alevel only)
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Weak acids and weak bases dissociate only slightly in aqueous solution. K_{a} is the dissociation constant for a weak acid. pK_{a} = –log_{10} K_{a} Students should be able to:

MS 0.0 Students carry out pK_{a} calculations and give appropriate units. MS 0.1 Students understand standard form when applied to areas such as (but not limited to) K_{a} AT a, c, d, e, f and k PS 2.3 Students could calculate K_{a} of a weak acid by measuring the pH at half neutralisation. 
3.1.12.5 pH curves, titrations and indicators (Alevel only)
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Titrations of acids with bases. Students should be able to:
Typical pH curves for acid–base titrations in all combinations of weak and strong monoprotic acids and bases. Students should be able to:

MS 3.2 AT a, c, d and k PS 3.2 and 4.1Students could plot pH curves to show how pH changes during reactions. 
Required practical 9 Investigate how pH changes when a weak acid reacts with a strong base and when a strong acid reacts with a weak base. 
3.1.12.6 Buffer action (Alevel only)
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A buffer solution maintains an approximately constant pH, despite dilution or addition of small amounts of acid or base. Acidic buffer solutions contain a weak acid and the salt of that weak acid. Basic buffer solutions contain a weak base and the salt of that weak base. Applications of buffer solutions. Students should be able to:

AT a, c, e and k PS 1.1 Students could be asked to prepare and test a buffer solution with a specific pH value. MS 0.4 Students make appropriate mathematical approximations in buffer calculations. 
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