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

Content

Opportunities for skills development

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

Content

Opportunities for skills development

Mass number (A) and atomic (proton) number (Z).

Students should be able to:

  • determine the number of fundamental particles in atoms and ions using mass number, atomic number and charge
  • explain the existence of isotopes.

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:

  • interpret simple mass spectra of elements
  • calculate relative atomic mass from isotopic abundance, limited to mononuclear ions.

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

Content

Opportunities for skills development

Electron configurations of atoms and ions up to Z = 36 in terms of shells and sub-shells (orbitals) s, p and d.

Ionisation energies.

Students should be able to:

  • define first ionisation energy
  • write equations for first and successive ionisation energies
  • explain how first and successive ionisation energies in Period 3 (Na–Ar) and in Group 2 (Be–Ba) give evidence for electron configuration in sub-shells and in shells.
 

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 dm3 of a solution of known concentration and by volume in dm3 of a gas.

3.1.2.1 Relative atomic mass and relative molecular mass

Content

Opportunities for skills development

Relative atomic mass and relative molecular mass in terms of 12C.

The term relative formula mass will be used for ionic compounds.

Students should be able to:

  • define relative atomic mass (Ar)
  • define relative molecular mass (Mr).
 

3.1.2.2 The mole and the Avogadro constant

Content

Opportunities for skills development

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:

  • using the Avogadro constant
  • using mass of substance, Mr, and amount in moles
  • using concentration, volume and amount of substance in a solution.

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

Content

Opportunities for skills development

The ideal gas equation pV = nRT with the variables in SI units.

Students should be able to:

  • use the equation in calculations.

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 Mr 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

Content

Opportunities for skills development

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:

  • calculate empirical formula from data giving composition by mass or percentage by mass
  • calculate molecular formula from the empirical formula and relative molecular mass.

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

Content

Opportunities for skills development

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:
  • write balanced equations for reactions studied
  • balance equations for unfamiliar reactions when reactants and products are specified.

Students should be able to use balanced equations to calculate:

  • masses
  • volumes of gases
  • percentage yields
  • percentage atom economies
  • concentrations and volumes for reactions in solutions.

AT a, d, e, f and k

PS 4.1

Students could be asked to find:

  • the concentration of ethanoic acid in vinegar
  • the mass of calcium carbonate in an indigestion tablet
  • the Mr of MHCO3
  • the Mr of succinic acid
  • the mass of aspirin in an aspirin tablet
  • the yield for the conversion of magnesium to magnesium oxide
  • the Mr of a hydrated salt (eg magnesium sulfate) by heating to constant mass.

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

Content

Opportunities for skills development

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:

  • predict the charge on a simple ion using the position of the element in the Periodic Table
  • construct formulas for ionic compounds.
 

3.1.3.2 Nature of covalent and dative covalent bonds

Content

Opportunities for skills development

A single covalent bond contains a shared pair of electrons.

Multiple bonds contain multiple pairs of electrons.

A co-ordinate (dative covalent) bond contains a shared pair of electrons with both electrons supplied by one atom.

Students should be able to represent:

  • a covalent bond using a line
  • a co-ordinate bond using an arrow.
 

3.1.3.3 Metallic bonding

Content

Opportunities for skills development

Metallic bonding involves attraction between delocalised electrons and positive ions arranged in a lattice.

 

3.1.3.4 Bonding and physical properties

Content

Opportunities for skills development

The four types of crystal structure:

  • ionic
  • metallic
  • macromolecular (giant covalent)
  • molecular.

The structures of the following crystals as examples of these four types of crystal structure:

  • diamond
  • graphite
  • ice
  • iodine
  • magnesium
  • sodium chloride.

Students should be able to:

  • relate the melting point and conductivity of materials to the type of structure and the bonding present
  • explain the energy changes associated with changes of state
  • draw diagrams to represent these structures involving specified numbers of particles.

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

Content

Opportunities for skills development

Bonding pairs and lone (non-bonding) 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:

  • explain the shapes of, and bond angles in, simple molecules and ions with up to six electron pairs (including lone pairs of electrons) surrounding the central atom.

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

Content

Opportunities for skills development

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:

  • use partial charges to show that a bond is polar
  • explain why some molecules with polar bonds do not have a permanent dipole.
 

3.1.3.7 Forces between molecules

Content

Opportunities for skills development

Forces between molecules:

  • permanent dipole–dipole forces
  • induced dipole–dipole (van der Waals, dispersion, London) forces
  • hydrogen bonding.

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:

  • explain the existence of these forces between familiar and unfamiliar molecules
  • explain how melting and boiling points are influenced by these intermolecular forces.

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

Content

Opportunities for skills development

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 ∆H298Ɵ).

Students should be able to:

  • define standard enthalpy of combustion (∆cHƟ)
  • define standard enthalpy of formation (∆fHƟ).
 

3.1.4.2 Calorimetry

Content

Opportunities for skills development

The heat change, q, in a reaction is given by the equation q = mcT

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:

  • use this equation to calculate the molar enthalpy change for a reaction
  • use this equation in related calculations.

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 = mcT

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:
  • dissolution of potassium chloride
  • dissolution of sodium carbonate
  • neutralising NaOH with HCl
  • displacement reaction between CuSO4 + Zn
  • combustion of alcohols.

3.1.4.3 Applications of Hess’s law

Content

Opportunities for skills development

Hess’s law.

Students should be able to:

  • use Hess’s law to perform calculations, including calculation of enthalpy changes for reactions from enthalpies of combustion or from enthalpies of formation.

MS 2.4

Students carry out Hess's law calculations.

AT a and k

PS 2.4, 3.2 and 4.1

Students 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:

  • thermal decomposition of NaHCO3
  • hydration of MgSO4
  • hydration of CuSO4

3.1.4.4 Bond enthalpies

Content

Opportunities for skills development

Mean bond enthalpy.

Students should be able to:

  • define the term mean bond enthalpy
  • use mean bond enthalpies to calculate an approximate value of ∆H for reactions in the gaseous phase
  • explain why values from mean bond enthalpy calculations differ from those determined using Hess’s law.

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

Content

Opportunities for skills development

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:

  • define the term activation energy
  • explain why most collisions do not lead to a reaction.
 

3.1.5.2 Maxwell–Boltzmann distribution

Content

Opportunities for skills development

Maxwell–Boltzmann distribution of molecular energies in gases.

Students should be able to:

  • draw and interpret distribution curves for different temperatures.
 

3.1.5.3 Effect of temperature on reaction rate

Content

Opportunities for skills development

Meaning of the term rate of reaction.

The qualitative effect of temperature changes on the rate of reaction.

Students should be able to:

  • use the Maxwell–Boltzmann distribution to explain why a small temperature increase can lead to a large increase in rate.

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

Content

Opportunities for skills development

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:

  • explain how a change in concentration or a change in pressure influences the rate of a reaction.
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

Content

Opportunities for skills development

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:

  • use a Maxwell–Boltzmann distribution to help explain how a catalyst increases the rate of a reaction involving a gas.
 

3.1.6 Chemical equilibria, Le Chatelier’s principle and Kc

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, Kc, 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

Content

Opportunities for skills development

Many chemical reactions are reversible.

In a reversible reaction at equilibrium:

  • forward and reverse reactions proceed at equal rates
  • the concentrations of reactants and products remain constant

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:

  • use Le Chatelier’s principle to predict qualitatively the effect of changes in temperature, pressure and concentration on the position of equilibrium
  • explain why, for a reversible reaction used in an industrial process, a compromise temperature and pressure may be used.

PS 1.1

Students could carry out test-tube equilibrium shifts to show the effect of concentration and temperature (eg Cu(H2O)62+ with concentrated HCl).

3.1.6.2 Equilibrium constant Kc for homogeneous systems

Content

Opportunities for skills development

The equilibrium constant Kc is deduced from the equation for a reversible reaction.

The concentration, in mol dm–3, of a species X involved in the expression for Kc 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:

  • construct an expression for Kc for a homogeneous system in equilibrium
  • calculate a value for Kc from the equilibrium concentrations for a homogeneous system at constant temperature
  • perform calculations involving Kc
  • predict the qualitative effects of changes of temperature on the value of Kc

MS 0.3

Students estimate the effect of changing experimental parameters on a measurable value, eg how the value of Kc 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 Kc

PS 2.3

Students could determine the equilibrium constant, Kc, 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 half-equations are written for the oxidation or reduction processes. These half-equations can then be combined to give an overall equation for any redox reaction.

Content

Opportunities for skills development

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:

  • work out the oxidation state of an element in a compound or ion from the formula
  • write half-equations identifying the oxidation and reduction processes in redox reactions
  • combine half-equations to give an overall redox equation.
 

3.1.8 Thermodynamics (A-level 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 free-energy change to be calculated.

3.1.8.1 Born–Haber cycles (A-level only)

Content

Opportunities for skills development

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:

  • enthalpy of formation
  • ionisation energy
  • enthalpy of atomisation
  • bond enthalpy
  • electron affinity.

Students should be able to:

  • define each of the above terms and lattice enthalpy
  • construct Born–Haber cycles to calculate lattice enthalpies using these enthalpy changes
  • construct Born–Haber cycles to calculate one of the other enthalpy changes
  • compare lattice enthalpies from Born–Haber cycles with those from calculations based on a perfect ionic model to provide evidence for covalent character in ionic compounds.

Cycles are used to calculate enthalpies of solution for ionic compounds from lattice enthalpies and enthalpies of hydration.

Students should be able to:

  • define the term enthalpy of hydration
  • perform calculations of an enthalpy change using these cycles.
 

3.1.8.2 Gibbs free-energy change, ∆ G, and entropy change, ∆ S (A-level only)

Content

Opportunities for skills development

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 = ∆HTS (derivation not required).

For a reaction to be feasible, the value of ∆G must be zero or negative.

Students should be able to:

  • calculate entropy changes from absolute entropy values
  • use the relationship

    G = ∆HTS to determine how ∆G varies with temperature

  • use the relationship ∆G = ∆HTS to determine the temperature at which a reaction becomes feasible.

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 = ∆HTS to find unknown values.

MS 3.3

Students determine ∆S and ∆H from a graph of ∆G versus T.

3.1.9 Rate equations (A-level 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 (A-level only)

Content

Opportunities for skills development

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 = AeEa/RT

where A is a constant, known as the Arrhenius constant, Ea is the activation energy and T is the temperature in K.

Students should be able to:

  • define the terms order of reaction and rate constant
  • perform calculations using the rate equation
  • explain the qualitative effect of changes in temperature on the rate constant k
  • perform calculations using the equation k = Ae–Ea/RT
  • understand that the equation k = Ae–Ea/RT can be rearranged into the form ln k = –Ea /RT + ln A and know how to use this rearranged equation with experimental data to plot a straight line graph with slope –Ea/R

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 zero-order reaction by determination of the gradient.

3.1.9.2 Determination of rate equation (A-level only)

Content

Opportunities for skills development

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:

  • use concentration–time graphs to deduce the rate of a reaction
  • use initial concentration–time data to deduce the initial rate of a reaction
  • use rate–concentration data or graphs to deduce the order (0, 1 or 2) with respect to a reactant
  • derive the rate equation for a reaction from the orders with respect to each of the reactants
  • use the orders with respect to reactants to provide information about the rate determining/limiting step of a reaction.

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 zero-order 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 best-fit curve.

Students draw tangents to such curves to deduce rates at different times.

Required practical 7Measuring the rate of reaction:
  • by an initial rate method
  • by a continuous monitoring method.
 

3.1.10 Equilibrium constant Kp for homogeneous systems (A-level only)

The further study of equilibria considers how the mathematical expression for the equilibrium constant Kp 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.

Content

Opportunities for skills development

The equilibrium constant Kp is deduced from the equation for a reversible reaction occurring in the gas phase.

Kp is the equilibrium constant calculated from partial pressures for a system at constant temperature.

Students should be able to:

  • derive partial pressure from mole fraction and total pressure
  • construct an expression for Kp for a homogeneous system in equilibrium
  • perform calculations involving Kp
  • predict the qualitative effects of changes in temperature and pressure on the position of equilibrium
  • predict the qualitative effects of changes in temperature on the value of Kp
  • understand that, whilst a catalyst can affect the rate of attainment of an equilibrium, it does not affect the value of the equilibrium 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.

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 Kp

3.1.11 Electrode potentials and electrochemical cells (A-level 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 (A-level only)

Content

Opportunities for skills development

IUPAC convention for writing half-equations 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, EƟ, 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:

  • use EƟ values to predict the direction of simple redox reactions
  • calculate the EMF of a cell
  • write and apply the conventional representation of a cell.

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 Zn|Zn2+||Cu2+|Cu

AT j and k

PS 2.2

Students could use EƟ values to predict the direction of simple redox reactions, then test these predictions by simple test-tube reactions.

Required practical 8

Measuring the EMF of an electrochemical cell.

 

3.1.11.2 Commercial applications of electrochemical cells (A-level only)

Content

Opportunities for skills development

Electrochemical cells can be used as a commercial source of electrical energy.

The simplified electrode reactions in a lithium cell:

Positive electrode: Li+ + CoO2 + e → Li+[CoO2]

Negative electrode: Li → Li+ + e

Cells can be non-rechargeable (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:

  • use given electrode data to deduce the reactions occurring in non-rechargeable and rechargeable cells
  • deduce the EMF of a cell
  • explain how the electrode reactions can be used to generate an electric current.
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 (A-level 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 Brønsted–Lowry acid–base equilibria in aqueous solution (A-level only)

Content

Opportunities for skills development

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 (A-level only)

Content

Opportunities for skills development

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 = –log10[H+]

Students should be able to:

  • convert concentration of hydrogen ions into pH and vice versa
  • calculate the pH of a solution of a strong acid from its concentration.

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, Kw (A-level only)

Content

Opportunities for skills development

Water is slightly dissociated.

Kw is derived from the equilibrium constant for this dissociation.

Kw = [H+][OH]

The value of Kw varies with temperature.

Students should be able to:

  • use Kw to calculate the pH of a strong base from its concentration.

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) Kw

MS 2.2

Students use Kw = [H+][OH] to find the pH of strong bases.

3.1.12.4 Weak acids and bases Ka for weak acids (A-level only)

Content

Opportunities for skills development

Weak acids and weak bases dissociate only slightly in aqueous solution.

Ka is the dissociation constant for a weak acid.

pKa = –log10 Ka

Students should be able to:

  • construct an expression for Ka
  • perform calculations relating the pH of a weak acid to the concentration of the acid and the dissociation constant, Ka
  • convert Ka into pKa and vice versa.

MS 0.0

Students carry out pKa calculations and give appropriate units.

MS 0.1

Students understand standard form when applied to areas such as (but not limited to) Ka

AT a, c, d, e, f and k

PS 2.3

Students could calculate Ka of a weak acid by measuring the pH at half neutralisation.

3.1.12.5 pH curves, titrations and indicators (A-level only)

Content

Opportunities for skills development

Titrations of acids with bases.

Students should be able to:

  • perform calculations for these titrations based on experimental results.

Typical pH curves for acid–base titrations in all combinations of weak and strong monoprotic acids and bases.

Students should be able to:

  • sketch and explain the shapes of typical pH curves
  • use pH curves to select an appropriate indicator.

MS 3.2

AT a, c, d and k

PS 3.2 and 4.1

Students 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 (A-level only)

Content

Opportunities for skills development

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:

  • explain qualitatively the action of acidic and basic buffers
  • calculate the pH of acidic buffer solutions.

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.