AS Chemistry₇₄₀₄

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.

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.

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.

Relative atomic mass and relative molecular mass in terms of¹² C.

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

Students should be able to:

• define relative atomic mass ( A _r )
• define relative molecular mass ( M _r ).

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, M _r , 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.

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 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.

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.

Equations (full and ionic).

Percentage atom economy is:

$\frac{\mathit{molecular\; mass\; of\; desired\; product}}{\mathit{sum\; of\; molecular\; masses\; of\; all\; reactants}}\mathsf{\times }\mathsf{100}$

Economic, ethical and environmental advantages for society and for industry of developing chemical processes with a high atom economy.

• 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

Students could be asked to find:

• the concentration of ethanoic acid in vinegar
• the mass of calcium carbonate in an indigestion tablet
• the M _r of MHCO₃
• the M _r of succinic acid
• the mass of aspirin in an aspirin tablet
• the yield for the conversion of magnesium to magnesium oxide
• the M _r 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.

Make up a volumetric solution and carry out a simple acid–base titration.

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.

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.

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

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).

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.

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.

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.

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:

• define standard enthalpy of combustion (∆_c H ^Ɵ )
• define standard enthalpy of formation (∆_f H ^Ɵ ).

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:

• 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 = 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

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

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

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

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.

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.

Maxwell–Boltzmann distribution of molecular energies in gases.

Students should be able to:

• draw and interpret distribution curves for different temperatures.

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.

Investigation of how the rate of a reaction changes with temperature.

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.

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.

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.

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(H₂ O)₆ ²⁺ with concentrated HCl).

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:

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

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.

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.