Atomic Structure

Atomic structure refers to the structure of an atom comprising a nucleus (centre) in which the protons (positively charged) and neutrons (neutral) are present. The negatively charged particles called electrons revolve around the centre of the nucleus.

Stoichiometry

Content

• Relative masses of atoms and molecules

• The mole, the Avogadro constant

• The calculation of empirical and molecular formulae

• Reacting masses and volumes (of solutions and gases)

Learning Outcomes

[the term relative formula mass or M r will be used for ionic compounds]

Candidates should be able to:

(a) define the terms relative atomic, isotopic, molecular and formula mass

(b) define the term mole in terms of the Avogadro constant

(c) calculate the relative atomic mass of an element given the relative abundances of its isotopes

(d) define the terms empirical and molecular formula

(e) calculate empirical and molecular formulae using combustion data or composition by mass

(f) write and/or construct balanced equations

(g) perform calculations, including use of the mole concept, involving:

(i) reacting masses (from formulae and equations)

(ii) volumes of gases (e.g. in the burning of hydrocarbons)

(iii) volumes and concentrations of solutions
[when performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question]

(h) deduce stoichiometric relationships from calculations such as those in (g)

Redox

Content

● Redox processes: electron transfer and changes in oxidation number (oxidation state)

Learning Outcomes

Candidates should be able to:

(a) describe and explain redox processes in terms of electron transfer and/or of changes in oxidation number (oxidation state)

Chemical Bonding

Content

• Ionic bonding, metallic bonding, covalent bonding and co-ordinate (dative covalent) bonding

• Shapes of simple molecules and bond angles

• Bond polarities and polarity of molecules

• Intermolecular forces, including hydrogen bonding

• Bond energies and bond lengths

• Lattice structure of solids

• Bonding and physical properties

Learning Outcomes

Candidates should be able to:

(a) show understanding that all chemical bonds are electrostatic in nature and describe:

(i) ionic bond as the electrostatic attraction between oppositely charged ions

(ii) covalent bond as the electrostatic attraction between a shared pair of electrons and positively charged nuclei

(iii) metallic bond as the electrostatic attraction between a lattice of positive ions and delocalised electrons

(b) describe, including the use of ‘dot-and-cross’ diagrams,

(i) ionic bonding as in sodium chloride and magnesium oxide

(ii) covalent bonding as in hydrogen; oxygen; nitrogen; chlorine; hydrogen chloride; carbon dioxide; methane; ethene

(iii) co-ordinate (dative covalent) bonding, as in formation of the ammonium ion and in the Al2 Cl6 molecule

(c) describe covalent bonding in terms of orbital overlap (limited to s and p orbitals only), giving σ and π bonds (see also Section 11.1)

(d) explain the shapes of, and bond angles in, molecules such as BF 3 (trigonal planar); CO 2 (linear); CH4 (tetrahedral); NH 3 (trigonal pyramidal); H2 O (bent); SF 6 (octahedral) by using the Valence Shell Electron Pair Repulsion theory

(e) predict the shapes of, and bond angles in, molecules analogous to those specified in (d)

(f) explain and deduce bond polarity using the concept of electronegativity [quantitative treatment of electronegativity is not required]

(g) deduce the polarity of a molecule using bond polarity and its molecular shape (analogous to those specified in (d));

(h) describe the following forces of attraction (electrostatic in nature):

(i) intermolecular forces, based on permanent and induced dipoles, as in CHCl3 (l); Br2 (l) and the liquid noble gases

(ii) hydrogen bonding, using ammonia and water as examples of molecules containing –NH and –OH groups

(i) outline the importance of hydrogen bonding to the physical properties of substances, including ice and water

(j) explain the terms bond energy and bond length for covalent bonds

(k) compare the reactivities of covalent bonds in terms of bond energy, bond length and bond polarity

(l) describe, in simple terms, the lattice structure of a crystalline solid which is:

(i) ionic, as in sodium chloride and magnesium oxide

(ii) simple molecular, as in iodine

(iii) giant molecular, as in graphite and diamond

(iv) hydrogen-bonded, as in ice

(v) metallic, as in copper

[the concept of the ‘unit cell’ is not required]

(m) describe, interpret and/or predict the effect of different types of structure and bonding on the physical properties of substances

(n) suggest the type of structure and bonding present in a substance from given information

Hybridisation

Learning Outcomes

Candidates should be able to:

(d) describe sp 3 hybridisation, as in ethane molecule, sp2 hybridisation, as in ethene and benzene molecules, and sp hybridisation, as in ethyne molecule

(e) explain the shapes of, and bond angles in, the ethane, ethene, benzene, and ethyne molecules in relation to σ and π carbon-carbon bonds

(f) predict the shapes of, and bond angles in, molecules analogous to those specified in (e)

The Gaseous State

Content

• Ideal gas behaviour and deviations from it

• pV = nRT and its use in determining a value for Mr

• Dalton’s Law and its use in determining the partial pressures of gases in a mixture

Learning Outcomes

Candidates should be able to:

(a) state the basic assumptions of the kinetic theory as applied to an ideal gas

(b) explain qualitatively in terms of intermolecular forces and molecular size:

(i) the conditions necessary for a gas to approach ideal behaviour

(ii) the limitations of ideality at very high pressures and very low temperatures

(c) state and use the general gas equation pV = nRT in calculations, including the determination of Mr

(d) use Dalton’s Law to determine the partial pressures of gases in a mixture (see also Section 9)

Energetics

Content

• Enthalpy changes: ∆H, of formation; combustion; hydration; solution; neutralisation; atomisation; bond energy; lattice energy; electron affinity

• Hess’ Law, including Born-Haber cycles

• Entropy and Free Energy

Learning Outcomes

Candidates should be able to:

(a) explain that most chemical reactions are accompanied by energy changes, principally in the form of heat usually associated with the breaking and forming of chemical bonds; the reaction can be exothermic (∆H negative) or endothermic (∆H positive)

(b) construct and interpret an energy profile diagram, in terms of the enthalpy change of the reaction and of the activation energy (see also Section 8)

(c) explain and use the terms:

(i) enthalpy change of reaction and standard conditions, with particular reference to: formation; combustion; hydration; solution; neutralisation; atomisation

(ii) bond energy (∆H positive, i.e. bond breaking) (see also Section 2)

(iii) lattice energy (∆H negative, i.e. gaseous ions to solid lattice)

(d) calculate enthalpy changes from appropriate experimental results, including the use of the relationship:

heat change = mc∆T

(e) explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy

(f) apply Hess’ Law to construct simple energy cycles, e.g. Born-Haber cycle, and carry out calculations involving such cycles and relevant energy terms (including ionisation energy and electron affinity), with particular reference to:

(i) determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion

(ii) the formation of a simple ionic solid and of its aqueous solution

(iii) average bond energies

(g) explain and use the term entropy

(h)  discuss the effects on the entropy of a chemical system by the following:

(i) change in temperature

(ii) change in phase

(iii) change in the number of particles (especially for gaseous systems)

(iv) mixing of particles

[quantitative treatment is not required]

(i) predict whether the entropy change for a given process or reaction is positive or negative

(j) state and use the equation involving standard Gibbs free energy change of reaction, ∆G⦵ : ∆G ⦵ = ∆H⦵ – T∆S⦵ [the calculation of standard entropy change, ∆S⦵ , for a reaction using standard entropies, S⦵ , is not required]

(k) state whether a reaction or process will be spontaneous by using the sign of ∆G⦵

(l) understand the limitations in the use of ∆G⦵ to predict the spontaneity of a reaction

(m) predict the effect of temperature change on the spontaneity of a reaction, given standard enthalpy and entropy changes