4.2.NoS1 Looking for trends and discrepancies—compounds containing non-metals have different properties than compounds that contain non-metals and metals. (2.5)

4.2.NoS2 Use theories to explain natural phenomena—Lewis introduced a class of compounds which share electrons. Pauling used the idea of electronegativity to explain unequal sharing of electrons. (2.2)

4.2.U1 A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei.

4.2.U2 Single, double and triple covalent bonds involve one, two and three shared pairs of electrons respectively.

4.2.U3 Bond length decreases and bond strength increases as the number of shared electrons increases.

4.2.U4 Bond polarity results from the difference in electronegativities of the bonded atoms.

4.2.AS1 Deduction of the polar nature of a covalent bond from electronegativity values.

4.2.G1 Bond polarity can be shown either with partial charges, dipoles or vectors.

4.2.G2 Electronegativity values are given in the data booklet in section 8.

4.2.Uz1 Microwaves—cooking with polar molecules.

4.2.Aims1 Aim 3: Use naming conventions to name covalently bonded compounds.

4.3.NoS Scientists use models as representations of the real world—the development of the model of molecular shape (VSEPR) to explain observable properties. (1.10)

4.3.U1 Lewis (electron dot) structures show all the valence electrons in a covalently bonded species.

4.3.U2 The “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8 electrons.

4.3.U3 Some atoms, like Be and B, might form stable compounds with incomplete octets of electrons.

4.3.U4 Resonance structures occur when there is more than one possible position for a double bond in a molecule.

4.3.U5 Shapes of species are determined by the repulsion of electron pairs according to VSEPR theory.

4.3.U6 Carbon and silicon form giant covalent/network covalent structures.

4.3.AS1 Deduction of Lewis (electron dot) structure of molecules and ions showing all valence electrons for up to four electron pairs on each atom.

4.3.AS2 The use of VSEPR theory to predict the electron domain geometry and the molecular geometry for species with two, three and four electron domains.

4.3.AS3 Prediction of bond angles from molecular geometry and presence of non- bonding pairs of electrons.

4.3.AS4 Prediction of molecular polarity from bond polarity and molecular geometry.

4.3.AS5 Deduction of resonance structures, examples include but are not limited to C_{6}H_{6}, CO_{3}^{2-} and O_{3}.

4.3.AS6 Explanation of the properties of giant covalent compounds in terms of their structures.

4.3.G1 The term “electron domain” should be used in place of “negative charge centre”.

4.3.G2 Electron pairs in a Lewis (electron dot) structure can be shown as dots, crosses, a dash or any combination.

4.3.G3 Allotropes of carbon (diamond, graphite, graphene, C60 buckminsterfullerene) and SiO2 should be covered.

4.3.G4 Coordinate covalent bonds should be covered.

4.3.ToK1 Does the need for resonance structures decrease the value or validity of Lewis (electron dot) theory? What criteria do we use in assessing the validity of a scientific theory?

4.3.Aims1 Aim 7: Computer simulations could be used to model VSEPR structures

14.1.NoS Principle of Occam’s razor—bonding theories have been modified over time. Newer theories need to remain as simple as possible while maximizing explanatory power, for example the idea of formal charge. (2.7)

14.1.U1 Covalent bonds result from the overlap of atomic orbitals. A sigma bond (σ) is formed by the direct head-on/end-to-end overlap of atomic orbitals, resulting in electron density concentrated between the nuclei of the bonding atoms. A pi bond (π) is formed by the sideways overlap of atomic orbitals, resulting in electron density above and below the plane of the nuclei of the bonding atoms.

14.1.U2 Formal charge (FC) can be used to decide which Lewis (electron dot) structure is preferred from several. The FC is the charge an atom would have if all atoms in the molecule had the same electronegativity. FC = (Number of valence electrons)-½(Number of bonding electrons)-(Number of non-bonding electrons). The Lewis (electron dot) structure with the atoms having FC values closest to zero is preferred.

14.1.U3 Exceptions to the octet rule include some species having incomplete octets and expanded octets.

14.1.U4 Delocalization involves electrons that are shared by/between all atoms in a molecule or ion as opposed to being localized between a pair of atoms.

14.1.U5 Resonance involves using two or more Lewis (electron dot) structures to represent a particular molecule or ion. A resonance structure is one of two or more alternative Lewis (electron dot) structures for a molecule or ion that cannot be described fully with one Lewis (electron dot) structure alone.

14.1.AS1 Prediction whether sigma (σ) or pi (π) bonds are formed from the linear combination of atomic orbitals.

14.1.AS2 Deduction of the Lewis (electron dot) structures of molecules and ions showing all valence electrons for up to six electron pairs on each atom.

14.1.AS3 Application of FC to ascertain which Lewis (electron dot) structure is preferred from different Lewis (electron dot) structures.

14.1.AS4 Deduction using VSEPR theory of the electron domain geometry and molecular geometry with five and six electron domains and associated bond angles.

14.1.AS5 Explanation of the wavelength of light required to dissociate oxygen and ozone

14.1.AS6 Description of the mechanism of the catalysis of ozone depletion when catalysed by CFCs and NOx.

14.1.G1 The linear combination of atomic orbitals to form molecular orbitals should be covered in the context of the formation of sigma (σ) and pi (π) bonds.

14.1.G2 Molecular polarities of geometries corresponding to five and six electron domains should also be covered.

14.1.IM1 How has ozone depletion changed over time? What have we done as a global community to reduce ozone depletion?

14.1.IM2 To what extent is ozone depletion an example of both a success and a failure for solving an international environmental concern?

14.1.ToK1 Covalent bonding can be described using valence bond or molecular orbital theory. To what extent is having alternative ways of describing the same phenomena a strength or a weakness?

14.1.Uz1 Drug action and links to a molecule’s structure.

14.1.Uz2 Vision science and links to a molecule’s structure.

14.1.Aims1 Aim 1: Global impact of ozone depletion.

14.1.Aims2 Aim 7: Computer simulations can be used to model structures predicted by VSEPR theory

14.1.Aims3 Aim 8: Moral, ethical, social, economic and environmental implications of ozone depletion and its solution.

14.2.NoS The need to regard theories as uncertain—hybridization in valence bond theory can help explain molecular geometries, but is limited. Quantum mechanics involves several theories explaining the same phenomena, depending on specific requirements. (2.2)

14.2.U1 A hybrid orbital results from the mixing of different types of atomic orbitals on the same atom

14.2.AS1 Explanation of the formation of sp3, sp2 and sp hybrid orbitals in methane, ethene and ethyne.

14.2.AS2 Identification and explanation of the relationships between Lewis (electron dot)structures, electron domains, molecular geometries and types of hybridization.

14.2.G1 Students need only consider species with sp3, sp2 and sp hybridization.

14.2.ToK1 Hybridization is a mathematical device which allows us to relate the bonding in a molecule to its symmetry. What is the relationship between the natural sciences, mathematics and the natural world? Which role does symmetry play in the different areas of knowledge?

14.2.Aims1 Aim 7: Computer simulations could be used to model hybrid orbitals.