ACSI 2023 Chemical Bonding and Structure AHL

Year 6 (2023) IBDP Chemistry HL TOPIC 4 Chemical bonding and structure 14.1 14.2 Further aspect of covalent bonding and structure Larger structures and more in-depth explanations of bonding systems often require more sophisticated concepts and theories of bonding. Hybridization Hybridization results from the mixing of atomic orbitals to form the same number of new equivalent hybrid orbitals that can have the same mean energy as the contributing atomic orbitals.

2 14.1 Further aspects of covalent bonding and structure Nature of science: 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) Understandings: • 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. • 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. • Exceptions to the octet rule include some species having incomplete octets and expanded octets. • 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. • 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. Applications and skills: • Prediction whether sigma (σ) or pi (π) bonds are formed from the linear combination of atomic orbitals. • Deduction of the Lewis (electron dot) structures of molecules and ions showing all valence electrons for up to six electron pairs on each atom. • Application of FC to ascertain which Lewis (electron dot) structure is preferred from different Lewis (electron dot) structures. • Deduction using VSEPR theory of the electron domain geometry and molecular geometry with five and six electron domains and associated bond angles. • Explanation of the wavelength of light required to dissociate oxygen and ozone.# • Description of the mechanism of the catalysis of ozone depletion when catalysed by CFCs and NOx.# • # Refer to page 343 of Chemistry Course Companion: An environment perspective: Catalysis of ozone depletion. Guidance: • 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. • Molecular polarities of geometries corresponding to five and six electron domains should also be covered.  Theories of bonding and structure (Refer to page 330 of Chemistry Course Companion)

3 14.2 Hybridization Nature of science: 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) Understandings: • A hybrid orbital results from the mixing of different types of atomic orbitals on the same atom. Applications and skills: • Explanation of the formation of sp3, sp2 and sp hybrid orbitals in methane, ethene and ethyne. • Identification and explanation of the relationships between Lewis (electron dot) structures, electron domains, molecular geometries and types of hybridization. Guidance: • Students need only consider species with sp3, sp2 and sp hybridization.

4 4.3.5 Hybridization Theory Valence bond theory (VBT) and molecular orbital theory (MOT) are basic chemistry theories that were developed to explain chemical bonding using quantum mechanics. In VBT, a covalent bond is formed by the overlap of half-filled valence atomic orbitals of two atoms. These atomic orbitals contain one unpaired electron each. The overlapping of the atomic orbitals result in in the formation of a bonding orbital with electrons localized between the two atoms. In VBT, it is assumed that when the atoms interact they retain their own respective atomic orbitals but the overlapping atomic orbitals can differ. The two types of overlapping orbitals are sigma and pi; which will be elaborated in a later the section. The diagram below shows the formation of a  bond when two 1s orbitals from 2 H atoms overlap. http://www.sparknotes.com/chemistry/organic1/orbitals/section1.rhtml In MOT however, the overlap of atomic orbitals results in the formation of new orbitals known as molecular orbitals. The electrons are assigned to these molecular orbitals, and associated with the whole molecule rather than individual atoms. http://www.sparknotes.com/chemistry/bonding/molecularorbital/section1.rhtml The figure above shows the formation of the molecular orbitals for H2 molecule. The two 1s orbitals on the hydrogen atoms combine to form two new molecular orbitals. One combination results in a bonding molecular orbital (), which is of lower energy, and the other combination forms an anti-bonding molecular orbital (*) which is of higher energy. http://www.sparknotes.com/chemistry/bonding/molecularorbital/section1.rhtml

5 1. Types of Bonds • A Lewis structure is a simple model showing how the valence electrons are distributed in a molecule or a polyatomic ion. But the MOT helps to visualize the difference between a sigma () and a pi () bond. • For atomic orbitals to overlap and form molecular orbitals ✓ they must be relatively close in energy; ✓ and the symmetry of the atomic orbitals must be identical. • When X number of atomic orbitals combine, it will form x number of new molecular orbitals. There are 3 possible combinations: ✓ bonding orbital: sigma () or pi () orbital ✓ anti-bonding orbital: sigma* (*) or pi* (*) orbital ✓ non-bonding situation. (See table 2 on page 335 of Chemistry Course Companion for the different possible combinations of atomic orbitals.) • There are two main ways atomic orbitals can overlap to make a new bond: ✓ Head on (or end-on) overlap resulting in a sigma () bond. ✓ Sideways overlap resulting in a pi () bond. (a) Sigma () bond • Sigma (σ) bond is formed from the head on (axial) overlap of orbitals. • In σ bond, the bonding electrons are localized symmetrically along the inter-nuclear axis of the bonding atoms. • The electron density in a -bond is found mainly in the space between the two nuclei. • There are three different types of -Bond. i. s – s overlap (e.g. H2) ii. s – p overlap (e.g.HCl) H H H or + 1s orbital 3p orbital overlap of orbitals -bond (molecular orbital) Cl Cl Cl + H or H H 1s orbital 1s orbital overlap of orbitals -bond (molecular orbital) H H H A molecular orbital is a new orbital created from the combination of atomic orbitals used to make a bond. Hence, covalent bonds are formed by the overlapping of atomic orbitals.

6 iii. p – p overlap (eg Cl2) Note that a bonding molecular orbital can only hold 2 electrons (like an atomic orbital), and hence a -bond is made up of a doubly-filled -orbital. (b) Pi () bond • Pi () bond is formed from the sideways overlap of parallel p orbitals. It is much weaker than a σ bond. • The  bond electron cloud consists of two lobes, one above and one below the internuclear axis of the bonding atoms. • Each  bond (a molecular orbital) c