RI 2022 Chemical Bonding II v2.0

Raffles Institution Year 5 H2 Chemistry 2022 Lecture Notes 8 - Chemical Bonding II (Hybridisation and R,:sonance) Learning outcomes Students should be able to: (a) describe sp3 hybridisation, as in ethane molecule, sp2 hybridisation, as in ethene and benzene molecules, and sp hybridisation, as in ethyne molecule (b) explain the shapes of, and bond angles in, the ethane, ethene, benzene, and ethyne molecules in relation to cr and 1t carbon-carbon bonds (c) predict the shapes of, and bond angles in, molecules analogous to those specified in (b) (d) interpret, and use the term "delocalisation" associated with organic reactivities. Contents 1. Hybridisation ........................................................................................................................ 1 1.1 sp3 hybridisation of carbon in methane, CH4 ............................................................................................... 2 1.2 sp2 hybridisation of carbon in ethene, C2H4 .................................................................. 3 1.3 sp hybridisation of carbon in ethyne, C2H2 .................................................................... 4 1.4 Effect of Hybridisation on Bond Length and Bond Strength ......................................... 7 2. Resonance .......................................................................................................................... 8 3. Structures of some giant molecular structures and their properties ................................. 12 1 Hybridisation • The theory of covalent bond formation via overlapping of atomic orbitals does not always give good agreement with observations . • Hybridisation is a concept used to explain the observed shapes of molecules ( determined by experimental studies) . • On hybridisation , a set of atomic orbitals are mixed to generate a set of equivalent hybrid orbitals with the same shape and energy. (Note: number of hybrid orbitals formed always equals number of atomic orbitals used for mixing.) • Consider the electron ic configuration of carbon at ground state (1 s2 2s2 2p2). (i) (ii) (iii) To form a normal covalent bond, an orbital containing one electron is required . These orbitals overlap to form a covalent bond. Carbon in its ground state contains only two singly occupied orbitals, thus it should form 2 covalent bonds. However, carbon tends to form 4 covalent bonds . To have four singly occupied orbitals for forming 4 covalent bonds , an electron must first be promoted from the 2s orbital to the empty 2p orbital. 1s 2s 2p C in ground state [!] I 1 11 I C in excited state 1s [I] 2s [I] 2p 11 11 11 I Carbon now has 4 unpaired electrons and can form 4 covalent bonds, but the atomic orbitals do not point in the correct direction for bonding . (iv) For example , CH4 is tetrahedral with H-C-H bond angles of 109.5°, but the p orbitals (Px, py, p,) are perpendicular (at 90°) to one another . (v) To explain the fou