1. LCAO-molecular orbitals oIn MO Theory, atomic orbitals on the constituent atoms are combined to form bonding, non-bonding and anti-bonding orbitals for the molecule H2H2 He 2

2. Valence Bond (“hybridization”) theory oValence Bond (VB) theory is driven by the shape of the molecule (Chapter 10.7) oVB Theory begins with two steps:  hybridization (combination of AOs on the same atom such that new AOs, known as hybrid orbitals, are formed that POINT IN THE RIGHT DIRECTION)  hybrid orbitals and/or AOs on different atoms are combined to make sigma bonds with electron density localized between the two bonding atoms  Pi bonds are formed from unhybridized atomic p orbitals oKey differences between MO and VB theory:  MO theory has electrons distributed over molecule  VB theory localizes an electron pair between two atoms  MO theory combines AOs on DIFFERENT atoms to make MOs (LCAO)  VB theory combines AOs on the SAME atom to make hybridized atomic orbitals (hybridization)  In MO theory, the symmetry (or antisymmetry) must be retained in each orbital.  In VB theory, all orbitals must be looked at at once to see retention of the molecule’s symmetry.

3. H—Be—H and sp hybridization oIn the method of hybrid orbitals, we “create” two hybrid atomic orbitals specifically designed to fit the shape of the molecule, in this case linear, using atomic orbitals of an excited state Be atom! oUnused AOs are left behind as unhybridized atomic orbitals oThe energy of the hybrid atomic orbitals are intermediate between those of the original constituent AO’s oThe hybrid orbitals combine with other orbitals, atomic or hybrid, in the usual fashion, creating both bonding and anti-bonding molecular orbitals, which are localized molecular orbitals

5. BH 3 and sp 2 hybridization oBH 3 is trigonal planar with three equal B—H bonds oTo get this shape, we need to combine the 2s with two 2p AO’s to generate three equivalent hybrid atomic orbitals oCombination with the H 1s leads to bonding and anti-bonding molecular orbitals, which are localized molecular orbitals pointing to the corners of a triangle

7. CH 4 and sp 3 hybridization oCH 4 is tetrahedral with four equal C—H bonds oTo get this shape, we need to combine all the n=2 AO’s to generate four equivalent hybrid atomic orbitals oCombination with the H 1s leads to bonding and anti-bonding molecular orbitals, which are localized molecular orbitals pointing to the corners of a tetrahedron

8. Compare to the LCAO model of CH 4 1σ1σ 2σ2σ 3σ3σ 4σ4σ 2σ2σ 1σ1σ 1σ1σ 2σ2σ 3σ3σ 4σ4σ

9. VB theory is built on VSEPR shapes oRecap: Molecular Geometry and Electron Group Geometry (Chapter 10) sp sp 2 Hybridization:

10. VB theory is built on VSEPR shapes oVB theory is useless beyond 4 electron groups! oThus we need only consider 2, 3, and 4 electron groups from Chapter 10, section 7 sp 3 Hybridization:

11. Summary of the VB (Hybrid AO) method oThe normal atomic orbitals have shapes that do not correspond to chemical bonds oAtomic orbitals are essentially spherical oBonds are oriented towards the Terminal Atoms 1.Hybrid Atomic Orbitals (HAO) are mathematically valid mixtures of the original atomic orbitals 2.They are “manufactured” by promoting electrons to the desired excited state 3.This energy has to be “paid” for – it is obtained by the bond energy that results from good chemical bonds formed when the geometry “fits” 4.Remember that the method results in multiple equal – and hence degenerate – HAO’s that differ only in their orientation oBonds can be formed from the “overlap” between any of the following: 1.HAO with HAO 2.HAO with AO 3.AO with AO

12. Two central atoms: ethane oIf we treat ethane by the VSEPR theory, we find that both carbon atoms are tetrahedral oThe shape of the molecule is shown in the diagram: the only additional information required is the conformation which is adjusted to minimize contacts between the atoms – this is known as the staggered conformation oWe can thus explain the bonding in ethane by using sp 3 hybrid orbitals on each carbon atom oThe H atoms bond using their 1s atomic orbitals oIn all there are 14 electrons or 7 electron pair bonds in the molecule Ignore the “tail ends” of the HAORotation…

13. Bonding in Ethane Lewis structure (including all lone pairs)VSEPR Geometry (including 3D) Hybridization at central atoms

14. Double bonds: ethene oIf we treat ethane by the VSEPR theory, we find that both carbon atoms are trigonal planar oThe shape of the molecule is shown in the diagram: the only additional information required is the conformation which is planar. Why does this geometry occur? oWe can explain the bonding in ethene by using sp 2 hybrid orbitals on each carbon atom, which leaves one atomic p orbital unused on each C atom, while H atoms use their 1s atomic orbitals oIn all there are 6 electron pair bonds in the molecule, 5 in σ orbitals, 1 in the π orbital The sigma skeleton of etheneThe pi bond of ethene

15. Planarity in double bonds: ethene again oWe can now explain the origin of the planar structure of ethene oOnly when the two CH 2 fragments are co-planar can there be efficient overlap between the unhybridized p orbitals leading to the π bond oAs shown in the bottom diagram: if ethene is rotated by 90° along the C—C bond, the atomic p orbitals have zero net overlap oSuch an arrangement is know as an orthogonal interaction of wavefunctions, and does not lead to any net bonding oDouble bonds impose coplanar conformations on the joining atoms oThis is true for all double-bonded molecules, and is a powerful confirmation of the bonding theories that we have developed oNote that a double bond is always the sum of a sigma + a pi bond oSingle bonds are always sigma bonds, so that in ethane, all the bonds are sigma

19. Bonding in Allene (H 2 C=C=CH 2 ) Lewis StructureVSEPR Geometry σ bonding π bonding Hybridization