1. Molecular Orbitals

2. Organic compounds are organised into groups according to similarities and differences in their structure. Groups of atoms within a molecule give the molecule specific characteristic properties. These groups of atoms are known as FUNCTIONAL GROUPS.

3. Compounds containing only carbon and hydrogen. Hydrocarbons Sub-divisions. Alkanes Alkenes Alkynes Cycloalkanes

4. Saturated Alkanes General formula C n H 2n+2 All carbon atoms form 4 single covalent bonds 4 bonding pairs of electrons round each C These pairs mutually repel and take up a TETRAHEDRAL arrangement C H HHH

5. Alkanes Carbon atom Ground state Electron arrangement 1s 2 2s 2 2p 2 The 2p 2 electrons occupy 2 of the 3 p orbital (Hund’s Rule) 2 unpaired electrons  only 2 covalent bonds? No! Methane = CH 4 => 4 bonds 1s 2 2p 2 2s 2

6. This is explained by the idea of Hybrid orbitals. Hybrid orbitals are the result of ‘mixing’ orbitals to produce a new set of orbitals. (Mathematical basis for this doing this is based on the Schr Ö dinger wave equations – not needed at this level!) Spectroscopic measurements show that all four bonds in methane are identical. Alkanes

7. One of the 2s electron is promoted to the third p orbital. This produces 4 Hybrid orbitals of equal energy containing single electrons. This is known as sp 3 hybridisation Atom after hybridisation sp 3 Atom before hybridisation s p Alkanes

8. Atom after hybridisation sp 3 Atom before hybridisation s p Compare the shape of the p orbital for carbon with that of the sp 3 orbital. The hybrid orbital is more directional in shape  better overlap when forming bonds + stronger bonds + forms 4 bonds. Alkanes

9. For simplicity, the smaller lobe is often omitted from diagrams Alkanes

10. Alkane - methane H C H H H sp 3 hybrid orbital of carbon 1s orbital of hydrogen C H H HH molecular orbitals of methane molecule

11. Overlap of a half-filled sp 3 hybrid orbital with a half-filled 1s orbital of a hydrogen atom forms a new molecular orbital The shared pair of electrons are now under the influence of both nuclei.molecular orbital The new molecular orbital lies along the axis joining the two nuclei and is known as a SIGMA (σ ) BOND. A sigma bond is formed by end-on overlap of atomic orbitals. This model is consistent with the fact that all the bond angles in alkanes are 109.5° (the tetrahedral bond angle). Alkanes

12. Alkane - ethane molecular orbitals of ethane molecule C C HH HH HH

13. Bonding in alkenes Alkenes  unsaturated hydrocarbons  molecules contain at least one carbon to carbon double bond. How is a double bond formed? The explanation must fit the observed facts.

14. Bonding in alkenes How is a C = C bond formed? The observed facts. The ethene molecule is flat and all bond angles are 120°. The C = C bond is shorter than the single bond. The C=C bond strength is intermediate between a single and a triple bond but not twice as strong as a single bond. There is restricted rotation around the C=C bond. There are two isomers of but-2-ene which differ only in the position of the methyl groups. In one, they are on the same side of the double bond, while in the other they are on opposite sides. (Stereoisomerism).

15. The two isomers of but-2-ene C C H H CH 3 C CHH Trans - isomerCis - isomer The formation of the C=C is explained by hybridisation. Bonding in alkenes

16. The formation of the C=C involves sp 2 hybridisation Bonding in alkenes Atom after hybridisation Atom uses the 2s orbital and only two of the 2p orbitals Atom before hybridisation 1s 2 2p 2 2s 2 1s 2 p 3 x sp 2 hybrid orbitals Unhybridised p orbital

17. sp 2 hybridisation Bonding in alkenes On mathematically mixing these, three identical hybrid orbitals are obtained. These lie in the same plane at 120° to each other and at 90° to the remaining p orbital This is called sp 2 hybridisation and the new orbitals are called sp 2 hybrid orbitals.

18. C Three sp 2 hybrid orbitals unhybridised 2p orbital

19. H H HH SIGMA (σ ) BONDS A sigma bond is formed by end-on overlap of atomic orbitals.

20. H H H H Unhybridised p orbitals are close enough to overlap above and below the plane of the sp 2 bonds. This produces a new molecular orbital between the 2 carbon atoms with lobes above and below the molecular plane. This bond is called a pi (π) bond This extra bond pulls the carbon atoms closer together – bond length shortens

21. Sideways overlap produces a weaker bond than the end-on overlap C=C isn’t 2 x as strong as C-C Resistance to rotation π overlap is more effective when the p orbitals are parallel Any other position and the overlap is reduced – the bond will eventually be broken

22. The Bonding Continuum The molecular orbital formed from overlapping atomic orbitals is symmetrical around a mid-point where the bonding electrons are most likely to be found. Pure Covalent Bonds – Non -polar Polar Covalent Bonds When there is a large difference between the electronegativities of the two elements involved in the bond, the bonding molecular orbital will be asymmetrical.

23. Colour in Molecules

27. Organic molecules contain bonding orbitals (s and p ) and non-bonding (lone pair) orbitals ( usually designated with the letter n ). In addition they also have anti-bonding orbitals (s* and p*).

28. Colour in Molecules

29. With increased possibilities for both bonding and antibonding orbitals in larger conjugated systems, the gap between the highest occupied (bonding) molecular orbital (HOMO) and the lowest unoccupied (antibonding) molecular orbital (LUMO) is decreasing. Buta-1,3-diene, again, only absorbs in the UV region of the electromagnetic spectrum but molecules with larger conjugated systems will absorb from the Visible spectrum and produce colours. For example, lycopene which is responsible for the red colour in tomatoes.

30. Colour in Molecules