The Orbital Approach to addition to the Carbonyl Group - Taking Things a Little Further
The curly arrow representation of the mechanism is enough for first years and non-chemistry majors but a fuller picture of the addition can be obtained by considering the orbitals involved. The carbonyl group is made of a σ bond and a π bond. The latter is formed by the overlap of non-hybridized 2p orbitals. It is not directly between the carbon and oxygen atoms but instead above and below the σ bond. This means it is further from the atoms, and is weaker. This is why it breaks first.
The valence bond model of the carbonyl group. The lowest unoccupied molecular orbital (LUMO) is shown with a larger coefficient on the carbon. The highest occupied molecular orbital or HOMO is larger on the oxygen. This has been shown as the two 2p orbitals instead of the normal π bond to make the size difference more obvious (the π* antibonding orbital is drawn correctly).
The electronegative oxygen atom skews the π bond so that there is a larger orbital coefficient on the oxygen indicating a greater chance of finding electrons closer to the oxygen. The empty, antibonding or π* orbital is the opposite, it has a larger coefficient on carbon. In the diagram the orbitals have been represented in two ways. In the bottom left corner, I’ve tried to show the π bond shared between the atoms. The right-hand column uses the 2p orbitals to emphasize the size difference. These are the orbitals that overlap to make the π bond.
When the carbonyl group reacts with a nucleophile, a pair of electrons from the highest occupied molecular orbital or HOMO of the nucleophile overlap with the lowest unoccupied molecular orbital or LUMO of the carbonyl electrophile. As the LUMO has the larger coefficient on the carbon, the best molecule overlap occurs here. Filling the antibonding π* orbital breaks the π bond and the electrons flow onto the oxygen.
The addition of a nucleophile to a carbonyl group. Top row shows the standard line diagram. The middle row shows the interaction of the HOMO of the nucleophile with the LUMO or π* orbital of the carbonyl group. The bottom row shows the consequence of this interaction and the breaking of the π bond.
The flat, trigonal planar, sp2 hybridized carbon is now a tetrahedral sp3 hybridized atom (hence the name ‘tetrahedral intermediate’).
The nucleophile always approaches the carbonyl from the same angle and this is approximately 107°, which is similar to the where the new bond will be. This angle is known as the Bürgi-Dunitz angle. This angle of approach maximizes the overlap of the nucleophiles HOMO and the LUMO or π* orbital. The π* orbital is not perpendicular to the σ bond due to the node running through the centre. A second argument suggests this angle of attack reduces electrostatic repulsion between the electrons of the nucleophile and the electrons of the π bond.
The nucleophile will approach at an angle of 107° (the Bürgi-Dunitz angle).
This angle of attack has important consequences for reactivity (and stereoselectivity).