Given the electronic configuration of carbon, at first glance it would seem that it has only two electrons available for bonding:
However, carbon has the ability to rearrange the shape and energy of its atomic orbitals to maximise the opportunities for bonding, both to itself and to other atoms. This rearranging is called hybridisation.
- carbon hybridises the 2s and 2p orbitals to form four sp3 hybrid orbitals which have the same energy as the original s and p orbitals had between them
- in methane each sp3 hybrid orbital interacts with a 1s orbital from a hydrogen atom to form a single sigma (𝜎) C-H bond
- in ethene the 2s, 2px and 2py orbitals on each carbon atom hybridise to give three sp2 hybrid orbitals
- one of these sp2 hybrid orbitals interacts with an sp2 hybrid orbital from another carbon atom to form a regular single (sigma) C-C bond. The other two sp2 hybrid orbitals interact with 1s orbitals from hydrogen atoms forming single (sigma) C-H bonds.
- the unhybridised 2pz atomic orbitals on each carbon atom merge to make a ∏-bond, containing a pair of electrons, with electron density extending above and below the plane of the sigma bonds / carbon-hydrogen skeleton.
- a double bond is therefore the combination of a sigma bond and a pi bond between the same two carbon atoms. We can break the pi bond in a reaction, using the electrons to make new bonds to other atoms, and the molecule stays intact because the sigma C-C bond is unaffected.
There is a deeper dive in to the valence theory of bonding and molecular orbital theory – click the links to read more 🙃.