Essentially, hybridisation is the mixing of standard atomic orbitals to form new orbitals – which can be used to describe bonding in molecules.
Most importantly we have sp3, sp2 and sp hybridisation.
sp3 Hybridisation in Methane (CH4):
The best way I can describe sp3 hybridisation is in Methane (also the most basic choice!). This is simplified for expression. Remember that Carbon has 6 electrons.
- In methane (CH4), 1 Carbon binds with 4 Hydrogens. The carbon atom itself has only 2 electrons available for bonding in the 2p subshell.
- In order for 4 hydrogens to bind there need to be 4 electrons available for bonding, which cannot be achieved at the moment. The pull of a hydrogen nucleus results in an electron being excited from the 2s subshell into the 2p subshell, where it is available for bonding.
- This excitation changes the forces on the valence (bonding) electrons as the nucleus now exerts a stronger effective core portential upon them. This and other factors leads to the creation of a new ‘hybridised orbital’, called sp3.
This leaves 4 valence electrons which will each overlap with the s orbital of a Hydrogen to form a σ (sigma) bond. These hydrogens space themselves as far apart as possible, leading to the tetrahedral structure of methane.
Each of the bonds in the image above are σ-bonds.
sp2 Hybridisation in Ethene (C2H4):
This is similar to sp3 hybridisation, except there are only 2 hydrogen nuclei pulling on the bonding electrons (which need an electron each) and the other 2 electrons are required for the π (pi) bond (double bond) between the two Carbons.
The electron configuration in carbon starts the same:
but the resulting spread is different:
Only 2 of the 2p orbitals are used in sp2 hybridisation; in contrast to the 3 used in sp3 hybridisation (you should be seeing where the numbers come from!).
This leaves us with 3 sp3-orbitals and 1 p-orbital to bond with. 2 of the sp3 orbitals are used for forming σ-bonds with the 2 hydrogens, while the remaining sp3 orbital binds with the other carbon to form a σ-bond and the p-orbital bonds with a p-orbital from the other carbon to form a π-bond.
Every double bond (regardless of what atoms it joins) consists of a π-bond and a σ-bond.
sp Hybridisation in Ethyne (C2H2):
This can occur on an atom with a triple bond such as the alkynes. Ethyne is the simplest.
In this case we only have 1 hydrogen attached to a carbon, and three bonds between each carbon. That’s 1 hybridised bond between the carbon and hydrogen with another hybridised bond between the carbons. The other two p-orbitals form two more bonds between the carbons.
This essentially means that the triple bond consists of 1 σ-bond and 2 π-bonds.
Essentially, the hybridisation of the carbon atom is based on the number of bonds to other carbons or identical atoms.
sp3 = single bond
sp2 = double bond
sp = triple bond