Hybridisation involving s, p and d orbitals only
Hybridisation means making something new from an amalgamation or combination of other parts. A hybrid plant is one made from two different plants blended together. The hybrid shows the characteristics of both plants.
In terms of chemistry we refer to the hybridisation of atomic orbitals to explain the change that seems to happen between the atomic orbitals in an uncombined atom and the orbitals used by the same atom when bonding.
The 's' orbital is spherical about the nucleus and the 'p' orbitals are like double headed balloons arranged along the axis of (imaginary) three dimensional coordinates.
However, it is apparent that the shapes of these orbitals are inadequate to explain the orientation of the bonds produced in molecules. The 'p' orbitals are oriented at 90º to one another and yet there are few molecules that show a bond angle of 90º (in fact the bond angle 90º does appear in some of the larger moolecules but that is due to different reasons).
The classic molecule to consider is methane CH4. In this molecule the bond angles indicate that the shape of the molecule is a perfect tetrahedron with bond angles of 109º 28' (approximately 109.5º)
It seems that the orbitals used for bonding are arranged as far apart as possible suggesting that they have the same energy (degenerate). We know that the orbitals on the carbon atom do not have the same energy - the 2s orbital is of lower energy that the three 2p orbitals. Hybridisation is a model that allows us to combine the atomic orbitals and then produce four degenerate orbitals to be used for bonding.
In order for the electrons to be ready for this process one of them must be promoted from the 2s orbital to the 2pz orbital.
It should be emphasised at this point that this is the norm rather than the exception. It seems that all elements undergo this hybridisation process (or a similar one) when bonding. Logically, the 2s orbital is in no position to overlap directly with another orbital from another atom without interfering with the p orbitals. This hybridisation process allows the 2s electrons to be involved in bonding.
If we study the shape of the water molecule we find that the four electron pairs around the oxygen are tetrahedrally arranged. They have hybridised and the sp3 orbitals so formed overlap directly with the 1s orbitals of the two hydrogens. The two lone pairs on the oxygen remain in the sp3 orbitals that are not used in bonding.
Other types of hybridisation
Carbon can also bond to three other atoms instead of four (as in methane) and it seems that it hybridised its orbitals using only the 2s and two of the 2p orbitals to do this.
There is the formation of a double bond in molecules such as ethene. The three sp2 hybrid orbitals are degenerate (same energy) and consequently arrange as far apart as possible in space i.e. at 120º to each other. This creates a trigonal shape that is planar leaving the remaining 2pz orbital to orientate itself above and below the plane of the other orbitals. This 2p orbital can then laterally overlap with adjacent singly occupied 'p' orbitals on adjacent atoms.
In sp hybridisation, carbon bonds to two other atoms by hybridising the 2s and only one of the 2p orbitals to produce two sp orbitals arranged at 180º to one another. The remaining two 2p prbitals can overlap with suitable orbitals on adjacent atoms to produce pi systems. Examples include ethyne, the nitrogen molecule, hydrogen cyanide, and any other triple bond systems.
Other forms of hybridisation
It should be mentioned that this hybridisation process can be extended to allow atoms to bond with more than four other atoms (octet expansion). In this case the hybridisation invariably involves one or more of the 'd' orbitals. Sulphur hexafluoride forms six attacments to the six fluorines and consequently needs six available orbitals. It gets these by promoting one electron from the 3s and 3px orbitals into two of the 3d orbitals. It can then hybridise the 3s, 3px, 3py, 3pz, and 3dxy, 3dxz orbitals into an octagonal arrangement each with one electron.
Atoms rearrange their atomic orbitals when bonding to produce orbitals with shapes more suitable for the bonding process. This is called hybridisation.
It is performed by almost all atoms when bonding although carbon provided the easiest examples to show.
It is easy to recognise the hybridisation used by simply observing the double or triple bonds.
Only single bonds = sp3 hybridisation
1 double bond = sp2 hybridisation
1 triple bond =sp hybridisation