Previously, we saw the construction of a sigma bond in ethane from two sp³ hybridised orbitals.  However, although sp³ hybridisation is perhaps the most common form, it is certainly not the only form of hybridisation available to carbon.

Now let us take a look at sp² hybridisation:  As the name indicates, this is hybridisation involving one s orbital, and two p orbitals.

Here we can see more clearly the triangular displacement of the hybrid orbitals.

This arrangement of sp² orbitals allows us to explain some of the observed features of ethene.  It was known for a long time that in order for the carbon atoms in ethene to be tetravalent (to accept four electrons into its valence shell),  they must share a total of four electrons (thus we have 4 valence electrons plus 1 from each hydrogen plus two from the other carbon, which equals 8).  It is not easy to imagine this without the use of hybridised orbitals.  However, let us imagine two CH₂ fragments, with an sp² carbon in each:

Now, as we bring them closer together, it is obvious that there will be overlap of one of the sp² orbitals from each fragment.  This forms a sigma bond as before in ethane.  However, there will also be overlap of the unhybridised p orbitals above and below the plane of the molecule, to from a pi bond:

This combination of sigma and pi bond is known as a double bond.  It is important to note that the use of hybridised orbitals has automatically predicted that ethene is planar, which it is, and we might also imagine that ethene will not twist about its C-C axis, because to do so will destroy the pi bond, which will be energetically unfavourable.  This is indeed the case.