Valentiebondtheorie en gehybridiseerde orbitalen

Valence bond theory proposes that a chemical bond results through an overlap of partially filled atomic orbitals, including those other than the spherical s orbital.

When a single bond forms between two non-spherical orbitals, the two orbitals will have head-to-head overlap.

The type of covalent bond formed by head-to-head overlap of atomic orbitals is called a sigma bond.

s and p orbitals overlapping to form covalent bonds cannot yield the various molecular shapes in the VSEPR model. Valence bond theory helps to explain this molecular geometry through the hybridization, or mixing, of atomic orbitals.

Some atomic orbitals involved in bonding recombine to form new orbitals whose shapes are a hybrid of the originals. The initial number of atomic orbitals and the number of hybrid orbitals generated is always the same.

The hybrid orbitals have a different shape from their constituent atomic orbitals, with one lobe that is significantly larger than the other. Thus, the electron probability density is highly concentrated in a directional lobe, which leads to a more effective overlap with the orbitals of other atoms. For clarity, these orbitals are often shown without the minor lobes.

For instance, in the excited state of carbon, the one s and three p orbitals that contain the four unpaired electrons undergo hybridization, yielding four equivalent hybrid s-p-three orbitals that occupy the vertices of a regular tetrahedron.

The mixing of one s and two p orbitals in the excited state of carbon generates three equivalent s-p-two hybrid orbitals with trigonal planar geometry.

Similarly, the hybridization of one s and one p orbital can create two sp orbitals oriented at 180 degrees to each other.  

The concept of hybridization also provides an explanation for the formation of multiple bonds. The side-on overlap of two p orbitals gives rise to a pi bond.

However, a pi bond can only be formed in double and triple bonds when a sigma bond already exists between two atoms. Because the pi bond exists on opposite sides of the internuclear axis, pi bonds are unable to rotate around this axis.