AP Chemistry
Hybridization
Hybridization:
1. The mixing of nonequivalent atomic orbitals to give rise to new hybrid orbitals with
different characteristics. Hybridization is necessary to explain the bonding in many
molecules, especially that of carbon.
2. The number of hybrid orbitals created must equal the number of atomic orbitals mixed.
3. The type of hybrid orbital created depends on the type of atomic orbitals mixed.
4. Thus, from (2) and (3) above. . .
o 2 sp hybrid orbitals arise from combinations of one s and one p orbital
o 3 sp2 hybrid orbitals arise from combinations of one s and two p orbitals
o 4 sp3 hybrid orbitals arise from combinations of one s and three p orbitals
o 5 sp3d hybrid orbitals arise from combinations of one s, three p, and one d orbital
o 6 sp3d2 hybrid orbitals arise from combinations of one s, three p, and two d orbitals.
5. The number of regions of electron density around an atom determines the type of
hybridization it undergoes; thus to determine hybridization:
a. Draw the Lewis structure of the molecule
b. Determine the number of regions of electron density around the atom (regions of
electron density = any type of bond or lone pair)
c. Match the # regions of electron density to the total number of orbitals needed
(starting with an s). This will give the hybridization
6. Strength of bond in hybridization depends on the amount and type of overlap, while the
overlap depends on the shape and direction of orbitals involved. Two types of overlap:
o Sigma () bond:
 Results from end to end (head to head) overlap of orbitals.
 In  bonds, the highest electron density overlap occurs along the bond axis.
o Pi () bond:
 Results from side to side, parallel overlap of orbitals.
 Electron density is located above and below bonding axis
 Usually occurs between unhybridized p orbitals.
7. Single, double, and triple bonds can now be represented as follows:
o Single bond = 1  bond
o Double bond = 1 , 1  bond
o Triple bond = 1 , 2  bonds