Ch 8: Hybridization Orbitals

Hybrid Orbitals
The VSEPR theory works well when accounting for molecular shapes, but it does not help much in describing the types of bonds formed. Orbital hybridization provides information about both molecular bonding and molecular shape. In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals.
Hybridization Involving Single Bonds
Recall that the carbon atom’s outer electron configuration is 2s2 2p2, but one of the 2s electrons is promoted to a 2p orbital to give one 2s electron and three 2p electrons, allowing it to bond to four hydrogen atoms in methane. You might suspect that one bond would be different from the other three. In fact, all the bonds are identical. This is explained by orbital hybridization.
The one 2s orbital and three 2p orbitals of a carbon atom mix to form four sp3 hybrid orbitals. These are at the tetrahedral angle of 109.5°. As you can see in Figure 8.19, the four sp3 orbitals of carbon overlap with the 1s orbitals of the four hydrogen atoms. The sp3 orbitals extend farther into space than either s or p orbitals, allowing a great deal of overlap with the hydrogen 1s orbitals. The eight available valence electrons fill the molecular orbitals to form four C —H sigma bonds. The extent of overlap results in unusually strong covalent bonds.

Figure 8.19

Hybridization Involving Double Bonds
Hybridization is also useful in describing double covalent bonds. Ethene is a relatively simple molecule that has one carbon–carbon double bond and four carbon–hydrogen single bonds.

Experimental evidence indicates that the H—C—H bond angles in ethene are about 120°. In ethene, sp2 hybrid orbitals form from the combination of one 2s and two 2p atomic orbitals of carbon. As you can see in Figure 8.20, each hybrid orbital is separated from the other two by 120°. Two sp2 hybrid orbitals of each carbon form sigma-bonding molecular orbitals with the four available hydrogen 1s orbitals. The third sp2 orbitals of each of the two carbons overlap to form a carbon–carbon sigma-bonding orbital. The nonhybridized 2p carbon orbitals overlap side-by-side to form a pi-bonding orbital. A total of twelve electrons fill six bonding molecular orbitals. Thus five sigma bonds and one pi bond hold the ethene molecule together. The sigma bonds and the pi bond are two-electron covalent bonds. Although they are drawn alike in structural formulas, pi bonds are weaker than sigma bonds. In chemical reactions that involve breaking one bond of a carbon–carbon double bond, the pi bond is more likely to break than the sigma bond.


Hybridization Involving Triple Bonds
A third type of covalent bond is a triple bond, such as is found in ethyne (C2H2), also called acetylene.

As with other molecules, the hybrid orbital description of ethyne is guided by an understanding of the properties of the molecule. Ethyne is a linear molecule. The best hybrid orbital description is obtained if a 2s atomic orbital of carbon mixes with only one of the three 2p atomic orbitals. The result is two sp hybrid orbitals for each carbon.
The carbon–carbon sigma-bonding molecular orbital of the ethyne molecule in Figure 8.21 forms from the overlap of one sp orbital from each carbon. The other sp orbital of each carbon overlaps with the 1s orbital of each hydrogen, also forming sigma-bonding molecular orbitals. The remaining pair of p atomic orbitals on each carbon overlap side-by-side. They form two pi-bonding molecular orbitals that surround the central carbons. The ten available electrons completely fill five bonding molecular orbitals. The bonding of ethyne consists of three sigma bonds and two pi bonds