Multiple Bonds

Objectives

By the end of this section, you will be able to:

  • Describe multiple covalent bonding in terms of atomic orbital overlap
  • Relate the concept of resonance to π-bonding and electron delocalization

Hybridization and Molecular Geometry

The hybrid orbital model appears to account well for the geometry of molecules involving single covalent bonds. Is it also capable of describing molecules containing double and triple bonds? We have already discussed that multiple bonds consist of σ and π bonds. Next we can consider how we visualize these components and how they relate to hybrid orbitals.

The Lewis structure of ethene, C2H4, reveals that one other carbon atom and two hydrogen atoms surrounds each carbon atom.

The three bonding regions form a trigonal planar electron-pair geometry. Thus, we expect the σ bonds from each carbon atom to form using a set of sp2 hybrid orbitals. These hybrid result from hybridization of two of the 2p orbitals and the 2s orbital. These orbitals form the C–H single bonds and the σ bond in the C=C double bond.

The remaining 2p orbital on each carbon atom, not involved in hybridization, overlaps to create the π bond in the C=C double bond. This unhybridized p orbital (lobes shown in red and blue in Figure 3, below) is perpendicular to the plane of the sp2 hybrid orbitals. As a result, the unhybridized 2p orbitals overlap in a side-by-side fashion. They overlap above and below the internuclear axis (Figure 3 (b)) and form a π bond.

In ethene, each carbon atom is sp2 hybridized, and the sp2 orbitals and the p orbital are singly occupied. The hybrid orbitals overlap to form σ bonds, while the p orbitals on each carbon atom overlap to form a π bond.
In the ethene molecule, C2H4, there are (a) five σ bonds. One C–C σ bond results from overlap of sp2 hybrid orbitals on the carbon atom with one sp2 hybrid orbital on the other carbon atom. Four C–H bonds result from the overlap between the C atoms’ sp2 orbitals with s orbitals on the hydrogen atoms. (b) The π bond is formed by the side-by-side overlap of the two unhybridized p orbitals in the two carbon atoms. The two lobes of the π bond are above and below the plane of the σ system.

In an ethene molecule, the four hydrogen atoms and the two carbon atoms are all in the same plane. If the two planes of sp2 hybrid orbitals tilt relative to each other, the p orbitals will not align properly to overlap and form the π bond.

Differences Between σ and π Bonds

The planar configuration for the ethene molecule occurs because it is the most stable bonding arrangement. This is a significant difference between σ and π bonds. Rotation around single (σ) bonds occurs easily because the end-to-end orbital overlap does not depend on the relative orientation of the orbitals on each atom in the bond. In other words, rotating around the internuclear axis does not change the overlap of the σ bonding orbitals. This is because the bonding electron density is symmetric about the axis.

In contrast, rotation about the internuclear axis is much more difficult for multiple bonds. Such rotation would drastically alter the off-axis overlap of the π bonding orbitals, essentially breaking the π bond.

In molecules with sp hybrid orbitals, two unhybridized p orbitals remain on the atom (Figure 4). We find this situation in acetylene, H−C≡C−H, which is a linear molecule. The sp hybrid orbitals of the two carbon atoms overlap end to end to form a σ bond between the carbon atoms (Figure 5a). The remaining sp orbitals form σ bonds with hydrogen atoms. The two unhybridized p orbitals on each carbon align side by side to form two π bonds through overlapping. The two carbon atoms of acetylene bond together with one σ bond and two π bonds, creating a triple bond.

Diagram of the two linear sp hybrid orbitals of a carbon atom, which lie in a straight line, and the two unhybridized p orbitals at perpendicular angles.
(a) In the acetylene molecule, C2H2, there are two C–H σ bonds and a $C≡C$ triple bond involving one C–C σ bond and two C–C π bonds. The dashed lines, each connecting two lobes, indicate the side-by-side overlap of the four unhybridized p orbitals. (b) This shows the overall outline of the bonds in C2H2. The two lobes of each of the π bonds are positioned across from each other around the line of the C–C σ bond.

Resonance and Delocalization

Hybridization involves only σ bonds, lone pairs of electrons, and single unpaired electrons (radicals). Structures that account for these features describe the correct hybridization of the atoms. However, many structures also include resonance forms. Remember that resonance forms occur when various arrangements of π bonds are possible.

Since the arrangement of π bonds involves only the unhybridized orbitals, resonance does not influence the assignment of hybridization. For example, molecule benzene has two resonance forms (Figure 6). We can use either of these forms to determine that each of the carbon atoms bonds to three other atoms with no lone pairs. Therefore, the correct hybridization is sp2. The electrons in the unhybridized p orbitals form π bonds. Neither resonance structure completely describes the electrons in the π bonds. They are not located in one position or the other, but in reality are delocalized throughout the ring. Valence bond theory does not easily address delocalization. The molecular orbital theory better describes bonding in molecules with resonance forms.

Each carbon atom in benzene, C6H6, is sp2 hybridized, independently of which resonance form is considered. The electrons in the π bonds are not located in one set of p orbitals or the other, but rather delocalized throughout the molecule.

Assignment of Hybridization Involving Resonance
Some acid rain results from the reaction of sulfur dioxide with atmospheric water vapor. This reaction leads to the formation of sulfuric acid. Sulfur dioxide, SO2, is a major component of volcanic gases as well as a product of the combustion of sulfur-containing coal. What is the hybridization of the S atom in SO2?

Solution
The minor resonance structures of SO2 are

In each resonance structure, two bonds and one lone pair of electrons surround the sulfur atom. Therefore, the electron-pair geometry is trigonal planar, and the hybridization of the sulfur atom is sp2.

Check Your Learning
Another acid in acid rain is nitric acid, HNO3, which is produced by the reaction of nitrogen dioxide, NO2, with atmospheric water vapor. What is the hybridization of the nitrogen atom in NO2? (Note: the lone electron on nitrogen occupies a hybridized orbital just as a lone pair would.)

Answer:
sp2