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 $\text{C=C}$ double bond.

The remaining 2p orbital on each carbon atom, not involved in hybridization, overlaps to create the π bond in the $\text{C=C}$ 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 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, $\text{H−C≡C−H,}$ 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.

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.