Chemistry - A Molecular Science

Chapter 6 Molecular Structure & Bonding

LUMO HOMO

p*^4 p^2 p

1

p*^3

Figure 6.27 MOs for C

H 6

viewed from the top 6

Hydrogen atoms have been omitted. Only the relative phases of the orbitals, not their relative contributions to the MOs, are shown.

so

π^2

is the highest occupied MO (HOMO) and

π^3

is the lowest unoccupied MO (LUMO).

Note that both pairs of electrons reside in bonding orbitals, consis

tent with the Lewis

structure that shows two double bonds. However,

the orbitals are delocalized over all four

carbon atoms, not localized between two atoms as shown in the Lewis structure.

As our last example, we examine the delocalized

system in benzene (Cπ

H 6

) shown in 6

Figure 6.27. Recall that the double bonds in benzene are sometimes represented by a circle (Figure 6.8b) due to resonance in the molecu

le. Although the construction of the MOs is

beyond the scope of this text, an examin

ation of them demonstrates the rules for

construction and provides a be

tter understanding of the bonding in this very important

molecule. There are six carbon atoms and six p orbitals, so there are six

MOs. The π

lowest energy MO has no nodal planes and is

a bonding orbital delocalized over all six

atoms. The highest energy orbital has a nodal plane between each pair of atoms, but, due to the symmetry of benzene, this requir

es only three nodal planes. Thus, the four

remaining MOs must contain either one or two nodal planes. In fact, two MOs have one nodal plane, and two MOs have two nodal planes. The

system has six electrons, so only π

the three bonding MOs are occupied, which gi

ves rise to the three double bonds in the

Lewis structure. The three MOs are delocali

zed over all six carbon atoms, so, consistent

with representing the double bonds with a circle, the

electron density is spread over the π

entire molecule with no localized double bonds.

6.6

CHAPTER SUMMARY AND OBJECTIVES According to the valence shell electron pa

ir repulsion (VSEPR) model, the valence

electron groups, or regions (lone pairs and

bonds) adopt positions around the atom so as σ

to minimize electron-electron repulsion. The geometries that minimize the interactions require bond angles of 109

o for four groups, 120

o for three groups, and 180

o for two

groups. Thus, the application of VSEPR theory to a Lewis structure provides an excellent way to determine the geometry of atoms around a central atom.

In valence bond theory, bonds are produced by the overlap of two orbitals on the

bound atoms. The bond that results is called a

bond if the internuclear axis is enveloped σ

by the bonding electron density or a

bond if the axis lies in a nodal plane of the bond. π

All bonds contain one and only one

bond, but multiple bonds also contain σ

bonds. The π

bond order of a bond is the sum of the

and σ

bonds that it contains. Overlap of simple π

atomic orbitals does not account for the bond

angles obtained from VSEPR, so the atomic

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