Chemistry - A Molecular Science

6.4

VALENCE BOND THEORY AND HYBRIDIZATION

Figure 6.10 Overlap of atomic orbitals When atoms get close enough to bond,

their atomic orbitals overlap.

The overlap region is shown in yellow. NodalPlane

(a) p orbitals do not interact at

large separations.

(b) They are weakly interacting at

smaller separations.

(c) They are strongly interacting at

bonding separations. There isno electron density on theinternuclear axis, so theirinteraction is a bond.

p

Figure 6.11

π bond formation

The internuclear axis is shown as the green line in Figure c. There is electron density above and below it, but none resides on it.

+ + +

H+H

H^2

F+F

F^2

H+ F

HF

(a)

(b)

(c)

In

valence bond theory

, each bond results from the overlap of two atomic orbitals on two

adjacent atoms. The bonding electrons in such bonds are localized in the region between the two atoms. A single bond is composed of two bonding electrons, so the total number of electrons in the two overlapping atomic orbitals used to produce a bond cannot exceed two. In most cases, each bonding orbital contains one electron and the two electrons pair when the orbitals overlap. However, both electrons

can reside in one of the atomic orbitals

(a lone pair), but in this case, the other orb

ital must be empty. A bond in which a lone pair

is converted into a covalent bond is called a

coordinate covalent bond

. Coordinate

covalent bonds are produced in Lewis acid-b

ase reactions, which are discussed in Chapter

12. We limit our discussion here to cases where each overlapping orbital has one electron. SIMPLE OVERLAP MODEL The H-H bond discussed in Section 5.1 is produ

ced when the distance between the two H

atoms is so small that their 1s orbitals overl

ap (the overlap region is highlighted in yellow

in Figure 6.10a). The valence electron configuration of a fluorine atom is 2s

2 2p

5 , so an

unpaired electron resides in one of the 2p orbitals, and the F-F bond results from the overlap of two p orbitals as shown in Fi

gure 6.10b. In these two examples, the bonding

atoms were the same, so the overlapping orbita

ls were the same type (both s or both p

orbitals). However, the overlapping orbitals do not have to be the same type. The H-F bond is the result of overlap between the 1s or

bital of H and the 2p orbital of F (Figure

6.10c). The lone pairs on fluorine would then reside in its s and remaining p orbitals.

The line between the two atoms in a bond is called the

internuclear axis

, and bonds

in which the bonding electron density falls

on the internuclear axis are called

sigma

(σ

)

bonds

. The bonds in Figure 6.10 are

σ bonds. The

σ bond in F

(Figure 6.10b) results 2

from the end-on overlap of two p orbitals, but

p orbitals can also overlap in a side-on

fashion as shown in Figure 6.11. In Figure 6.11a, the two orbitals are not interacting. In Figure 6.11b, the p orbitals begin to overlap

(yellow region) and to distort into the bonding

orbital shown in Figure 6.11c. Note that the

overlap lies above and below the internuclear

axis, but not on it, so the bond is not a

σ bond. Bonds that place bonding electron density

above and below, but not on, the internuclear axis are called

pi

(

π)

bonds

. The

internuclear axis lies in the nodal plane of a

π bond. We conclude that

σ^

bonds result from

end-on overlap and place electron dens

ity on the internuclear axis, while

π

bonds are

produced from side-on overlap

and contain a nodal plane th

rough the internuclear axis

.

Chapter 6 Molecular Structure & Bonding

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