5 Steps to a 5 AP Chemistry

entire molecule. These are called molecular orbitals (MOs). These molecular orbitals have

definite shapes and energies associated with them. When two atomic orbitals are added, two

molecular orbitals are formed, one bonding and one antibonding. The bonding MO is of

lower energy than the antibonding MO. In the molecular orbital model the atomic orbitals

are added together to form the molecular orbitals. Then the electrons are added to the

molecular orbitals, following the rules used previously when filling orbitals: lowest-energy

orbitals get filled first, maximum of two electrons per orbital, and half fill orbitals of equal

energy before pairing electrons (see Chapter 3). When s atomic orbitals are added, one

sigma bonding (σ) and one sigma antibonding (σ*) molecular orbital are formed. Figure 11.9

shows the molecular orbital diagram for H 2.

Note that the two electrons (one from each hydrogen) have both gone into the sigma

bonding MO. The bonding situation can be calculated in the molecular orbital theory by

calculating the MO bond order. The MO bond order is the number of electrons in bonding

MOs minus the number of electrons in antibonding MOs, divided by 2. For H 2 in Figure 11.9

the bond order would be (2 – 0)/2 =1. A stable bonding situation exists between two atoms

when the bond order is greater than zero. The larger the bond order, the stronger the bond.

When 2 sets of p orbitals combine, one sigma bonding and one sigma antibonding MO

are formed, along with two bonding pi MOs and two pi antibonding (π*) MOs. Figure 11.10

shows the MO diagram for O 2. For the sake of simplicity, the 1s orbitals of each oxygen and

the MOs for these elections are not shown, just the valence-electron orbitals.

The bond order for O 2 would be (10 – 6)/2 =2. (Don’t forget to count the bonding

and antibonding electrons at energy level 1.)

Resonance

Sometimes when writing the Lewis structure of a compound, more than one possible struc-

ture is generated for a given molecule. The nitrate ion, NO 3 −, is a good example. Three pos-

sible Lewis structures can be written for this polyatomic anion, differing in which oxygen is

double bonded to the nitrogen. None truly represents the actual structure of the nitrate ion;

that would require an average of all three Lewis structures. Resonancetheory is used to

describe this situation. Resonance occurs when more than one Lewis structure can be written

for a molecule. The individual structures are called resonance structures (or forms) and are

156  Step 4. Review the Knowledge You Need to Score High

Increasing Energy

Atomic

Orbital

of H

Atomic

Orbital

of H

1 s 1 s

Molecular

Orbital

of H 2

σ

σ∗

Figure 11.9 Molecular orbital diagram of H 2.

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