The molecular orbital (MO) theory provides a comprehensive way to describe the bonding in homodiatomic molecules like O2, O2+, O2–, and O22-. Let’s analyze each of these species to determine the nature of their bonding.
O2 (Dioxygen)
In the case of O2, the molecular orbital diagram shows that there are 12 valence electrons. These electrons fill the molecular orbitals as follows: σ2s2, σ2s*2, σ2p2, π2p4, π2p*2. The presence of two unpaired electrons in the π2p* orbitals indicates that O2 is paramagnetic. The bond order is calculated as (Number of bonding electrons – Number of antibonding electrons)/2, which for O2 is (8 – 4)/2 = 2. This means O2 has a double bond.
O2+ (Dioxygen Cation)
For O2+, one electron is removed from the π2p* orbital. This results in 11 valence electrons. The electron configuration is σ2s2, σ2s*2, σ2p2, π2p4, π2p*1. The bond order is (8 – 3)/2 = 2.5, indicating a stronger bond than in O2. O2+ is also paramagnetic due to the presence of one unpaired electron.
O2– (Superoxide Ion)
O2– has an extra electron compared to O2, making it 13 valence electrons. The electron configuration is σ2s2, σ2s*2, σ2p2, π2p4, π2p*3. The bond order is (8 – 5)/2 = 1.5, indicating a weaker bond than in O2. O2– is paramagnetic with one unpaired electron.
O22- (Peroxide Ion)
O22- has two extra electrons compared to O2, resulting in 14 valence electrons. The electron configuration is σ2s2, σ2s*2, σ2p2, π2p4, π2p*4. The bond order is (8 – 6)/2 = 1, indicating a single bond. O22- is diamagnetic as all electrons are paired.
In summary, the bond order decreases as we move from O2+ to O22-, indicating weaker bonds. The magnetic properties also change, with O2 and O2+ being paramagnetic, while O22- is diamagnetic.