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Consider the molecular orbital configurations of O2, O2+, and O2−. Determine the bond order and explain the relationship between bond order and bond length.

1. O2+: (σ1s)2 (σ*1s)2 (σ2s)2 (σ*2s)2 (σ2pz)2 (π2px = π2py)2 (π*2px = π*2py)1, B.O.: 2.5
2. O2: (σ1s)2 (σ*1s)2 (σ2s)2 (σ*2s)2 (σ2pz)2 (π2px = π2py)2 (π*2px = π*2py)2, B.O.: 2
3. O2−: (σ1s)2 (σ*1s)2 (σ2s)2 (σ*2s)2 (σ2pz)2 (π2px = π2py)2 (π*2px = π*2py)3, B.O.: 1.5
4. Since the bond length decreases as the bond order increases, hence, O2+ has the least bond length.

Correct answer: Since the bond length decreases as the bond order increases, hence, O2+ has the least bond length.

Solution

Bond order is directly proportional to bond strength and inversely proportional to bond length. O2+ has the highest bond order (2.5), so it has the shortest bond length.

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