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Question
Using molecular orbital theory, compare the bond energy and magnetic character of \[\ce{O^{+}2}\] and \[\ce{O^{-}2}\] species.
Solution
The electronic configurations of \[\ce{O^{+}2}\] and \[\ce{O^{-}2}\] to molecular orbital theory is:
\[\ce{O^{+}2}\]: σ1s2, σ∗1s2, σ2s2, σ∗2s2, σ2pz2, π2py2, π2px2, π∗2px1
\[\ce{O^{-}2}\]: σ1s2, σ∗1s2, σ2s2, σ∗2s2, σ2pz2, π2py2, π2px2, π∗2px2, π∗2py1
The bond order of \[\ce{O^{+}2}\]:
BO = `1/2 [N_b - N_a]`
BO = `1/2[10 - 5] = 5/2`
BO = 2.5
The bond order of \[\ce{O^{-}2}\]:
BO = `1/2[N_b - N_a]`
BO = `1/2[10 - 7] = 3/2`
BO = 1.5
As the bond order of \[\ce{O^{+}2}\] is higher, it is more stable than \[\ce{O^{-}2}\], because higher the bond order more stable is the bond. Both the molecular species have the presence of unpaired electrons. So, they both are paramagnetic in nature.
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| (ii) \[\ce{CO}\] | (b) 2.0 |
| (iii) \[\ce{O^{-}2}\] | (c) 2.5 |
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