Question

Using molecular orbital theory, compare the bond energy and magnetic character of \[\ce{O^{+}2}\] and \[\ce{O^{-}2}\] species.

Answer 1

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, σ2p_z², π2p_y², π2p_x², π^∗2p_x¹

\[\ce{O^{-}2}\]: σ1s², σ^∗1s², σ2s², σ^∗2s², σ2p_z², π2p_y², π2p_x², π^∗2p_x², π^∗2p_y¹

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.