Correct order of C-O bond length is: | vedclass

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(B) To determine the bond length,we first calculate the bond order for each species.$1$. For $CO$: The bond order is $3.0$.$2$. For $CO^{+}$: The bond

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$C_2O_{4}^{2-} > CO^{+} > CO$

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(B) To determine the bond length,we first calculate the bond order for each species.$1$. For $CO$: The bond order is $3.0$.$2$. For $CO^{+}$: The bond order is $2.5$ (as one electron is removed from the antibonding orbital of $CO$).$3$. For $C_2O_{4}^{2-}$ (oxalate ion): The $C-O$ bond has partial double bond character due to resonance,with an average bond order of $1.5$.Since bond length is inversely proportional to bond order $(Bond \ Length \propto \frac{1}{Bond \ Order})$,the order of bond length is: $C_2O_{4}^{2-} (1.5) > CO^{+} (2.5) > CO (3.0)$.Therefore,the correct order is $C_2O_{4}^{2-} > CO^{+} > CO$.

Column $I$	Column $II$
$(A)$ $B_2$	$(p)$ Paramagnetic
$(B)$ $N_2$	$(q)$ Undergoes oxidation
$(C)$ $O_2^{-}$	$(r)$ Undergoes reduction
$(D)$ $O_2$	$(s)$ Bond order $\geq 2$
	$(t)$ Mixing of '$s$' and '$p$' orbitals

Correct order of $C-O$ bond length is:

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