Today, I had a friendly debate with my friend on what should be the oxidation state of $\ce {Cl}$ in the chlorine monoxide radical. I took the view that it should be $\ce {+2}$ while he took the view that it should be $\ce {+1}$. My reason for taking this view is simply that conventional state rules tell us that $\ce {O}$ should have a $\ce {-2}$ oxidation state assigned to it, when it is not part of a peroxide or in the superoxide ion, and when its bonding partner is less electronegative than it. His reason, on the other hand, is that one of the Lewis structures that can be drawn of this radical, features a $\ce {Cl-O}$ single bond. Hence, considering this single bond to be an ionic bond, we would thus observe charge separation to produce $\ce {Cl^+}$ and $\ce {O^-}$.

I would think that his perspective of the issue is flawed. Firstly, it is because there is more than one way of representing the covalent bonding between the two atoms. In fact, a simplistic MO diagram would allow us to deduce a bond order of $\ce {+1.5}$. Furthermore, a Lewis structure with a $\ce {Cl=O}$ double bond can also be drawn, although with $\ce {Cl}$ carrying a total of nine electrons. Hence, I believe that the approach of assigning states based on one's interpretation of the covalent bonding between the two atoms is inaccurate in this case. By following the conventional oxidation states, there would thus be no ambiguity in the assigning of oxidation states in this case.

Which view do you think is correct and why?

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    $\begingroup$ I'd say it would be correct not to apply the very idea of oxidation states here. $\endgroup$ – Ivan Neretin Nov 8 '19 at 15:30
  • $\begingroup$ @IvanNeretin I agree with that the use of oxidation states to understand the bonding within this molecule would not be correct. However, there is some usefulness of this assignment when we consider the redox reactions of the this radical such as: $\ce {4ClO + 2H2O -> 3HClO2 + HCl}$. Clearly, if we take the oxidation state of $\ce {Cl}$ to be +2 in $\ce {ClO}$, then this reaction would just be a mere disproportionation of $\ce {ClO}$. We would have to consider reduction of $\ce {O}$ in $\ce {ClO}$ if we took the alternative view. $\endgroup$ – Tan Yong Boon Nov 8 '19 at 16:00
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    $\begingroup$ For the same reason, I'm reluctant to use the word "redox" when speaking about the reactions of this radical. Besides, the radical in question is not even a compound; it must have been obtained a microsecond or so ago from some real compounds, and it is their reaction in the first place. $\endgroup$ – Ivan Neretin Nov 8 '19 at 16:03
  • $\begingroup$ @IvanNeretin The context is the photolysis of $\ce {HOCl}$ into $\ce {Cl}$ and $\ce {OH}$ and then, $\ce {Cl}$ goes on to abstract $\ce {H}$ from another $\ce {HOCl}$ to produce $\ce {OCl}$ radical. $\endgroup$ – Tan Yong Boon Nov 8 '19 at 16:06
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    $\begingroup$ A similar problem is assigning an oxidation state to iron in hemoglobin, when it is bound to oxygen. Depending on different resonance structures, you could take it to be +III, +IV or maybe even +V. Therefore, it is usually not assigned - and everything still works fine. $\endgroup$ – FusRoDah Nov 8 '19 at 16:18

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