RADICALS

Species with an odd number of electrons also cannot satisfy the octet rule.  Here's how we determine their Lewis structures and formal charges.

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In the case of a radical species, if you don’t initially realize that the species is a radical, then you will find that it is not possible to draw a valid Lewis structure since the difference between the sum of formal charges and the actual charge will be an odd number. Consider the radical NO2. One double bond will give a formal charge of –1 on the O atom and give the structure of [NO2]–, whereas two double bonds will give a +1 charge on the N atom and give the structure of [NO2]+. The simplest solution for constructing the Lewis structure of a radical is to temporarily add one electron (NO2 becomes [NO2]– in this case), draw the Lewis structure, and then remove one electron from the most electropositive atom (N in this case). The atom with an electron removed will have an increase in formal charge by +1.

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We can use our chemical principles to see why NO2 and NO don’t readily dimerise: N2O4 has positive charges next to each other whereas N2O2 has a long N–N bond due to repulsion between two lone pairs and two double bonds. N2O3, formed by combination of NO and NO2, has only one lone pair and two double bonds and is a more stable molecule.