Nitrogen dioxide, NO2, is a brown, paramagnetic gas. Because the molecule possesses an odd number of valence electrons, its Lewis depiction inevitably contains one unpaired electron.
(1) Construct two distinct Lewis structures for NO2 that differ only in the placement of the unpaired electron and of any double bond(s). Clearly indicate the location of the unpaired electron in each structure.
(2) For each of the two structures, calculate the formal charge on nitrogen and on each oxygen atom.
(3) Based on your calculations, identify which of the two resonance forms makes the greater contribution to the actual electronic structure of NO2. Justify your answer fully, referencing both formal charges and the location of the unpaired electron.
(4) Discuss how the coexistence of these resonance forms accounts for the experimentally observed intermediate N–O bond lengths in NO2.