3.45

[Cindy Nguyen 1L]

3.45 asks us to draw the resonance structures for ClNO2. I drew that N would have a double bond with an oxygen atom and a single bond with the other oxygen atom. With the oxygen with a single bond with Nitrogen, I drew chlorine bonded to it so that that oxygen would now have two bonds. This way the nitrogen doesnt need a formal charge of +1. How are we supposed to know that N would need to have a formal charge of +1 in this problem? Why would my answer be incorrect?

[KateCaldwell 1A]

It asked specifically for a resonance structure, so I think you have to show what the structure looks like with a charge on certain atoms. The delocalized electrons have to be show in two different Lewis structures. I'm not fully sure, but I hoped that helped.

[Sarah Brecher 1I]

I'm a little bit confused on how you drew your original Lewis structure but the way I drew it was you start with N in the middle since it is the least electronegative atom with it connected to Cl and the 2 O's. As you mentioned, you have a single bond connected to Cl with a full octet surrounded it. Then, you have a double bond on one of the O's and then a full octet. Because N cannot have an expanded octet, there's a single bond on the other O with a full octet rather than a double bond.

Therefore, the N has a formal charge of +1 and the single bond O has a formal charge of -1. The reason why you have these formal charges still is because this is the best Lewis structure that will minimize the formal charges. Then, you draw a resonance structure and move the double bond.