(Sorry, I can not link to the picture, so just put URL here.)
Now we are learning Lewis Structure and Formal Charges. I encounter a problem when I was asked to draw the Lewis structure of N2O5. The problem tells me that there are no N-N or O-O bonds, and the structure should be symmetrical.
Formal Charge=valence electrons-nonbonding electrons-bongding electrons/2
According to the formal charge rule, a structure should be most stable without any formal charges. So I draw a Lewis Structure different from the one which is correct (See picture). I used N=O bonds to connect all the oxygens at four corners (sorry I can not show my work, since I can't upload here). And the oxygen at four corners have another 2 lone pairs of electrons. Calculate their formal charges, I got 6-4-4/2=0. And the middle part of the structure is the same as the picture shows, which is N-O-N. Nitrogen has 5 valence electrons, so if it has 5 single bonds and no lone electron, it has no formal charge. And the oxygen in the middle, which has two single bonds and two lone pairs of electrons, the formal charge is also 0. Every one is happy, and the electrons in total is 40, which corresponds to 6*5+5*2=40.
However, the correct structure has formal charges. Why we use this structure instead of the one I thought was right? I'm so confused.