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QUESTION

# How do you calculate the formal charge of NO3-?

The of the nitrate anion is of course -1. In the Lewis representation at least 3 of the four participating atoms bear a formal charge.

Formal charge is by definition a formalism; it has no physical reality, but may nevertheless be useful for calculation.

We can write the Lewis representation of the nitrate anion as, (O=)N^+(-O^-)_2. The nitrogen centre is quaternized, and bears a positive charge. Why? Because it has a share in 4 bonding electrons: 2 from the the doubly bound oxygen, and 1 each from the 2 N-O bonds. So with the 2 inner core electrons, the nitrogen centre is associated with 6 electrons ONLY rather than the 7 it requires for electrical neutrality. The nitrogen centre thus bears a positive charge.

Compare this with ammonia, a neutral molecule. We write :NH_3. The nitrogen centre owns (or has a share in) 5 electrons (2 from the lone pair are entirely associated with the nitrogen; and 3 electrons from the N-H bonds). Ammonia is rightly depicted as a neutral atom with no formal charges.

So back to nitrate: nitrogen has dibs on only 6 electrons, and therefore bears a formal positive charge. The doubly bound oxygen has a share in or owns 8 electrons, and so is depicted as neutral. The singly bound oxygen atoms have 9 electrons associated with them, and so each bears a negative charge. The overall charge on the nitrate ion is of course -1, which this representation is designed to suggest. Can you treat the sulfate anion, SO_4^(2-) in the same way? Where does the formal charge lie?