Lewis structures represent the last layer of an atom or ion occupied by electrons :
If it has in its last layer the same number of electrons than in its Lewis structure, the atom is neutral, if it has n electrons more (less) , it becomes a negatively (positively) charged ion having n negative (positive) charges, by example $O^{2 -}$ has $6 + 2 = 8\;e^-$ in its last layer.
To each atom, we attribute:
- Two electrons for each of its own doublets.
- An electron for each covalent bond.
Comparing then with the Lewis structure of the neutral aton, its charge is deduced, for example:
- To each $H$ atom is assigned $1\;e^-$ (one of the covalent bond it forms), thus the formal charge is equal to $0$ whenever
- To the $C$ atom are assigned $4\;e^-$ (those that form 4 covalent bonds ), thus the formal charge is equal to $0$
- To the $O$ atom are assigned $7\; e^-$ (3 own doublets and one electron of the covalent bond ), so the formal charge is $-1$
- To the $N$ atom are assigned $4\; e^-$ (one for each of the covalent bonds), thus the formal charge is equal to $1$
Often for reasons of convenience, the own doublets are not represented, for example:
- Obviously the $O$ atom above has two doublets and forms two covalent bonds, therefore its formal charge is equal to $0$
- Obviously the $O$ atom below has three doublets and forms one covalent bond, and therefore its formal charge is $-1$
- The central $C$ atom forms four covalent bonds, so its formal charge is $0$
- The lower $C$ atom forms four covalent bonds (one shown, the other three with $H$ atoms), therefore its formal charge is equal to $0$