Formal Charge

Consider an atom in a molecule, e.g., N in NH₃. If we assign half of the electrons shared with H atoms to N and add the two unshared electrons, we have 6/2 + 2 = 5 electrons assignable to N. Since N contributed 5 valence electrons, we say that the N formal charge is zero. Similarly, the other neutral molecules of Figure 8-1 have zero formal charge on each atom. With O3, however, the central O has 1 unshared pair and half of 3 bonding pairs, 2 + 6/2 = 5; since O contributed 6 electrons, the formal charge on the central O of O₃ is +1. In each resonance structure, the double-bonded terminal O has 2 unshared pairs and half of 2 bonding pairs, 4 + 4/2 = 6, so that this O atom is neutral. The singly bonded terminal O atom has 3 unshared pairs and half of a bonding pair, 6 + 2/2 = 7, so that this O atom has a formal charge of -1. On the average, the terminal O atoms each have formal charges of -1/2. In CO₃^-2, the C atom has half of 4 bonding pairs, and thus C is neutral. As in O3, the doubly bonded O atom is neutral and the singly bonded O atoms have formal charges of -1. On the average each O atom has a formal charge of -2/3. Note that in each case the total charge is the sum of the formal charges.