Lewis Structures, Formal Charge and Resonance

Bookkeeping for Electrons

A Lewis structure accounts for every valence electron as either a bonding pair or a lone pair. The procedure is mechanical:

1. Count all valence electrons, adding one per negative charge and subtracting one per positive charge.
2. Join the atoms with single bonds, putting the least electronegative atom (never hydrogen) at the centre.
3. Complete octets on the outer atoms with lone pairs.
4. Any electrons left go on the central atom.
5. If the centre is short of an octet, convert lone pairs from neighbours into double or triple bonds.

Formal charge then tests whether the drawing is sensible:

$$FC = V - N - \tfrac{1}{2}B$$

where $V$ is the free atom's valence electrons, $N$ the nonbonding electrons on it, and $B$ the bonding electrons around it. The best structure keeps formal charges as close to zero as possible and puts any negative formal charge on the most electronegative atom.

Some molecules cannot be drawn correctly at all with one structure. Ozone, $\mathrm{O_3}$, appears to have one single and one double bond, yet both bonds are measurably identical at 128 pm, between a single (148 pm) and a double (121 pm). The truth is a resonance hybrid: the real molecule is a single averaged structure, not a rapid flicker between drawings.

Three families break the octet rule. Incomplete octets occur in $\mathrm{BF_3}$ (boron with 6) and $\mathrm{BeCl_2}$. Odd-electron species such as NO cannot pair everything. Expanded octets appear from period 3 downward, where $\mathrm{PCl_5}$ and $\mathrm{SF_6}$ hold 10 and 12 electrons.

Common pitfall: reading resonance as oscillation. The molecule does not spend half its time in each drawing. There is one structure, and the separate drawings are a limitation of a notation that insists bonds be whole numbers.