Every arrow in organic chemistry starts from a lone pair or bond, and every charge comes from a formal-charge count — so resonance, acids and bases, and mechanisms all follow from these two skills.
The species we will build and charge-count below — drawn live.
1. A Lewis Structure Shows Every Bonding Pair and Every Lone Pair
A Lewis structure draws a line for each bonding pair and a dot-pair for each lone pair; as bonds fall 4 → 3 → 2, lone pairs climb 0 → 1 → 2 to keep a full octet.
2. Build Any Lewis Structure in the Same Five Steps
(1) Count valence electrons (CO₂ = 4 + 2·6 = 16); (2) draw a single-bonded skeleton, least electronegative atom central; (3) fill outer octets with lone pairs; (4) make multiple bonds to satisfy a short central atom; (5) check formal charge.
3. Formal Charge = Valence Electrons − Lone-Pair Electrons − Half the Bonding Electrons
Take valence electrons minus lone-pair electrons minus half the bonding electrons: water's O = 6 − 4 − 2 = 0, hydronium's O = 6 − 2 − 3 = +1, ammonium's N = 5 − 0 − 4 = +1 — the extra bond, not a missing electron, carries the charge.
4. The Formal Charges Always Add Up to the Overall Charge
The formal charges on every atom must sum to the overall charge: nitromethane's +1 N and −1 O cancel to neutral, and methoxide's lone −1 O matches its −1 — if they don't sum, you miscounted at step 1.
5. Formal Charge Picks the Best Resonance Contributor
The major contributor has the fewest and smallest formal charges, with negatives on the most electronegative atom — though carbon monoxide's octet-satisfying triple bond forces a −1 on carbon, and ozone splits into two equivalent charged forms.
6. Where the Charge Sits: Reading Real Ions
In common ions the charge is delocalized by resonance: nitrate spreads −1 over three oxygens (N is +1), carbonate spreads −2 over three, and acetate spreads −1 over two — see resonance structures and hybridization.
7. Summary
Count valence electrons · single-bonded skeleton · fill outer octets · multiple bonds for a short center · formal charge = valence − lone-pair − ½(bonding) · use it to check (charges sum to total), locate, and judge the best structure.
Worked example
- Formal charge = (valence e⁻) − (nonbonding e⁻) − ½(bonding e⁻).
- Nitrogen: 5 valence electrons; in O=N=O it has 0 lone-pair electrons and 8 bonding electrons (two double bonds).
- FC = 5 − 0 − ½(8) = 5 − 4 = +1.
Answer. Nitrogen is +1 (each oxygen is 0) — consistent with the overall NO2+ charge.
Quiz yourself
Tap a question to reveal the answer — free, no login.
+1. Oxygen's group valence = 6. It keeps one lone pair (2 nonbonding electrons) and makes three bonds (6 bonding electrons, half = 3). Formal charge = 6 − 2 − 3 = +1. The extra bond, not a missing electron, is what makes it positive.
Zero. Formal charges always sum to the overall charge: (+1) + (−1) + zeros on every other atom = 0, so the molecule is neutral even though it contains charged atoms.
The one with the −1 on oxygen. When charges are otherwise equal, the major resonance contributor places the negative charge on the more electronegative atom, and oxygen is more electronegative than nitrogen.
Because giving both atoms a full octet requires a carbon–oxygen triple bond. With three bonds and one lone pair, carbon counts 4 − 2 − 3 = −1 and oxygen counts 6 − 2 − 3 = +1. Satisfying the octet outranks the slightly awkward charge placement.
Draw this on the whiteboard
Open the OChem Board whiteboard — benzene rings, wedge/dash bonds, and a clickable periodic table built in. No account needed.