How to Draw Lewis Structures: Step-by-Step Guide with Examples

Lewis structures (also called Lewis dot structures or electron dot diagrams) are 2D representations of how atoms in a molecule share or hold onto valence electrons. They show bonding pairs (shared electrons between atoms) and lone pairs (unshared electrons on individual atoms). Lewis structures are essential for predicting molecular geometry (via VSEPR theory), understanding chemical reactivity, and determining whether a molecule is polar or nonpolar.

Add up the valence electrons for every atom in the molecule. For main group elements, the group number gives the valence electron count: Group 1 = 1, Group 2 = 2, Groups 13-18 = 3-8. For polyatomic ions, add electrons for negative charges and subtract for positive charges. Example: CO2 has 4 (from C) + 6 + 6 (from two O) = 16 valence electrons. NH4+ has 5 (from N) + 4(1) (from four H) − 1 (positive charge) = 8 valence electrons.

The central atom is usually the least electronegative element (excluding hydrogen, which is always terminal). Carbon, nitrogen, sulfur, and phosphorus are common central atoms. Hydrogen is NEVER a central atom because it can only form one bond. When in doubt, the atom that appears only once in the formula is usually the central atom. Example: In CO2, carbon is the central atom. In H2O, oxygen is the central atom.

Connect the central atom to each surrounding atom with a single bond (one line = 2 electrons). Subtract the electrons used in bonds from your total. Example: CO2 — draw C with a single bond to each O. That uses 4 electrons (2 bonds × 2 electrons each), leaving 12 electrons to distribute.

Place remaining electrons as lone pairs on the outer atoms first, giving each a full octet (8 electrons). Then place any leftover electrons on the central atom. Remember: hydrogen only needs 2 electrons (a duet), not 8. Example: For CO2 with 12 remaining electrons — place 6 on each oxygen (3 lone pairs each). This uses all 12 electrons. But check: does carbon have an octet? With only 2 single bonds, carbon has 4 electrons, not 8.

If the central atom doesn't have a full octet, convert lone pairs from adjacent atoms into bonding pairs (double or triple bonds). Move one lone pair from an outer atom to form a double bond, or two lone pairs for a triple bond. Example: In CO2, carbon needs 4 more electrons. Convert one lone pair from each oxygen into a bonding pair, creating O=C=O (two double bonds). Now carbon has 8 electrons (4 from double bonds), and each oxygen has 8 (2 from the double bond + 2 lone pairs).

Calculate formal charge for each atom: Formal Charge = Valence Electrons − Lone Pair Electrons − (1/2 × Bonding Electrons). The best Lewis structure has formal charges closest to zero. If multiple valid structures exist, choose the one with the smallest formal charges, and any negative formal charges should be on the more electronegative atom. Need to verify your Lewis structure? ChemistryIQ can analyze Lewis structures from a photo and provide instant feedback at three detail levels.

Not all molecules follow the octet rule. Incomplete octets: Boron and beryllium compounds often have fewer than 8 electrons (e.g., BF3 has 6 electrons on boron). Expanded octets: Elements in period 3 and beyond can hold more than 8 electrons because they have accessible d orbitals (e.g., SF6 has 12 electrons on sulfur, PCl5 has 10). Odd-electron species: Molecules with an odd number of total valence electrons (like NO with 11 electrons) cannot satisfy the octet rule for every atom — these are called radicals.