Lewis Structure Rules
Learn the systematic rules for drawing Lewis dot structures.
Count valence electrons, apply the octet rule, and handle exceptions.
The Formula
Number of bonds = bonding electrons / 2
Lone pair electrons = valence electrons - bonding electrons
Lewis structures (also called Lewis dot diagrams) are drawings that show the bonding between atoms in a molecule and any lone pairs of electrons. They were introduced by American chemist Gilbert N. Lewis in 1916.
The basic steps for drawing a Lewis structure are: Count the total number of valence electrons for all atoms. Determine the central atom (usually the least electronegative, never hydrogen). Place single bonds between the central atom and surrounding atoms, using 2 electrons per bond. Distribute remaining electrons as lone pairs to satisfy the octet rule (8 electrons around each atom, or 2 for hydrogen).
If atoms lack a full octet after placing all electrons, convert lone pairs on adjacent atoms into double or triple bonds. Some elements (B, Be, P, S, and others) can be exceptions to the octet rule.
Variables
| Term | Meaning |
|---|---|
| Valence electrons | Electrons in the outermost shell. Group 1 = 1, Group 14 = 4, Group 17 = 7, etc. |
| Octet rule | Most atoms need 8 electrons in their outer shell for stability (2 for hydrogen) |
| Bonding pair | A pair of electrons shared between two atoms (one bond = 2 electrons) |
| Lone pair | A pair of non-bonding electrons on a single atom |
| Formal charge | Valence electrons - lone pair electrons - ½(bonding electrons) |
Example 1
Draw the Lewis structure for water (H₂O).
Count valence electrons: O has 6, each H has 1. Total = 6 + 1 + 1 = 8
O is the central atom. Place single bonds to each H: O-H and O-H (uses 4 electrons)
Remaining electrons: 8 - 4 = 4. Place as 2 lone pairs on oxygen.
Check octets: O has 2 bonds (4e) + 2 lone pairs (4e) = 8. Each H has 2. All satisfied.
H-O-H with 2 lone pairs on the oxygen atom
Example 2
Draw the Lewis structure for carbon dioxide (CO₂).
Count valence electrons: C has 4, each O has 6. Total = 4 + 6 + 6 = 16
C is central. Place single bonds: O-C-O (uses 4 electrons). Remaining = 12.
Distribute 12 electrons as lone pairs: 3 pairs on each O (6e each). Check C: only 4 electrons — needs more.
Convert one lone pair from each O into a bond: O=C=O. Now C has 8e (two double bonds), each O has 8e (2 bond pairs + 2 lone pairs).
O=C=O with 2 lone pairs on each oxygen
When to Use It
Lewis structures are essential for understanding molecular geometry and chemical bonding.
- Predicting molecular shape using VSEPR theory
- Determining bond order (single, double, or triple bonds)
- Identifying polar and nonpolar molecules
- Understanding reactivity and acid-base behavior