Every molecule has a skeleton made of shared electrons. Lewis structures are how chemists draw that skeleton — a map of where the valence electrons live. Here's the method, from the first dot to the final check.
A Lewis structure (also called a Lewis dot structure) is a diagram that shows how valence electrons are arranged in a molecule. Dots represent lone pair electrons that belong to a single atom, while lines represent bonding pairs shared between two atoms. It's the simplest way to see which atoms connect, how many bonds they form, and where unshared electrons sit.
Valence electrons
Only the outermost electrons matter. Count them using the group number for main-group elements.
Two dots = one line
A shared pair of electrons is drawn as a single line between two atoms. Two lines mean a double bond.
Unshared electrons
Pairs that stay on one atom are drawn as dots. They affect shape, polarity, and reactivity.
Step 1 — Count the Valence Electrons
THE FOUNDATIONBefore you draw anything, you need to know how many valence electrons you have to work with. For main-group elements, the group number tells you directly.
| Group | Valence electrons | Example elements |
|---|---|---|
| Group 1 | 1 | H, Li, Na |
| Group 2 | 2 | Be, Mg, Ca |
| Group 13 | 3 | B, Al |
| Group 14 | 4 | C, Si |
| Group 15 | 5 | N, P |
| Group 16 | 6 | O, S |
| Group 17 | 7 | F, Cl, Br, I |
| Group 18 | 8 | He, Ne, Ar (noble gases) |
For a molecule, add the valence electrons of every atom together. Then adjust for charge:
- Add one extra electron for each negative charge (anion).
- Subtract one electron for each positive charge (cation).
Hydrogen (group 1) = 1 electron × 2 atoms = 2. Oxygen (group 16) = 6 electrons. Total valence electrons = 2 + 6 = 8 electrons, or 4 pairs. That's your entire budget for the whole molecule.
Step 2 — Choose the Skeleton
WHO CONNECTS TO WHOMNow decide which atom goes in the center and which atoms surround it. The central atom is usually the least electronegative atom that is not hydrogen. Hydrogen is always terminal because it can form only one bond.
- Carbon is almost always central when present — it forms four bonds and loves to be in the middle.
- Oxygen is often terminal, but becomes central when bonded to hydrogen (like in H₂O).
- Hydrogen and halogens (F, Cl, Br, I) are usually terminal.
- In simple oxyanions like CO₃²⁻, the carbon is central with oxygens around it.
Connect the outer atoms to the central atom with single bonds first — two dots per bond, drawn as one line. Don't worry about double bonds yet.
Step 3 — Add Lone Pairs to Satisfy the Octet Rule
FILL THE OUTER SHELLSAtoms are most stable when they have eight electrons in their valence shell — the octet rule. Hydrogen is the exception: it only needs two electrons (a duet).
After placing single bonds, count how many electrons each atom currently has. Add lone pairs as dots around the atoms until each one reaches its octet (or duet for hydrogen).
Nitrogen (group 15) = 5 valence electrons. Hydrogen × 3 = 3 electrons. Total = 8 electrons, or 4 pairs. Nitrogen is central; three single bonds to H use 6 electrons. The remaining pair (2 electrons) becomes a lone pair on nitrogen. Nitrogen now has 8 electrons: 3 bonds + 1 lone pair. Each H has 2 electrons from its bond. Octet satisfied.
Step 4 — Form Multiple Bonds if Needed
DOUBLE AND TRIPLE BONDSSometimes single bonds and lone pairs don't get every atom to an octet — especially when the central atom is short on electrons. In those cases, you convert lone pairs from outer atoms into extra bonds.
- Take a lone pair from an outer atom and move it between that atom and the central atom.
- That creates a double bond (four shared electrons, drawn as two lines).
- If two pairs are moved, you get a triple bond (six shared electrons, three lines).
Carbon dioxide (CO₂) is the classic example: carbon is central, each oxygen is terminal. Two single bonds use 4 electrons; each oxygen then has two lone pairs. Carbon still only has 4 electrons — not an octet. Moving one lone pair from each oxygen into a second bond gives two double bonds: O=C=O. Now carbon has 8 electrons and each oxygen has 8. Octet rule satisfied.
Step 5 — Check Formal Charges
PICKING THE BEST STRUCTUREWhen more than one Lewis structure is possible, formal charge tells you which one is best. It's a bookkeeping number — not a real charge — but it reveals how evenly the electrons are shared.
The best Lewis structure usually has:
- Formal charges closest to zero on every atom.
- Any negative formal charge on the most electronegative atom.
- No separation of positive and negative charges if a neutral structure is possible.
Ozone has two possible resonance structures. The central oxygen has a formal charge of +1, one terminal oxygen has −1, and the other has 0. Neither structure is a complete picture — the real molecule is a hybrid, often drawn with a double-headed arrow between the two forms.
Common Molecules at a Glance
Here's a quick reference for some of the most frequently drawn Lewis structures.
| Molecule | Valence electrons | Bonds | Lone pairs | Notes |
|---|---|---|---|---|
| H₂O | 8 | 2 single (O–H) | 2 on O | Bent shape, polar |
| NH₃ | 8 | 3 single (N–H) | 1 on N | Trigonal pyramidal |
| CH₄ | 8 | 4 single (C–H) | 0 | Tetrahedral |
| CO₂ | 16 | 2 double (C=O) | 2 on each O | Linear, nonpolar |
| HCN | 10 | 1 single (H–C), 1 triple (C≡N) | 1 on N | Linear |
Try It Yourself
Use the visual below to see the electron count for a simple molecule. Adjust the number of valence electrons and watch how the dots and bonds fill in.
Each bond uses 2 electrons. The rest become lone pairs. This simple counter helps you keep track of your electron budget.
Common Mistakes to Avoid
- Forgetting to count all valence electrons — always add them up first, including charge adjustments.
- Putting hydrogen in the center — hydrogen can only form one bond, so it must be terminal.
- Giving hydrogen an octet — hydrogen only needs two electrons (a duet).
- Ignoring formal charges — a structure with smaller formal charges is usually preferred.
- Drawing all single bonds — double and triple bonds are often needed to satisfy the octet rule.
Lewis structures are a two-dimensional model. They don't show the actual three-dimensional shape of a molecule, and they can't explain paramagnetism or delocalized electrons in metals. For those, you need VSEPR theory, molecular orbital theory, or more advanced models.