How to Draw Lewis Dot Structures: A Step-by-Step Guide

In 1916 the American chemist Gilbert N. Lewis proposed a simple idea: atoms bond by sharing pairs of electrons, and they tend to end up with the same number of outer electrons as a noble gas. He drew molecules with dots for electrons around element symbols. More than a century later, those sketches — Lewis structures — are still the first tool chemists reach for when they want to understand a molecule.

What the dots mean

A Lewis structure shows only the valence electrons — the outer electrons involved in bonding.

• A pair of electrons shared between two atoms is a covalent bond, drawn as a line.
• A pair of electrons belonging to just one atom is a lone pair, drawn as two dots.

Most atoms in a stable molecule aim for eight electrons around them (the octet rule). Hydrogen aims for two.

The step-by-step method

1. Count the total valence electrons. Add up the valence electrons of every atom (use the group number: group 1 → 1, group 14 → 4, group 16 → 6, group 17 → 7). For a negative ion, add one electron per charge; for a positive ion, subtract. How to find valence electrons covers this in detail.
2. Draw the skeleton. Put the least electronegative atom in the centre (never hydrogen, which only forms one bond). Connect each outer atom to it with a single bond.
3. Subtract two electrons for each bond you drew.
4. Complete the outer atoms’ octets with lone pairs.
5. Put any leftover electrons on the central atom.
6. If the central atom still lacks an octet, move a lone pair from an outer atom to make a double (or triple) bond.
7. Check the total electron count and each atom’s octet.

Example 1: water, H₂O

1. Electrons: O has 6, each H has 1 → 8.
2. Skeleton: H–O–H (oxygen in the middle).
3. Two bonds use 4 → 4 left.
4. Hydrogens are complete with one bond each.
5. The 4 remaining electrons go on oxygen as two lone pairs.

Oxygen now has 2 bonds + 2 lone pairs = 8 electrons. Those lone pairs are why water is bent rather than straight — see VSEPR theory.

Example 2: carbon dioxide, CO₂

1. Electrons: C has 4, each O has 6 → 16.
2. Skeleton: O–C–O (carbon is less electronegative).
3. Two bonds use 4 → 12 left.
4. Give each oxygen three lone pairs: 12 used → 0 left.
5. Nothing left for carbon, which has only 4 electrons (two bonds).
6. Move one lone pair from each oxygen into a bond: O=C=O.

Now every atom has an octet: each oxygen has two bonds (4 electrons) plus two lone pairs, and carbon has two double bonds (8 electrons).

Example 3: nitrogen, N₂

10 valence electrons. A single N–N bond with lone pairs can’t give both atoms octets; you have to form a triple bond, N≡N, with one lone pair on each nitrogen. That triple bond is one of the strongest bonds in chemistry, which is why nitrogen gas is so unreactive.

Formal charge

Sometimes you can draw more than one valid structure. Formal charge helps choose between them:

formal charge = valence electrons − (lone-pair electrons + ½ × bonding electrons)

The best structure usually has formal charges as close to zero as possible, and any negative formal charge on the most electronegative atom.

Resonance: the nitrate ion

The nitrate ion, NO₃⁻, has 5 + 3(6) + 1 = 24 electrons. You end up with one N=O double bond and two N–O single bonds. But which oxygen gets the double bond? Any of the three.

These are resonance structures. The real ion isn’t flipping between them; it’s a blend of all three. Experiments confirm it: all three N–O bonds are identical in length, somewhere between a single and a double bond. The same thing happens in benzene, carbonate and ozone.

Exceptions to the octet rule

• Fewer than eight. Boron in BF₃ has only six electrons around it. This “electron-deficient” boron is why BF₃ readily accepts a lone pair, acting as a Lewis acid.
• More than eight. Elements in period 3 and beyond can have “expanded octets”: sulfur in SF₆ has six bonds, phosphorus in PCl₅ has five.
• Odd electron counts. Nitrogen monoxide, NO, has 11 valence electrons, so one must be unpaired. Molecules like this are called radicals and tend to be reactive.

Quick answers

Which atom goes in the middle? Usually the least electronegative one (excluding hydrogen), and often the one that appears only once in the formula.

Do Lewis structures show shape? No — they show connections and electron pairs. Shape comes from VSEPR theory applied to the Lewis structure.

Why eight electrons? Eight valence electrons fill the s and p orbitals of a shell, matching the stable configuration of a noble gas. See electron configuration rules.

Related

To check how many valence electrons an atom has, use its electron configuration — the outermost s and p electrons are the valence electrons for main-group elements.