Lewis Structures Step by Step: Drawing Molecules That Make Sense

What a Lewis Structure Actually Shows

A Lewis structure (also called a Lewis dot structure or electron dot structure) is a 2D diagram that shows how the valence electrons of a molecule are distributed: which atoms are bonded to which, how many bonds there are between each pair, and where any lone pairs sit. It is not a picture of the molecule's 3D shape — that comes later, from VSEPR theory — but it is the starting point for every shape prediction, bond polarity analysis, and reactivity argument you will make.

If you cannot draw a correct Lewis structure for a molecule, you cannot reason reliably about its chemistry. The skill is worth practicing until it is automatic.

The Valence Electron Count

Lewis structures only track valence electrons — the electrons in the outermost shell. For main group elements, the group number gives the valence count directly:

• Group 1 → 1 valence electron
• Group 2 → 2
• Group 13 → 3
• Group 14 → 4
• Group 15 → 5
• Group 16 → 6
• Group 17 → 7
• Group 18 → 8 (helium has 2)

Transition metals complicate things — this guide focuses on main group compounds, where Lewis structures behave predictably.

The Five-Step Method

A reliable procedure for drawing Lewis structures:

1. Count total valence electrons in the molecule or ion.
2. Choose a central atom and sketch a skeleton connecting the atoms.
3. Place single bonds between the central atom and each outer atom.
4. Distribute remaining electrons as lone pairs, filling outer atoms first to satisfy the octet rule.
5. Form multiple bonds if the central atom still lacks an octet.

Let's apply it.

Worked Example 1: Water (H₂O)

Step 1 — Count electrons. O has 6, each H has 1. Total = 6 + 2(1) = 8.

Step 2 — Skeleton. Oxygen is the central atom. Hydrogen can never be central (it forms only one bond). The skeleton is H–O–H.

Step 3 — Single bonds. Two O–H bonds use 2 × 2 = 4 electrons.

Step 4 — Distribute remaining. 8 − 4 = 4 electrons left. Hydrogen is already full (2 electrons per bond, which is hydrogen's maximum). Place all 4 remaining electrons on oxygen as two lone pairs.

Step 5 — Check octets. Oxygen has 2 bonding pairs + 2 lone pairs = 8 electrons total. Hydrogen has 2 electrons each, which is its complete shell. Structure is correct.

Worked Example 2: Carbon Dioxide (CO₂)

Step 1 — Count electrons. C has 4, each O has 6. Total = 4 + 2(6) = 16.

Step 2 — Skeleton. Carbon is central. Skeleton: O–C–O.

Step 3 — Single bonds. Two C–O single bonds use 4 electrons.

Step 4 — Distribute remaining. 16 − 4 = 12 electrons. Place 3 lone pairs on each oxygen (12 electrons total). Each oxygen now has an octet.

Step 5 — Check carbon. Carbon has only 2 bonding pairs = 4 electrons. Short by 4.

Fix with multiple bonds. Convert one lone pair on each oxygen into a bonding pair with carbon. Each O–C single bond becomes an O=C double bond. Final structure: O=C=O with 2 lone pairs on each oxygen.

Now carbon has 4 bonds = 8 electrons. Each oxygen has 2 bonds + 2 lone pairs = 8 electrons. Complete.

Worked Example 3: Ammonia (NH₃)

Step 1 — Count electrons. N has 5, each H has 1. Total = 5 + 3(1) = 8.

Step 2 — Skeleton. Nitrogen central, three H attached.

Step 3 — Single bonds. 3 × 2 = 6 electrons used.

Step 4 — Distribute. 8 − 6 = 2 electrons → one lone pair on nitrogen.

Step 5 — Check. N has 3 bonds + 1 lone pair = 8 electrons. Each H has 2. Done.

The nitrogen lone pair in ammonia is important — it is what makes NH₃ a base.

Handling Ions: Add or Subtract Electrons

For a polyatomic ion, adjust the valence electron count to account for the charge:

• Anion (negative charge): add electrons equal to the charge. SO₄²⁻ gets 2 extra.
• Cation (positive charge): subtract electrons equal to the charge. NH₄⁺ loses 1.

The final Lewis structure is drawn inside brackets with the charge written as a superscript outside: [structure]ⁿ⁻.

Resonance: When One Structure Is Not Enough

Some molecules cannot be represented by a single Lewis structure. A classic example is ozone (O₃). You can draw it with the double bond on either side — and neither drawing is the true structure. Experimentally, both oxygen-oxygen bonds in ozone are identical, intermediate in length between a single and double bond.

Resonance structures are a way to handle this. You draw each valid Lewis structure and connect them with double-headed arrows (↔). The real molecule is understood to be a resonance hybrid — a blend of the contributing structures, with the electron density delocalized across all the positions where lone pairs and multiple bonds move.

Important: resonance structures are not alternating in time. The molecule does not flip back and forth between them. The hybrid is a single, stable structure that the individual drawings together approximate.

Formal Charge: Picking the Best Structure

When multiple valid Lewis structures are possible, formal charge helps you identify the most likely one.

Formal charge formula:

Formal charge = (valence electrons in free atom) − (lone pair electrons) − ½(bonding electrons)

A structure with formal charges closest to zero is generally the most plausible. If negative formal charges are unavoidable, place them on the most electronegative atom.

Example — CO₂:

• Central C: 4 valence − 0 lone pair − ½(8 bonding) = 4 − 0 − 4 = 0
• Each O: 6 valence − 4 lone pair − ½(4 bonding) = 6 − 4 − 2 = 0

All formal charges are zero — a clean, favorable structure.

Exceptions to the Octet Rule

Three situations break the "eight electrons around each atom" rule:

1. Incomplete octets. Elements like boron and beryllium often form stable compounds with fewer than 8 valence electrons. BF₃ has only 6 electrons around boron and is still a normal molecule.

2. Odd-electron species (radicals). Molecules with an odd total valence electron count cannot satisfy the octet on every atom. Nitric oxide (NO) has 11 valence electrons total; one atom ends up with an unpaired electron. These molecules are called radicals and are often highly reactive.

3. Expanded octets. Elements in period 3 and below (phosphorus, sulfur, chlorine, and others) can accommodate more than 8 valence electrons by using d-orbitals. SF₆ has 12 electrons around the sulfur atom; PCl₅ has 10 around phosphorus. First-period elements (H, He) and second-period elements (C, N, O, F) cannot expand.

Recognizing when to use an exception matters: if you force an octet on every atom in SF₆, you will never get a correct structure.

Common Pitfalls

• Miscounting valence electrons. Double-check group numbers and ion charges.
• Forgetting that hydrogen only takes 2 electrons. Hydrogen is never central, never has lone pairs in normal molecules, and never takes part in double or triple bonds.
• Drawing too many bonds on second-period atoms. Carbon, nitrogen, oxygen, and fluorine cannot expand beyond 8 electrons.
• Skipping the formal charge check when multiple structures are possible.

From Lewis Structures to Real Chemistry

Once you have a correct Lewis structure, you can read an enormous amount from it: bond orders, lone pair positions, likely bond polarities, reactive sites, and (combined with VSEPR theory) the 3D shape. The structure itself is a hypothesis about how the electrons are arranged — and a good one, provided you followed the rules. Every subsequent step in understanding a molecule starts from the Lewis structure you draw.