Lewis Structures
A step-by-step guide to drawing dot structures, from simple molecules to resonance and beyond
What Lewis Structures Show
A Lewis structure is a 2D diagram that shows how valence electrons are arranged in a molecule. Lines represent shared electron pairs (bonds), and dots represent unshared electrons (lone pairs). Together they account for every valence electron in the molecule.
Lewis structures are the starting point for nearly everything in molecular chemistry. Once you have the correct Lewis structure, you can determine hybridization, predict 3D geometry with VSEPR, assess polarity, identify reactive sites, and draw resonance structures. Getting the Lewis structure wrong cascades errors through all of these.
Lewis structure of water (H₂O)
Each line = 2 shared electrons. Each pair of dots = 2 unshared electrons. Total: 8 valence electrons.
Reading the notation
Counting Valence Electrons
Every Lewis structure starts with the same question: how many valence electrons do I have to work with? Get this number wrong and nothing else will work. The count must be exact.
Valence electrons by group
For main group elements, the group number tells you the valence electron count:
Common elements you should memorize
| Element | Group | Valence e⁻ | Typical bonds |
|---|---|---|---|
| H | 1A | 1 | 1 bond, 0 lone pairs |
| C | 4A | 4 | 4 bonds, 0 lone pairs |
| N | 5A | 5 | 3 bonds, 1 lone pair |
| O | 6A | 6 | 2 bonds, 2 lone pairs |
| F, Cl, Br, I | 7A | 7 | 1 bond, 3 lone pairs |
| S | 6A | 6 | 2 bonds, 2 lone pairs (can expand) |
For neutral molecules
Add up the valence electrons of every atom. Example: CO₂ = 4 (from C) + 6 (from O) + 6 (from O) = 16 total.
For ions
Add electrons for negative charges, subtract for positive. NO₃⁻ = 5 + 6 + 6 + 6 + 1 (for the charge) = 24 total. NH₄⁺ = 5 + 1 + 1 + 1 + 1 - 1 (for the charge) = 8 total.
The Five-Step Drawing Method
This systematic approach works for the vast majority of molecules. Follow it exactly and you will get the correct structure.
Worked example: Carbon dioxide (CO₂)
C has 4, each O has 6. Total: 4 + 6 + 6 = 16 electrons (8 pairs).
Place the least electronegative atom in the center. Carbon is less electronegative than oxygen, so C goes in the middle. Connect with single bonds: O-C-O. This uses 4 electrons (2 per bond). 12 remain.
Each O needs 8 electrons total. Each already has 2 from the bond, so each needs 6 more as lone pairs (3 pairs each). That uses all 12 remaining electrons. 0 remain.
Carbon currently has only 4 electrons (2 from each single bond). It needs 8. Fix this by converting lone pairs from the outer atoms into bonding pairs. Move one lone pair from each O to form a double bond. Now C has 8 electrons: 4 from the left C=O + 4 from the right C=O.
C: 4 - 0 - (8/2) = 0. Each O: 6 - 4 - (4/2) = 0. All formal charges are zero. This is the best possible structure.
Final Lewis structure of CO₂
O=C=O with two double bonds. Each atom has a full octet. All formal charges are zero.
Lone Pairs and the Octet Rule
The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons. Hydrogen is the exception - it only needs two (a "duet"). The octet includes both bonding electrons and lone pair electrons.
Understanding how lone pairs fit into the octet is critical. An oxygen atom with 2 bonds and 2 lone pairs has 2 + 2 + 2 + 2 = 8 electrons around it. A nitrogen with 3 bonds and 1 lone pair also has 8. Each element has a predictable bonding pattern that satisfies its octet:
| Element | Valence e⁻ | Bonds needed | Lone pairs | Octet count |
|---|---|---|---|---|
| C | 4 | 4 | 0 | 4(2) = 8 |
| N | 5 | 3 | 1 | 3(2) + 1(2) = 8 |
| O | 6 | 2 | 2 | 2(2) + 2(2) = 8 |
| F, Cl | 7 | 1 | 3 | 1(2) + 3(2) = 8 |
| H | 1 | 1 | 0 | 1(2) = 2 (duet) |
Quick shortcut
For any main group element: bonds + lone pairs = 4 (for elements that follow the octet rule). Carbon makes 4 bonds and 0 lone pairs. Nitrogen makes 3 bonds and 1 lone pair. Oxygen makes 2 bonds and 2 lone pairs. This pattern holds in neutral molecules and is the fastest way to check your work.
When to Form Multiple Bonds
After placing all lone pairs (Step 3), if the central atom still doesn't have an octet, you need to convert lone pairs from neighboring atoms into bonding pairs. Each lone pair you convert creates an additional bond: single becomes double, double becomes triple.
You will know multiple bonds are needed when the central atom has fewer than 8 electrons after Step 3. This happens most often with carbon, nitrogen, and oxygen as central atoms. It never happens with hydrogen (which only needs 2) or halogens (which rarely serve as central atoms).
How to recognize the need for multiple bonds
N₂
Triple bond
N≡N with 1 lone pair on each N. 10 valence e⁻ total.
CH₂O
Double bond
H₂C=O with 2 lone pairs on O. 12 valence e⁻ total.
HCN
Triple bond
H-C≡N with 1 lone pair on N. 10 valence e⁻ total.
Formal Charge
Sometimes you can draw more than one valid Lewis structure for the same molecule. Formal charge helps you decide which structure best represents reality. It compares how many electrons an atom "owns" in the Lewis structure to how many it has as a free atom.
The formula
Formal charge = (valence electrons) - (lone pair electrons) - (1/2 bonding electrons)
An atom "owns" all of its lone pair electrons and half of its bonding electrons (since bonding electrons are shared). If this ownership differs from its normal valence count, the atom carries a formal charge.
Formal charge comparison: Carbon monoxide (CO)
Structure A: triple bond
C: 4 - 2 - 3 = -1
O: 6 - 2 - 3 = +1
Structure B: double bond
C: 4 - 2 - 2 = 0
O: 6 - 4 - 2 = 0
Structure B has zero formal charges, but carbon only has 6 electrons (no octet). Structure A is preferred because both atoms have complete octets. This is one case where satisfying the octet rule matters more than minimizing formal charges.
Rules for choosing the best structure
Resonance Structures
Sometimes one Lewis structure is not enough to describe a molecule. Resonance occurs when you can draw two or more valid Lewis structures that differ only in the position of electrons (never atoms). The real molecule is a blend of all the resonance structures, called the resonance hybrid.
You know resonance is possible when a molecule has a double bond adjacent to an atom with a lone pair, or when you can place a double bond in more than one position and get equally valid structures.
Resonance in ozone (O₃)
Both structures are equivalent. The real O₃ has two equal O-O bonds (bond order 1.5), not one single and one double.
Rules for drawing resonance structures
Evaluating resonance contributors
Not all resonance structures contribute equally. Structures with fewer formal charges, complete octets on every atom, and negative charges on electronegative atoms are more important contributors.
When all resonance structures are equivalent (like in O₃, CO₃²⁻, or NO₃⁻), each contributes equally, and the bond lengths are all identical. This is the most common type of resonance you will encounter in introductory chemistry.
Exceptions to the Octet Rule
The octet rule works for most molecules you will encounter, but there are three well-defined categories of exceptions. Recognizing which exception applies is straightforward once you know the patterns.
Incomplete octet (fewer than 8 electrons)
Boron and beryllium commonly form stable compounds with fewer than 8 electrons around the central atom. BF₃ has only 6 electrons around boron. This is not an error in your drawing - boron genuinely has an empty p orbital, which is why BF₃ is such a strong Lewis acid.
Elements: B (6 e⁻), Be (4 e⁻), sometimes Al
Expanded octet (more than 8 electrons)
Elements in period 3 and below can accommodate more than 8 electrons because they have accessible d orbitals. Phosphorus can hold 10 (as in PCl₅), sulfur can hold 12 (as in SF₆), and xenon can hold 10 or 12 (as in XeF₂ or XeF₄).
Key rule: only elements in period 3+ can expand their octet. Carbon, nitrogen, oxygen, and fluorine never exceed 8 electrons because they have no accessible d orbitals.
Elements: P, S, Cl, Br, I, Xe, Se
Odd-electron species (free radicals)
Molecules with an odd total number of valence electrons cannot give every atom a full octet because electrons pair up in twos. At least one atom will have an unpaired electron, making the molecule a free radical.
Examples: NO (11 e⁻), NO₂ (17 e⁻), ClO (13 e⁻). These species are typically reactive because of the unpaired electron.
How to handle expanded octets in practice
When drawing Lewis structures for molecules like SO₂, you may find that a structure with an expanded octet on sulfur (two S=O double bonds, formal charges all zero) competes with a structure obeying the octet (S-O single bonds with formal charges). For introductory courses, follow your textbook's convention - some prefer minimizing formal charges (expanded), others prefer obeying the octet rule. Both structures are used and the real electron distribution is somewhere between them.
Common Mistakes and How to Avoid Them
Miscounting valence electrons
The most common error. Double-check your count before drawing anything. Remember that charges on ions change the count: add for negatives, subtract for positives. A wrong count makes every subsequent step wrong.
Putting the wrong atom in the center
The central atom should be the least electronegative atom. Hydrogen is always terminal (it can only form one bond). Carbon is almost always central in organic molecules. A common error is putting oxygen in the center of CO₂ instead of carbon.
Exceeding the electron count
Every electron must be accounted for. If your structure uses more electrons than you calculated in Step 1, you have an error. Count all bonding pairs and lone pairs and verify the total matches your initial count.
Giving hydrogen more than 2 electrons
Hydrogen has only a 1s orbital and can hold a maximum of 2 electrons. It never has lone pairs and never forms more than one bond.
Expanding the octet for period 2 elements
Carbon, nitrogen, oxygen, and fluorine can never have more than 8 electrons. Only elements in period 3 and below (like S, P, Cl, Br, I, Xe) can expand their octets. If your structure shows 10 electrons around carbon, it is wrong.
Putting It Together
The Lewis structure is not the end goal - it is the foundation for everything that follows. Once you have a correct Lewis structure, you can determine:
Hybridization
Count the steric number (bonding groups + lone pairs) on each atom in your Lewis structure. 2 = sp, 3 = sp², 4 = sp³. The Lewis structure gives you this count directly.
VSEPR geometry
VSEPR uses the same electron group count from the Lewis structure to predict 3D shape. Lone pairs matter here - they occupy space even though they are invisible in the molecular shape. The Lewis structure tells you exactly how many lone pairs each atom has.
Polarity
Bond polarity comes from the electronegativity difference between bonded atoms, which you can read from the Lewis structure. Whether those bond dipoles cancel depends on the 3D geometry, which you predicted from the Lewis structure via VSEPR.
Reactivity
Lone pairs, formal charges, and multiple bonds in Lewis structures identify the sites where reactions happen. Nucleophiles attack from lone pairs. Electrophiles attack at electron-poor atoms (positive formal charge or incomplete octets). Pi bonds are reactive because they are weaker than sigma bonds.
Practice drawing Lewis structures
Work through 37 molecules step by step with guided walkthroughs, or test yourself with practice questions.