Lewis Structures
Lewis structures are simplified diagrams that show how atoms in a molecule are connected and how their valence electrons are arranged. They represent the bonding between atoms and the lone pairs of electrons that may exist in the molecule.
This guide covers key definitions, a step-by-step drawing process, formal charge, resonance structures, exceptions to the octet rule, essential formulas, memory aids, common mistakes, and a 10-question practice quiz.
1What Are Lewis Structures and Why Do They Matter?
Imagine you have a LEGO set, but no instruction manual. How would you know how to build the model? In chemistry, Lewis structures act as our instruction manual, providing a simplified diagram that shows how atoms are connected and how their valence electrons are arranged.
They are fundamental to understanding many chemical properties, including molecular geometry, polarity, reactivity, and bonding (single, double, and triple bonds).
Think of Lewis structures like architectural blueprints for molecules. Just as a blueprint shows you how rooms connect in a building, a Lewis structure shows you how atoms bond and where electron "rooms" (lone pairs) are located.
Molecular Geometry
Lewis structures help predict the 3D shape of molecules using VSEPR theory.
Polarity & Reactivity
The electron distribution reveals whether a molecule is polar and how it will react.
2What Are the Key Terms You Need to Know?
Mastering these terms is essential for understanding Lewis structures. Refer back here as needed.
Lewis Symbol
An element symbol with dots representing its valence electrons (e.g., C with 4 dots)
Valence Electrons
Electrons in the outermost shell of an atom, involved in chemical bonding
Octet Rule
Atoms tend to gain, lose, or share electrons to have 8 in their outermost shell
Duet Rule
Hydrogen and helium are stable with only 2 valence electrons
Lone Pair
A pair of valence electrons not shared with another atom, belonging to one atom
Bonding Pair
A pair of valence electrons shared between two atoms in a covalent bond
Single Bond
One shared pair of electrons (2e⁻), represented by a single line (C–C)
Double Bond
Two shared pairs (4e⁻), represented by two parallel lines (C=C)
Triple Bond
Three shared pairs (6e⁻), represented by three parallel lines (C≡C)
Formal Charge
Hypothetical charge on an atom assuming equal sharing of bonded electrons
Resonance
Multiple valid Lewis structures needed when one is insufficient to describe bonding
Electronegativity
An atom's ability to attract shared electrons; key for choosing the central atom
3How Do You Draw a Lewis Structure Step by Step?
Follow these steps meticulously to draw accurate Lewis structures for most molecules and polyatomic ions.
Step 1: Count Total Valence Electrons
- Sum the valence electrons for all atoms in the molecule
- For polyatomic ions: add 1e⁻ per negative charge, subtract 1e⁻ per positive charge
- Example: CO₂ = C(4) + O(6) + O(6) = 16 valence electrons
Step 2: Identify the Central Atom
- Usually the least electronegative atom (except H, which is never central)
- Often the atom that appears only once in the formula
- Example: C is central in CO₂; N is central in NH₃
Step 3: Draw a Skeletal Structure
- Connect the central atom to terminal atoms with single bonds
- Each single bond uses 2 valence electrons
- Subtract used electrons from your total count
Step 4: Distribute Remaining Electrons
- Place lone pairs on terminal atoms first until each has an octet (except H)
- Then place any remaining electrons on the central atom
Step 5: Form Multiple Bonds (if needed)
- If the central atom lacks an octet, move lone pairs from terminal atoms to form double or triple bonds
- Continue until the central atom achieves an octet (or satisfies its exception)
Step 6: Check Your Work
- Verify all atoms have octets (or duets for H)
- Confirm total electrons match your Step 1 count
- Calculate formal charges to find the best structure
Lewis Structure Builder: CO₂
Watch a step-by-step construction of the Lewis structure for carbon dioxide.
C has 4 valence e⁻, each O has 6. Total: 4 + 6 + 6 = 16 valence electrons.
4How Does Formal Charge Help Choose the Best Structure?
When multiple Lewis structures seem valid, formal charge helps determine which is most plausible. It is the hypothetical charge on an atom assuming electrons in bonds are shared equally.
FC = (Valence e⁻ of free atom) − (Non-bonding e⁻) − (½ × Bonding e⁻)
Rules for the Best Structure
Rule 1: Sum of formal charges = 0 for neutral molecules, or = ion charge for ions.
Rule 2: Structures with smaller (closer to zero) formal charges are preferred.
Rule 3: Negative formal charges should be on the more electronegative atoms.
Rule 4: Adjacent atoms should not have formal charges of the same sign.
Worked Example: CO₂
| Atom | Valence e⁻ | Non-bonding e⁻ | ½ Bonding e⁻ | FC |
|---|---|---|---|---|
| C (center) | 4 | 0 | 4 | 0 |
| O (left) | 6 | 4 | 2 | 0 |
| O (right) | 6 | 4 | 2 | 0 |
5What Are Resonance Structures?
Sometimes a single Lewis structure cannot accurately represent the true bonding in a molecule or ion. When there are multiple valid ways to place double or triple bonds and lone pairs, we draw resonance structures.
Resonance structures are not different molecules, nor do they interconvert. The actual molecule is a resonance hybrid — an average of all contributing structures. Electrons are delocalized over several atoms.
Double-Headed Arrow (↔)
Separates resonance structures. Do not use equilibrium arrows (⇌).
Delocalization
Electrons spread across multiple atoms, stabilizing the molecule or ion.
Resonance Visualization: NO₃⁻
Observe how electrons are delocalized in the nitrate ion.
The double bond is on the top oxygen. Each single-bonded oxygen carries a −1 formal charge.
6When Does the Octet Rule Not Apply?
While the octet rule is a powerful guideline, some molecules and ions deviate from it. There are three main categories of exceptions.
Octet Exceptions Comparison
Understand the three types of exceptions to the octet rule.
Incomplete Octet: BF₃
Boron has only 6 electrons (3 single bonds). Groups 2 & 13 elements can be stable with fewer than 8 electrons.
Central Atom
Boron (B)
Valence e⁻ on B
6 (3 bonds)
Period
2
Example
BF₃, BeH₂
Incomplete Octets (<8 electrons)
- Elements from Groups 2 and 13 (Be, B) can be stable with fewer than 8 electrons
- BF₃: Boron has only 6 valence electrons (3 single bonds)
- BeH₂: Beryllium has only 4 valence electrons (2 single bonds)
Expanded Octets (>8 electrons)
- Period 3+ elements (P, S, Cl, Br, I, Xe) have available d-orbitals
- PCl₅: Phosphorus has 10 valence electrons
- SF₆: Sulfur has 12 valence electrons
- Often seen when bonded to highly electronegative atoms (F, O)
Odd-Electron Molecules (Free Radicals)
- Molecules with an odd total number of valence electrons
- Cannot satisfy the octet rule for all atoms; contain an unpaired electron
- NO (11 e⁻) and NO₂ (17 e⁻) are common examples
- Highly reactive due to the unpaired electron
7Key Formulas and Rules
| Principle | Formula / Rule |
|---|---|
| Total Valence Electrons | Sum of (Group # for each atom) ± (Charge of ion) |
| Formal Charge (FC) | FC = (Valence e⁻) − (Non-bonding e⁻) − (½ × Bonding e⁻) |
| Octet Rule | 8 valence electrons for stability (most atoms) |
| Duet Rule | 2 valence electrons for stability (H and He) |
What Do These Mean?
The formal charge formula lets you evaluate each atom in a Lewis structure. Subtract non-bonding electrons and half the bonding electrons from the free atom's valence count. The best structure minimizes these charges.
The octet rule (8 electrons) guides most atoms, while the duet rule (2 electrons) applies to hydrogen and helium. Together with valence electron counting, these are the core tools for drawing accurate Lewis structures.
8Memory Aids
Count, Central, Connect, Complete, Create (multiple bonds), Check.
Hydrogen = 1 bond. Oxygen = 2 bonds. Nitrogen = 3 bonds. Carbon = 4 bonds.
"Less electronegative in the middle." Hydrogen is a lone wolf — never central, only 2 electrons.
"Formal charge wants to be zero (or on the most electronegative atom)."
"Period 3+ means expanded octets are possible." Period 2 elements (C, N, O, F) never expand.
9Common Mistakes Students Make
"Incorrectly counting total valence electrons."
This is the most common mistake. Double-check group numbers and remember to add electrons for negative ions and subtract for positive ions.
"Placing hydrogen as a central atom."
Hydrogen can only form one bond, so it can never be in the middle of a molecule. It is always terminal.
"Forgetting to form multiple bonds when the central atom lacks an octet."
If the central atom does not have 8 electrons after distributing lone pairs, you must convert lone pairs on terminal atoms into double or triple bonds.
"Trying to expand octets for Period 2 elements."
Carbon, Nitrogen, Oxygen, and Fluorine can never have more than 8 valence electrons. They lack d-orbitals.
"Ignoring formal charges."
While not always required, calculating formal charges helps confirm the most plausible structure, especially when resonance or exceptions are involved.
"Not enclosing ions in brackets with the charge outside."
For polyatomic ions, always draw brackets around the entire Lewis structure and place the overall charge as a superscript outside.
Frequently Asked Questions
- What's the difference between a Lewis structure and a structural formula?
- A structural formula simply shows the connectivity of atoms (e.g., CH₃CH₂OH). A Lewis structure goes a step further by showing all valence electrons, including lone pairs and distinguishing between single, double, and triple bonds, giving a more complete picture of electron distribution.
- Why can't Period 2 elements expand their octet?
- Period 2 elements (like C, N, O, F) only have 2s and 2p orbitals available for bonding. These orbitals can hold a maximum of 8 electrons. They lack accessible d-orbitals (which start from n=3) that would allow them to accommodate more than eight valence electrons.
- How do I know which atom is the central atom?
- Generally, the central atom is the least electronegative atom (excluding hydrogen, which is always terminal). It is also often the unique atom in the formula (e.g., C in CO₂, N in NH₃). If there are multiple possibilities, try drawing structures and checking formal charges.
- Are resonance structures real? Does the molecule constantly switch between them?
- No, resonance structures are not real distinct forms that interconvert. They are a limitation of our simple Lewis structure model. The actual molecule exists as a single, hybrid structure that is an average of all contributing resonance forms. Electrons are delocalized, meaning they are spread out over multiple atoms simultaneously.
- Why is formal charge important?
- Formal charge helps us choose the most plausible Lewis structure when multiple options exist. Structures with formal charges closest to zero are generally preferred, and if charges are present, negative charges should be on more electronegative atoms. It provides a way to assess electron distribution and stability.
Practice Quiz
Test your understanding of Lewis structures — select the correct answer for each question.
1.Which of the following elements can expand its octet?
2.What is the total number of valence electrons in the nitrate ion (NO₃⁻)?
3.According to the duet rule, which element is stable with only two valence electrons?
4.Which of the following statements about formal charge is TRUE?
5.What does a double-headed arrow (↔) between two Lewis structures signify?
6.Which molecule is an example of an incomplete octet?
7.How many lone pairs are present on the central nitrogen atom in ammonia (NH₃)?
8.Which of these elements can *never* be a central atom in a Lewis structure?
9.If a central atom has 10 valence electrons in its Lewis structure, it is considered to have a(n):
10.When drawing a Lewis structure for an ion, what should you remember to do?
Final Study Advice
- 1. Practice the 6 C's process on at least 10 different molecules until it becomes automatic.
- 2. Always check formal charges after drawing your structure — it is the quickest way to catch errors.
- 3. Memorize HONC 1234 — it lets you quickly verify bond counts for the most common atoms.
- 4. Draw all resonance structures for polyatomic ions like NO₃⁻, CO₃²⁻, and SO₃²⁻ to build intuition.
- 5. Know the three octet exceptions cold: incomplete (B, Be), expanded (Period 3+), and odd-electron (radicals).