Module 4 - 4.2: Drawing Lewis Structures for Molecular Compounds

Atomic symbol surrounded by dots showing number of valence electrons

Lewis Symbol

Diagram showing bonding pairs (lines) and lone pairs (dots) between atoms

Lewis Structure

Shared pair of electrons that forms a covalent bond (shown as a line)

Bonding Pair

Pair of non-bonding electrons shown as two dots on an atom

Lone Pair

To show how atoms bond and how valence electrons are distributed

Purpose of Lewis Structures

They connect directly to VSEPR theory and predict molecular shape

Why Draw Lewis Structures?

1️⃣ Determine total valence e⁻ using group numbers

Steps to Draw a Lewis Structure

2️⃣ Connect atoms with single bonds

3️⃣ Fill terminal atoms’ octets

4️⃣ Check and adjust central atom’s octet

H never has more than 2 e⁻

Hydrogen Rule

Atoms like C

Second Row Rule

Elements in the 3rd period or beyond can have more than 8 e⁻

Expanded Octet Rule

FC = (valence e⁻) − (lone-pair e⁻ + ½ bonding e⁻)

Formal Charge Formula

Helps identify most stable and reasonable structure

Formal Charge Purpose

Total FC must equal molecule or ion charge

Formal Charge Guidelines

place negative FC on more electronegative atoms

avoid like charges on neighbors

Has smallest possible formal charges and correct octets

Most Stable Lewis Structure

Two or more valid Lewis structures for the same molecule differing only in electron placement

Resonance Structures

Structures connected by a double-headed arrow (↔)

Representation of Resonance

Actual structure is a hybrid (average) of all resonance forms

Resonance Meaning

Has three equivalent resonance structures where negative charge is delocalized over all oxygens

Example – Chlorate Ion (ClO₃⁻)

Electrons shared between 3 or more atoms (not confined to one bond or atom)

Delocalized Electrons

Common bonding patterns that satisfy octet rule and have no formal charge

Ideal Neutral Fragments

Speeds up building Lewis structures by recognizing recurring stable arrangements

Purpose of Ideal Fragments

“Decorate” the central atom by attaching ideal fragments (like H

Structural Building Blocks – Step 1

Calculate FC on central atom to see if valence e⁻ are missing or extra

Step 2 – Check Formal Charge

Redistribute bonds or lone pairs to satisfy octet and minimize charges

Step 3 – Final Adjustments

• H → 1 bond
• B → 3 bonds (incomplete octet)
• C → 4 bonds
• N → 3 bonds + 1 lone pair
• O → 2 bonds + 2 lone pairs
• F → 1 bond + 3 lone pairs
These patterns help predict molecular frameworks

Common Bonding Patterns

Atoms or ions with the same total number of valence electrons

Isoelectronic Species

B (3 bonds) ↔ C⁺ (3 bonds) ↔ N²⁺ (3 bonds) — all behave similarly

Example of Isoelectronicity

N⁺ acts like neutral C

Expanded Understanding

Builds Lewis structures intuitively using patterns instead of trial and error

Why Method 2 is Useful

May expand octet (like S

Atoms Beyond Second Row

• NO₃⁻ and CO₃²⁻ have same valence e⁻ count

Isoelectronic Examples

• CH₄ and NH₄⁺ are isoelectronic

Recognizing these patterns helps when drawing structures quickly

For NCO⁻

Formal Charge Application

Ones with formal charges exceeding ±1 or placing like charges adjacent

Structures to Avoid