Topic 2.5 Notes – Lewis Diagrams

1. What Lewis Diagrams Are

A Lewis diagram shows only valence electrons and how atoms use them to bond.

It’s based on two big ideas:

• Octet rule: atoms tend to gain, lose, or share electrons to reach 8 valence electrons (like noble gases).
• Localized electron model: electrons are drawn in specific places, either:
  • Lone pairs (nonbonding electrons on one atom)
  • Bonding pairs (shared between atoms, drawn as lines)

Valence Electrons Refresher

You must count these correctly.

• For main-group elements, valence electrons = group number
  • Group 1 → 1 valence electron
  • Group 16 → 6 valence electrons
• Only valence electrons appear in Lewis structures.
• For ions:
  • Add electrons for negative charge
  • Subtract electrons for positive charge

If your total electron count is off, the entire structure will be wrong. This is one of the most common test mistakes.

2. The Two Types of Lewis Diagrams

a. Ionic Compounds

Ionic compounds form from electron transfer (metal → nonmetal).

You do not draw shared bonds. You show separate ions.

Example idea: $\ce{CaF2}$

• Ca loses 2 electrons → $\ce{Ca^{2+}}$
• Each F gains 1 electron → $\ce{F^{-}}$
• Two fluoride ions balance one calcium ion.

How to draw:

• Write correct formula (charges must cancel).
• Draw each ion separately.
• Use brackets around each ion.
• Put the charge outside the brackets.
• Show full octet on the anion.

You can draw two separate fluoride ions or draw one with a coefficient 2 outside the bracket.

Important pattern:

• Metals → cations
• Nonmetals → anions
• Total compound must be neutral

On multiple-choice questions, they love to test whether you balance the charges correctly before drawing anything.

b. Covalent Molecules and Polyatomic Ions

Covalent bonds form by sharing electrons.

Here electrons are placed between atoms as bonds.

Key structural rules:

• Count total valence electrons (adjust for charge).
• Choose central atom:
  • Usually least electronegative
  • Never hydrogen
  • Often carbon if present
• Connect atoms with single bonds first.
• Fill octets on outer atoms.
• Put remaining electrons on central atom.
• If central atom lacks octet, form multiple bonds.

For polyatomic ions:

• Enclose entire structure in brackets.
• Show overall charge outside.
• Make sure electron count matches the charge.

3. Step-by-Step Process for Drawing a Covalent Lewis Structure

Let’s organize it clearly. Follow this same order every time so you don’t miss points.

1. Count total valence electrons
   Include charge adjustment.
2. Draw skeletal structure
   Connect atoms with single bonds (2 electrons per bond).
3. Subtract bonding electrons
   Keep track as you go.
4. Complete octets on outer atoms
5. Place leftover electrons on central atom
6. If central atom does not have 8 electrons:
   • Convert lone pairs on neighboring atoms into double or triple bonds.
7. Final check:
   • Total electrons equal your starting count.
   • Octet rule satisfied (unless exception).
   • Brackets and charges shown if ion.

Here’s the whole process as a quick-reference flowchart you can picture during a test:

Lewis structure drawing flowchart

That last step saves people on tests. Always recount.

4. Exceptions to the Octet Rule

These show up a lot on quizzes.

a. Incomplete Octet

Some atoms are stable with fewer than 8 electrons:

• H → 2 electrons max
• Be → often 4 electrons
• B → often 6 electrons

If you force boron to have 8, you probably did something wrong.

b. Expanded Octet

Atoms in period 3 or below can exceed 8 electrons.

Common examples:

• P, S, Cl, Xe

Rules:

• Only the central atom expands.
• Period 2 elements (C, N, O, F) never expand.

c. Odd Number of Valence Electrons

If total valence electrons is odd, one atom will have an unpaired electron. Not every atom can satisfy the octet.

d. Multiple Bonds

Double bond = 4 shared electrons
Triple bond = 6 shared electrons

You form these when there aren’t enough electrons to give everyone an octet using single bonds alone.

5. Common Mistakes to Avoid

• Forgetting to adjust for ionic charge.
• Expanding octet on period 2 elements.
• Choosing hydrogen as central atom.
• Not using brackets for polyatomic ions.
• Forgetting final electron count check.

Most incorrect answers come from skipping that final recount.