Lewis Structures

Overview

Lewis structures (electron dot diagrams) show how valence electrons are arranged in molecules. They help predict molecular geometry, bonding, and reactivity.

Drawing Lewis Structures

Step-by-Step Method

1. Count total valence electrons

   • Add valence electrons for all atoms
   • Add electrons for negative charges
   • Subtract electrons for positive charges

2. Draw the skeleton structure

   • Place least electronegative atom in center
   • Hydrogen and fluorine are always terminal

3. Place bonding electrons

   • Connect atoms with single bonds (2 electrons each)

4. Distribute remaining electrons

   • Complete octets on terminal atoms first
   • Place remaining on central atom

5. Check octets and form multiple bonds if needed

   • If central atom lacks octet, convert lone pairs to bonds

Valence Electrons by Group

Group	Valence e⁻	Examples
1A	1	H, Li, Na
2A	2	Be, Mg, Ca
3A	3	B, Al
4A	4	C, Si
5A	5	N, P
6A	6	O, S
7A	7	F, Cl, Br
8A	8	He, Ne, Ar

Examples

Water (H₂O)

Total valence e⁻: $2(1) + 6 = 8$

    H—O—H
      ‥
      ‥

Carbon Dioxide (CO₂)

Total valence e⁻: $4 + 2(6) = 16$

O=C=O  or  :O=C=O:

Ammonia (NH₃)

Total valence e⁻: $5 + 3(1) = 8$

    H
    |
H—N—H
    ‥

Nitrate Ion (NO₃⁻)

Total valence e⁻: $5 + 3(6) + 1 = 24$ (One of three resonance structures)

Formal Charge

Helps determine the best Lewis structure.

$$\text{Formal Charge} = \text{Valence } e^- - \text{Lone pair } e^- - \frac{1}{2}(\text{Bonding } e^-)$$

Rules for Best Structure

1. Minimize formal charges
2. Place negative formal charges on more electronegative atoms
3. Avoid like charges on adjacent atoms

Example: CO₂

Structure: O=C=O

$$\text{O: } 6 - 4 - \frac{1}{2}(4) = 0$$ $$\text{C: } 4 - 0 - \frac{1}{2}(8) = 0$$ $$\text{Total formal charge} = 0 \checkmark$$

Octet Rule Exceptions

1. Incomplete Octets (< 8 electrons)

Common for: Be, B, Al

BF₃:    F—B—F    (B has 6 electrons)
            |
            F

2. Expanded Octets (> 8 electrons)

Only for Period 3+ elements with d orbitals

• SF₆: 6 bonds = 12 electrons around S
• PCl₅: 5 bonds = 10 electrons around P

3. Odd-Electron Molecules

Molecules with unpaired electrons (free radicals)

• NO: 11 valence electrons
• NO₂: 17 valence electrons

Resonance Structures

When a single Lewis structure cannot accurately represent a molecule.

Rules for Resonance

1. Atoms maintain same positions
2. Only electrons move
3. Total electrons remain constant
4. Formal charges may differ

Example: Benzene (C₆H₆)

The actual structure is a hybrid with delocalized electrons.

Common Polyatomic Ions

Ion	Formula	Lewis Structure Notes
Hydroxide	OH⁻	O has 3 lone pairs
Ammonium	NH₄⁺	N has no lone pairs
Carbonate	CO₃²⁻	3 resonance structures
Sulfate	SO₄²⁻	S has expanded octet
Phosphate	PO₄³⁻	P has expanded octet
Nitrate	NO₃⁻	3 resonance structures

Quick Tips

• Carbon always forms 4 bonds
• Nitrogen typically forms 3 bonds + 1 lone pair
• Oxygen typically forms 2 bonds + 2 lone pairs
• Halogens typically form 1 bond + 3 lone pairs
• Hydrogen always forms 1 bond, no lone pairs