Lecture Notes: Drawing Lewis Dot Structures

Introduction

• Comprehensive lesson on drawing Lewis dot structures.
• Topics include the octet rule, exceptions to the octet rule, resonance structures, and formal charges.
• Importance for understanding molecular geometry in future chapters.
• Speaker: Chad (Chad’s Prep).
• Courses available at chadsprep.com, including MCAT, DAT, and OAT prep.

Lewis Dot Structures Basics

• Represents valence electrons, the outermost shell of electrons involved in chemical reactions.
• Valence electrons in atoms:
  • Group 1 metals: 1 valence electron.
  • Group 2 metals: 2 valence electrons.
  • Group 13: 3 valence electrons.
  • Group 14: 4 valence electrons.
  • Group 15: 5 valence electrons.
  • Group 16: 6 valence electrons.
  • Halogens (Group 17): 7 valence electrons.
  • Noble gases (Group 18): 8 valence electrons (Helium has 2).
• Focus primarily on non-metals in molecular compounds.

Octet Rule

• Atoms aim for a filled octet of 8 valence electrons through transfer (ionic bonding) or sharing (covalent bonding).
• Exceptions to the octet rule:
  • Under the octet rule: Hydrogen (2 electrons), Beryllium (4 electrons), Boron and Aluminum (6 electrons).
  • Expanded octets: Atoms in period 3 or higher (e.g., Sulfur, Phosphorus) can have more than 8 electrons by utilizing d-orbitals.
  • Odd number of electrons: Molecules like NO have an odd number of electrons, violating the octet rule.

Drawing Lewis Structures: Step-by-Step

1. Count valence electrons: Total available electrons for bonding.
2. Determine central atom: Usually the least electronegative (not hydrogen).
3. Create skeleton structure: Connect atoms with single bonds.
4. Fill octets for outer atoms: Complete the valence shells for atoms surrounding the central atom.
5. Place leftover electrons on the central atom: Place any remaining electrons on the central atom as lone pairs.
6. Check central atom: Ensure the central atom has a full octet. If not, create double or triple bonds as needed.
7. Formal charges: Calculate to evaluate the best resonance structures.

Resonance Structures

• Occur when multiple valid Lewis structures exist for a molecule.
• Electrons are delocalized across multiple atoms.
• The best resonance structure minimizes formal charges.

Formal Charges

• Used to determine the most stable Lewis structure.
• Computed using the formula:
  • Valence electrons - (Non-bonding electrons + 1/2 Bonding electrons).
• Prefer structures with formal charges closest to zero.

Examples

1. CCl₄: Carbon as the central atom, chlorines bonded with single bonds.
2. NF₃: Nitrogen as the central atom, fluorines bonded with single bonds.
3. HCN: Carbon in the center, triple bond with nitrogen.
4. CO₂: Carbon with double bonds to two oxygens.
5. N₂O (tricky): Nitrogen central, shows resonance.
6. SF₄ and XeF₄: Examples of expanded octet structures.
7. SO₄²⁻: Use of formal charges to optimize structure.
8. NO₃⁻: Equivalent resonance structures, showing delocalized bonding.

Conclusion

• Importance of being proficient in drawing Lewis structures for understanding molecular shapes.
• Recommended practice for proficiency before exams.
• Encouragement to use available resources for further study.