Lecture on Lewis Dot Structures

Introduction

• Focus on drawing Lewis dot structures.
• Ions form ionic compounds; covalent bonds form networks.

Representing Atoms

• Use chemical symbols for atoms.
• Valence electrons are represented as dots around the atom.

Carbon Example

• Electron configuration: 1s², 2s², 2p² -> 4 valence electrons.
• There are four coordination sites for atoms like carbon.
• Fill each coordination site with one electron before pairing.

Similarity Across Groups

• Elements in the same group have similar Lewis dot symbols due to similar valence electrons.

Drawing Lewis Dot Structures for Molecules

• Draw each atom with its valence electrons.
• Unpaired electrons from different atoms form covalent bonds (lines).
• Electrons must be paired either in bonds or as lone pairs.

Bond Formation Tendencies

• Carbon: 4 bonds.
• Nitrogen: 3 bonds, 1 lone pair (5 valence electrons).
• Oxygen: 2 bonds, 2 lone pairs.
• Fluorine: 1 bond, 3 lone pairs.

Types of Bonds

• Sigma Bond: Single covalent bond.
• Double/Triple Bonds: Involve pi bonds.
  • Carbon Dioxide Example: Unpaired electrons form additional bonds.
  • Pi Bond: Second bond in a double bond and second/third in a triple bond.

Bond Lengths

• Single bonds are the longest.
• Double bonds are shorter.
• Triple bonds are the shortest.

Formal Charge

• Occurs when electron contribution differs from typical valence.
• Ammonia (NH₃): Neutral nitrogen atom.
• Ammonium Ion (NH₄⁺): Nitrogen with formal positive charge (contributes one less electron).

Octet Rule

• Atoms seek 8 electrons (n=2 shell filled).
• Hydrogen: Needs 2 electrons (n=1 shell).
• Larger Atoms (Phosphorus, Sulfur): Can form 5 or 6 bonds.

Conclusion

• Combine unpaired valence electrons to form bonds.
• For further questions, contact the lecturer.