Understanding Lewis Dot Structures

Sep 17, 2024

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.