Lecture Notes: Bonding and Hybridization

Overview

• Explanation of bonding through atomic orbitals
• Focus on understanding molecular geometries and hybridization

Lewis Structure and Bond Angles

• Example: Water (H₂O)
  • Hydrogen's s orbitals overlap with oxygen's p orbitals
  • Idealized bond angle from overlap: 90°
  • Actual bond angle: ~104°

Introduction to Hybridization

• Necessary to explain discrepancies in observed molecular geometries
• Learning Outcomes:
  • Explain atomic orbital hybridization
  • Determine hybrid orbitals for different molecular geometries

Key Concepts of Hybridization

• Existence: Only in covalent molecules (ignore ionic compounds)
• Shapes: Hybrid orbitals have new shapes, different from atomic orbitals
  • s: spherical
  • p: dumbbell
  • d: cloverleaf
• Formation: Combination of atomic orbitals
  • Mathematical linear combinations
• Properties:
  • Same shape and energy (degenerate)
  • Align with VSEPR theory
• Interactions:
  • Hybrid orbitals form sigma bonds
  • Unhybridized orbitals can form pi bonds

Example: Hybridization in Linear Molecules

• VSEPR Theory: Predicts 180° bond angle for linear molecules
• Hybridization Process:
  • Carbon atom: Take 2s and 2p orbitals
  • Form hybrid orbital: sp
  • Two sp orbitals formed
  • Two remaining p orbitals form pi bonds
• Bonding:
  • Triple bond: 1 sigma bond + 2 pi bonds
  • sp orbitals form sigma bonds

Visualizing Hybrid Orbitals

• Process:
  • Combine 2s and p orbitals of carbon
  • Consider wave character (positive and negative lobes)
  • Constructive interference forms new hybrid orbital shape
• Bonding with Hydrogen:
  • Hydrogen's s orbital bonds with carbon's hybrid orbitals
  • Aligns with VSEPR theory expectations

Conclusion

• Understanding hybridization helps reconcile discrepancies between simple Lewis structures and actual molecular shapes
• Provides a more accurate model of covalent bonding and molecular geometry