Hybridization and Molecular Geometry

Electron Geometry and Hybridization

• Valence Bond Theory: Explains molecular shape based on the hybridization of atomic orbitals.
• Hybridization Types:
  • sp Hybridization:
  • Geometry: Linear
  • Electron Groups: 2
  • Bond angles: $180^{\circ}$
  • Example: $\text{BeCl}_2$
  • sp² Hybridization:
  • Geometry: Trigonal planar
  • Electron Groups: 3
  • Bond angles: $120^{\circ}$
  • Example: $\text{BF}_3$
  • sp³ Hybridization:
  • Geometry: Tetrahedral
  • Electron Groups: 4
  • Bond angles: $109.5^{\circ}$
  • Example: $\text{CH}_4$

Hybridization with Lone Pairs

• Lone Pairs Count as Electron Groups:
  • For example, in $\text{NH}_3$, there are 4 electron groups (3 bonds + 1 lone pair) which means it is also sp³ hybridized despite having a lone pair.

Molecular Examples

• Ethane ($\text{C}2\text{H}6$):

  • Each carbon is sp³ hybridized (4 bonds)
  • All bonds are sigma bonds.

• Methanol ($\text{CH}_3\text{OH}$):

  • Carbon: sp³ hybridized (4 total bonds)
  • Oxygen: also sp³ hybridized (4 groups: 2 bonds + 2 lone pairs)
  • Shape is not tetrahedral due to lone pairs but electron geometry is tetrahedral.

General Trends in Bonding

• Counting Bonds:
  • Sigma Bonds: Every single bond is a sigma bond. In double bonds, the first bond is a sigma bond and second is a pi bond.
  • Triple Bonds: 1 sigma and 2 pi bonds.

Electrons and Energy

• Bond Strengths:
  • Strength increases from single to triple bonds:
  • Ethane (C-C single bond)
  • Ethylene (C=C double bond)
  • Acetylene (C≡C triple bond)
  • Energy needed to break bonds is not a simple multiple due to the nature of bonding between sigma and pi bonds.

Special Cases in Bonding

• Double Bonds:

  • Double bonds arise from sigma bonds plus an additional overlap of leftover p orbitals forming pi bonds.
  • Ethylene ($\text{C}2\text{H}4$):
  • sp² hybridization with 3 sigma bonds and 1 pi bond from leftover p orbital overlap.

• Triple Bonds:

  • Acetylene ($\text{C}2\text{H}2$):
  • sp hybridization with 2 pi bonds and 1 sigma bond due to two leftover p orbitals.

• Benzene ($\text{C}6\text{H}6$):

  • Each carbon is sp² hybridized, contributing to resonance and showing bond character across the entire ring.

Resonance Structures

• Sulfate Ion ($\text{SO}_4^{2-}$):
  • Hybridization can be determined by electron groups.
  • Can exhibit resonance due to different placement of double bonds among oxygens.

Summary

• Understanding hybridization allows for predicting molecular geometry and bond angles effectively.
• Recognizing lone pairs, bonding types (sigma vs. pi), and resonance is essential for mastering molecular structures.
• Concept application through examples helps solidify understanding of complex hybrid structures and molecular properties.