3.2 Valence Bond Theory and Hybridization

Valence Bond Theory explains how atoms form covalent bonds by overlapping orbitals. It's all about sharing electrons to create stable molecules. The theory helps us understand why some bonds are stronger than others.

Hybridization takes it a step further, showing how atoms mix their orbitals to form specific shapes. This process is key to predicting molecular geometry and understanding why molecules look the way they do.

Covalent bond formation

Valence bond theory explanation

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• Describes covalent bonding as the overlap of atomic orbitals to form a chemical bond
• Results in the sharing of electrons between atoms
• Overlapping atomic orbitals form a new molecular orbital that contains the shared electrons
  • Lowers the overall energy of the system
  • Stabilizes the molecule
• Strength of the covalent bond depends on the extent of orbital overlap
  • Greater overlap results in a stronger bond

Types of covalent bonds

• Sigma (σ) bonds
  • Formed by direct, head-on overlap of atomic orbitals along the internuclear axis
• Pi (π) bonds
  • Formed by lateral overlap of p orbitals above and below the internuclear axis
• Resonance describes molecules that cannot be accurately represented by a single Lewis structure
  • Actual structure is an average of multiple contributing structures (benzene, ozone)

Orbital hybridization and geometry

Hybridization process

• Combines atomic orbitals to form new hybrid orbitals
  • Shapes and energies different from original orbitals
• Occurs when energy difference between atomic orbitals involved in bonding is small
  • Allows orbitals to mix and form hybrid orbitals
• Hybridized orbitals are more energetically favorable for bonding
  • Results in more stable molecules with specific geometries (methane, ammonia)

Factors determining hybridization

• Type of hybridization (sp, sp², or sp³) depends on number of electron domains around central atom
  • Electron domains include bonding and nonbonding electrons
• Hybrid orbitals oriented in space to minimize electron repulsion and maximize bond angles
  • Determines molecular geometry of molecule
• Explains observed geometries that deviate from VSEPR theory predictions (water, carbon dioxide)

Hybridization state prediction

Determining hybridization state

• Based on number of electron domains (bonding and nonbonding) surrounding atom
• sp hybridization: two electron domains around central atom
  • Results in linear geometry with 180° bond angles (carbon dioxide, acetylene)
• sp² hybridization: three electron domains around central atom
  • Results in trigonal planar geometry with 120° bond angles (boron trifluoride, formaldehyde)
• sp³ hybridization: four electron domains around central atom
  • Results in tetrahedral geometry with 109.5° bond angles (methane, ammonia)

Factors affecting hybridization

• In molecules with multiple atoms, hybridization state determined independently for each atom
  • Based on its bonding environment
• Presence of multiple bonds (double or triple bonds) affects hybridization state
  • Additional pi bonds require unhybridized p orbitals for formation (ethylene, acetylene)
• Exceptions to general hybridization rules can occur
  • Molecules with significant resonance (benzene)
  • Presence of steric effects (hindered rotation in biphenyls)