Valence Bond Theory explains how atoms form covalent bonds through orbital overlap. It's like a handshake between atoms, where their electron-filled orbitals come together to share electrons and create stable molecules.

Hybridization takes this idea further, mixing atomic orbitals to form new hybrid orbitals. This explains why molecules have specific shapes and bond angles, like the bent shape of water or the tetrahedral structure of methane.

Valence Bond Theory and Bonding

Fundamentals of Valence Bond Theory

• Valence Bond Theory explains covalent bonding through orbital overlap
• Orbital overlap occurs when atomic orbitals of bonding atoms come close enough to share electrons
• Sigma bonds form through end-to-end overlap of atomic orbitals along the internuclear axis
• Pi bonds result from side-by-side overlap of p orbitals perpendicular to the sigma bond
• Electron domain refers to the region in a molecule where electrons are likely to be found
  • Includes both bonding and non-bonding electron pairs
  • Shapes molecular geometry based on electron-electron repulsion

Types of Bonds and Orbital Interactions

• Sigma bonds constitute the primary covalent bond in molecules
  • Formed by head-on overlap of atomic orbitals
  • Can involve s-s, s-p, or p-p orbital combinations
• Pi bonds are secondary bonds that form in multiple bonds
  • Created by parallel overlap of p orbitals
  • Occur in double and triple bonds (ethylene, acetylene)
• Multiple bonds consist of one sigma bond and one or more pi bonds
  • Double bond: one sigma and one pi bond
  • Triple bond: one sigma and two pi bonds

Hybridization

Principles of Orbital Hybridization

• Hybridization involves the mixing of atomic orbitals to form new hybrid orbitals
• Hybrid orbitals explain molecular shapes and bond angles not accounted for by pure atomic orbitals
• Number of hybrid orbitals formed equals the number of atomic orbitals mixed
• Hybrid orbitals have uniform energy levels, unlike the original atomic orbitals

Types of Hybridization

• sp hybridization produces two hybrid orbitals
  • Results in linear molecular geometry (180° bond angle)
  • Occurs in molecules like beryllium chloride (BeCl₂) and acetylene (C₂H₂)
• sp² hybridization creates three hybrid orbitals
  • Leads to trigonal planar geometry (120° bond angles)
  • Found in molecules such as boron trifluoride (BF₃) and ethylene (C₂H₄)
• sp³ hybridization forms four hybrid orbitals
  • Generates tetrahedral geometry (109.5° bond angles)
  • Present in molecules like methane (CH₄) and ammonia (NH₃)

Applications of Hybridization

• Hybridization explains bonding in organic compounds
  • Alkanes primarily involve sp³ hybridization
  • Alkenes feature sp² hybridization for carbon-carbon double bonds
  • Alkynes utilize sp hybridization for carbon-carbon triple bonds
• Hybridization concepts apply to inorganic compounds
  • PCl₅ exhibits sp³d hybridization, resulting in trigonal bipyramidal geometry
  • SF₆ shows sp³d² hybridization, leading to octahedral geometry

Molecular Geometry

VSEPR Theory and Molecular Shapes

• VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry
  • Based on the principle that electron pairs repel each other
  • Arranges electron domains to minimize repulsion
• Molecular geometry describes the three-dimensional arrangement of atoms in a molecule
  • Determined by the number of bonding and non-bonding electron pairs
  • Influences physical and chemical properties of compounds
• VSEPR geometries include linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral
  • Linear: two electron domains (CO₂)
  • Trigonal planar: three electron domains (BF₃)
  • Tetrahedral: four electron domains (CH₄)

Bond Angles and Molecular Structure

• Bond angle measures the angle formed between three atoms connected by chemical bonds
  • Influenced by electron domain repulsion and atomic size
  • Varies based on the type of hybridization and presence of lone pairs
• Ideal bond angles for common geometries:
  • Linear: 180° (BeF₂)
  • Trigonal planar: 120° (BF₃)
  • Tetrahedral: 109.5° (CH₄)
• Lone pairs affect bond angles due to increased repulsion
  • Water (H₂O) has a bent structure with a 104.5° bond angle, smaller than the ideal tetrahedral angle
  • Ammonia (NH₃) has a pyramidal shape with 107° bond angles, also smaller than the ideal tetrahedral angle