(CHEM)VSEPR Theory and Intermolecular Forces Overview

Introduction to VSEPR Theory

• Electron Domains: Evaluate where electrons exist around central atoms in molecules, which is critical in determining molecular shape.
• Lone Pairs: Count as one electron domain. Example: Oxygen's unshared electrons form a lone pair.
• Bonding: Each bond is also an electron domain. A double bond counts as one electron domain.

Counting Electron Domains

• To assess molecular shapes, determine electron domains around the central atom which can include lone pairs and bonds.
  • Example: In ozone (O₃), the central atom has:
  • 1 double bond counts as 1 domain
  • 1 single bond counts as 1 domain
  • 1 lone pair counts as 1 domain
  • Total: 3 electron domains.

VSEPR Model

• VSEPR Definition: Valence Shell Electron Pair Repulsion (VSEPR) predicts molecular shape based on electron pair repulsion.
• Principle: Electron pairs will arrange themselves to minimize repulsion, leading to specific 3D arrangements of atoms.
  • Linear Shape: If there are 2 electron domains, expect a linear molecular shape (180° bond angle).

Electron Domain Geometry vs. Molecular Geometry

• Electron Domain Geometry (EDG): Involves all electron domains including lone pairs.
• Molecular Geometry (MG): Refers only to the arrangement of bonded atoms, excluding lone pairs.
  • Example: Trigonal planar EGD might lead to a bent MG if lone pairs are present.

Common Geometries and Angles

• Trigonal Planar: 3 electron domains, 120° angles.
• Tetrahedral: 4 electron domains, 109.5° angles.
• Trigonal Bipyramidal: 5 electron domains:
  • Equatorial: 120°
  • Axial: 90°
• Bent: Results from lone pairs affecting bond angles (e.g., H₂O vs. CO₂).

Determining Molecular Geometry

• To derive molecular shapes:
  1. Draw Lewis Structures: Visualize the arrangement of all valence shell electrons.
  2. Identify Electron Domains: Count both lone pairs and bonding pairs around the central atom.
  3. Apply VSEPR Rules: Use the count to identify potential geometry.

Molecular Shape Examples

• Molecular Geometry Variations:
  • NH₃ (ammonia): Tetrahedral EDG but trigonal pyramidal MG due to one lone pair.
  • H₂O: Tetrahedral EDG but bent MG (2 lone pairs influence shape).

Intermolecular Forces Overview

• Systems: Influences between neighboring molecules in solids, liquids, and gases.
  • Strength Ranking: Solid > Liquid > Gas (stronger interactions in solids).

Types of Intermolecular Forces

• Dipole-Dipole Interactions: Forces between polar molecules. Strength depends on the magnitude of dipoles.
• Hydrogen Bonds: Special type of dipole-dipole interaction, occurs when hydrogen is bonded to electronegative atoms (N, O, F).
• Dispersion Forces (London Forces): Present in all molecules due to instantaneous dipoles; exceptionally significant in nonpolar molecules.
• Ion-Dipole Interactions: Colloquial attraction between ions and polar molecules.

Bonding Theory and Hybridization

• Hybridization Concept: During bond formation, atomic orbitals mix to form new hybrid orbitals. Steps include:
  1. Promotion: Electrons are excited to higher energy levels.
  2. Hybridization: Orbitals combine to form equivalent hybrid orbitals.
• Example: Beryllium hybridizes from a 1s2 2s2 state to a mixture of sp orbitals, allowing for equivalent bonding characteristics in molecules.

Energy Considerations in Bonding

• Energy Changes:
  • Energy is absorbed to break bonds and released when forming bonds. The stability of molecules correlates with their total potential energy.

Conclusion

• VSEPR Theory and bonding concepts outline fundamental chemical interactions that dictate molecular shapes and behavior in various states of matter. Understanding both intermolecular and intramolecular forces enhances comprehension of chemical properties and reactivity.