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Intermolecular Forces (IMF)

• Attractive forces between particles in a substance

• Generally weaker than intramolecular forces (ionic, covalent, metallic bonds) within compounds

• Strongest in polar molecules due to uneven charge distribution

Types of Intermolecular Forces

Dipole-Dipole Forces

• Attractive force between polar molecules

• Formed when equal and opposite charges are separated over a short distance

• Direction of dipole represented from positive to negative pole

Hydrogen Bonding

• Occurs when a hydrogen atom bonded to a highly electronegative atom (N, O, or F) attracts an unshared pair of electrons from neighboring electronegative atoms

• Represented by dotted lines

• Example: Attraction between hydrogen atoms in H₂O and oxygen atoms in another H₂O molecule

London Dispersion Forces

• Arises from instantaneous and temporary dipoles due to electron motion

• Present in all atoms and molecules

• Significant in noble gases and nonpolar molecules

Effects of IMF

• Viscosity: Resistance of a fluid to movement; increases with stronger IMFs

• Boiling Point: Amount of kinetic energy needed to overcome attractions between particles; increases with stronger IMFs

  • Example substances ranked by boiling points with ionic and various types of intermolecular forces

Summary of IMF Strength

Comparison of "Inter" and "Intra"

• Inter = between/among substances (intermolecular forces)

• Intra = within a substance (intramolecular forces = chemical bonds)

  • Examples include molecules that exhibit hydrogen bonds (e.g. adenine, thymine, cytosine, guanine)

Overview of Molecular Polarity

• Polarity is determined by the presence of partial charges within the molecule

  • Example: H₂O has partial charges leading to polar nature

• Nonpolar molecules do not have partial charges; e.g., O₂

Determining Polarity with BEND

• Bond Electronegativity Difference (BEND):

  • AEN = 0 → Nonpolar

  • AEN ≠ 0 → Polar; more significant AEN indicates increased polarity

• Example ranking (from least polar to most polar): O₂, CO₂, HCl, HF

Molecular Shapes

• Visualization methods include notation and 3D models

• Types of shapes:

  • Linear: 2 electron domains (Examples: O₂, HCN)

  • Trigonal Planar: 3 electron domains (Example: BF₃)

  • Tetrahedral: 4 electron domains (Example: CH₄)

  • Trigonal Bipyramidal: 5 electron domains (Example: PCl₅)

  • Octahedral: 6 electron domains (Example: SF₆)

VSEPR Theory

• Explains the three-dimensional arrangement of electron domains around a central atom

• Electron pairs repel each other, influencing molecular shape

  • Electron domains can be bonding pairs or lone pairs

Specific Molecular Shapes Examples

• Bent Shape: 4 electron domains with 2 bonding pairs, 2 lone pairs (e.g. H₂O)

• Trigonal Pyramidal Shape: 4 electron domains with 3 bonding pairs, 1 lone pair (e.g. NH₃)

This summary encompasses the various intermolecular forces, their impacts on physical properties, molecular polarity principles, and molecular geometries, which are fundamental for understanding chemical behavior and interactions.