Chapter12.IntermolecularForces

Chapter 12 Overview

• Topics Covered:

  • 12.1 Types of Intermolecular Forces

  • 12.2 Physical States and Phase Changes

  • 12.3 Quantitative Aspects of Phase Changes

  • 12.4 Properties of the Liquid State

  • 12.5 Properties of Water

  • 12.6 The Solid State (skip crystal lattice and unit cells)

Intra- vs Intermolecular Forces

• Intramolecular Forces:

  • Attractive forces within a molecule.

    • Examples:

      1. Ionic Bonding: Attraction between cations and anions.

      2. Covalent Bonding: Attraction between nuclei and shared electron pairs.

      3. Metallic Bonding: Attraction between metal cations and delocalized electrons.

• Intermolecular Forces:

  • Attractive forces between different molecules, atoms, or ions.

    • Caused by partial charges (δ+ and δ-).

    • Weaker than intramolecular forces, explained by Coulomb's Law.

    • Electrostatic energy is proportional to charge and inversely proportional to distance.

Types of Intermolecular Forces

• Categories:

  1. Ionic

  2. Covalent

  3. Metallic

  4. Ion-dipole

  5. Hydrogen bonding

  6. Dipole-dipole

  7. Dipole-induced dipole

  8. London (dispersion) forces

• Characteristics:

  • Intermolecular forces help explain physical/chemical behavior in compounds.

Example Interaction: HCl

• Visualize interactions: HCl (g) exhibits dipole-dipole interactions.

Ion-Dipole Interactions

• Definition: Interaction between a fully charged ion and a polar molecule.

• Common in solutions of ionic compounds in polar solvents (e.g., water).

• Examples include:

  • CaBr2 (aq) with H2O (l).

  • (NH4)2SO4 (aq) with H2O (l).

  • HCl (aq) with H2O (l).

Hydrogen Bonding

• Definition: Involves H directly attached to F, O, or N and another dipole containing one of these electronegative elements.

• Molecules with hydrogen bonding have higher boiling points and melting points.

• Examples:

  • Water (H2O)

  • Ammonia (NH3)

  • Hydrogen fluoride (HF)

Dipole-Dipole Forces

• Attraction between the positive end of one polar molecule and the negative end of another.

• More orderly in solids vs liquids.

• Example molecules:

  • CH3OCH3 is less polar but cannot H-bond.

  • HCl and its polar nature shows dipole-dipole interactions.

Difference Between H-bond and Dipole-Dipole Forces

• H-bonding is directional and involves N, O, F hydrogen; stronger influence due to lone pairs.

• Dipole-dipole is a general attraction without the same level of directionality.

Types of Solids

• Crystalline Solids: Orderly structure, defined shapes.

• Amorphous Solids: Non-crystalline, poorly defined shapes.

• Types Include:

  • Atomic solids

  • Molecular solids

  • Ionic solids

  • Metallic solids

  • Network covalent solids

Properties Affected by Intermolecular Forces

• Intermolecular forces influence:

  • Boiling point (bp)

  • Melting point (mp)

  • Viscosity

  • Vapor pressure

  • Solubility

• Rule of Thumb (ROT): Stronger intermolecular forces result in higher boiling points and melting points.

Vapor Pressure and Intermolecular Forces

• The vapor pressure of a liquid is affected by the strength of intermolecular forces.

• Strong forces yield lower vapor pressure and vice versa.

Solubility Principle: "Like Dissolves Like"

• Substances exhibiting similar intermolecular forces can dissolve in each other.

• Major types of interactions in solution include:

  • Ion-dipole

  • Hydrogen bond

  • Dipole-dipole

  • Dipole-induced dipole

  • Dispersion

Phase Changes and Enthalpy (∆H)

• Phase changes (e.g., melting, freezing) involve heat energy transfer:

  • Endothermic: ( ∆H = + ) (e.g., melting)

  • Exothermic: ( ∆H = - ) (e.g., freezing)

• Each phase change has an associated enthalpy value that governs heat flow and energy absorbed or released.