ch11_IMF

Chapter 11: Intermolecular Attractions and the Properties of Liquids and Solids

Introduction

  • Focus on intermolecular forces differentiating states of matter (gases, liquids, solids).

  • Intermolecular attractions significantly influence physical properties.

Intermolecular Forces

Differences Between States

  • Gases: Expand to fill their container.

  • Liquids: Retain volume, shape is not fixed.

  • Solids: Retain both volume and shape.

Physical State Determinants

  • Average kinetic energy of particles influences state.

    • Kinetic energy (KE) is proportional to average temperature (T).

  • Energy of intermolecular attractions affects physical properties.

    • Molecule packing and strength of intermolecular attractions.

Phase Changes

  • Gas to Liquid/Solid: Molecules must get closer by cooling or compressing.

  • Liquid/Solid to Gas: Molecules must move apart by heating or reducing pressure.

  • Temperature decrease leads to decreased kinetic energy; molecules lose enough energy to overcome attraction.

Types of Forces

Intermolecular vs Intramolecular Forces

  • Intramolecular forces:

    • Strong covalent bonds that hold atoms within a molecule (e.g., HCl).

  • Intermolecular forces:

    • Weaker attractions between molecules (e.g., vaporization of HCl).

Electronegativity

  • Measure of an atom's ability to attract electrons in covalent bonds.

  • Influences molecular polarity.

Bond Dipoles

  • Unequal sharing of electrons due to differing electronegativity leads to electron density unevenness.

  • Bond Dipoles indicated by delta (δ) notation, representing partial charges.

Net Dipoles

  • Symmetrical molecules with polar bonds can be non-polar if bond dipoles cancel each other out.

  • Asymmetrical molecules are polar, exhibiting permanent net dipoles, influencing intermolecular interactions.

Breaking Intermolecular Forces

  • Melting or boiling indicates breaking of intermolecular forces, NOT covalent bonds.

  • Essential for understanding gas behavior and condensed states of matter.

Types of Intermolecular Forces

  1. Dipole-dipole forces:

    • Occur between polar molecules with dipole moments.

    • Polar molecules align positive and negative charges.

    • Strength of force increases with polarity.

  2. Hydrogen bonds:

    • Special dipole-dipole interaction; strong due to hydrogen bonding with O, N, or F.

    • Examples: H-F, H-O, H-N.

  3. London dispersion forces:

    • Present in all substances; weakest force, result from temporary dipoles.

  4. Ion-dipole forces:

    • Occur between ions and polar molecules; strong due to full charges.

  5. Ion-induced dipole forces:

    • Interaction between an ion and a non-polar molecule, depends on ion charge and polarizability.

Physical Properties Predictions

  • Predict boiling and melting points based on strengths of intermolecular attractions.

    • Larger molecules typically exhibit stronger intermolecular forces.

Solubility Principles

  • "Like dissolves like": polar substances dissolve in polar solvents, non-polar in non-polar.

  • Absence of polarity leads to insolubility.

Phase Changes

General Principles

  • Phase changes involve motion and energy changes in molecules as temperature varies.

  • Changes include:

    • Liquid to gas (evaporation, heat added)

    • Solid to gas (sublimation, endothermic)

    • Gas to liquid (condensation, exothermic)

Phase Equilibria

  • At equilibrium, rates of evaporation and condensation equal in closed systems.

    • Equilibrium can also be observed in sublimation and deposition processes.

Boiling Points

  • Determined by vapor pressure equaling atmospheric pressure. Higher strength of intermolecular forces raises boiling points.

  • Normal boiling point defined as the boiling point at standard atmospheric pressure.

Energy Changes

  • Molar heat of fusion (heat absorbed when solid melts), vaporization, and sublimation increase with stronger intermolecular forces.

Phase Diagrams

  • Illustrate relationships between pressure, temperature, and phases; include critical and triple points where states coexist.

Supercritical Fluids

  • Exhibiting unique properties above a critical temperature and pressure, combining gas and liquid characteristics for efficiency in various applications.