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.