Liquids and Solids Notes

12.1 Intermolecular Forces

  • Intermolecular forces are the forces between molecules, influencing physical properties.
  • Intramolecular forces are the forces within molecules, influencing chemical properties.
  • Intermolecular forces are more important in condensed phases (liquids and solids).
  • Types of Intermolecular Forces (Weakest to Strongest):
    • Dispersion: Temporary dipole, occurs in all molecules, most important in nonpolar molecules (e.g., Ar).
    • Dipole-Dipole: Occurs between polar molecules (e.g., CH3-O-CH3).
    • Hydrogen Bonding: Strong dipole-dipole interaction between molecules with H bound to N, O, or F (e.g., H2O).
    • Ion-Dipole Interactions: Occurs between ions and polar molecules (e.g., Na+ & H2O).
  • Factors Affecting Strength of Dispersion Forces:
    • Size: ↑ size → ↑ polarizability → ↑ boiling point.
    • Shape: Linear shapes have stronger dispersion forces than bulky shapes.
  • ↑ Polarity → ↑ boiling point.
  • Hydrogen bonding requires H bound directly to N, O, or F.
  • Ion-dipole attractions are the strongest temporary attractions between particles.

12.2 Properties of Liquids

  • Liquids' fluid properties (viscosity, surface tension, capillary action) arise from weak, temporary intermolecular forces.
  • Viscosity:
    • Resistance to flow.
    • ↑ IM → ↑ viscosity
    • ↑ T → ↓ viscosity
  • Surface Tension:
    • Energy required to move a molecule to the surface.
    • ↑ IM forces → ↑surface tension
  • Cohesion: Attraction between like particles.
  • Adhesion: Attraction between different particles.
  • Capillary Action: Ability of a liquid to flow against gravity up a narrow tube.
  • Water is a unique compound:
    • Excellent solvent.
    • High heat capacity & ∆Hvap.
    • High surface tension and capillarity.
    • Ice is less dense than liquid water.

12.3 Phase Changes and Heating Curves

  • Phase Changes:
    • Solid → Liquid: Melting (fusion) (endothermic).
    • Liquid → Solid: Freezing (crystallization) (exothermic).
    • Liquid → Gas: Vaporization (evaporation) (endothermic).
    • Gas → Liquid: Condensation (exothermic).
    • Solid → Gas: Sublimation (endothermic).
    • Gas → Solid: Deposition (exothermic).
  • Calculating Heat (q) for Phase Changes:
    • q = n * ΔH (n = moles, ΔH = enthalpy of fusion or vaporization).
  • Heating Curves: Show temperature changes as a substance is heated.
    • At melting point (MP) or boiling point (BP), all heat added goes towards breaking intermolecular forces (IMFs).
    • No temperature change during phase change.
  • Calculations:
    • Heat the solid from -20 to 0oC: q = mcΔT
    • Convert s → l: q = nΔHºfusion
    • Heat the liquid from 0 to 70oC: q = mcΔT
    • Total heat required (by Hess’s Law) = sum of the steps

12.4 Vapor Pressure, Boiling Point, and the Clausius–Clapeyron Equation

  • Volatile substances vaporize easily; nonvolatile substances do not.
  • Vapor Pressure:
    • ↑ IMF → ↓ vapor pressure.
    • ↑ T → ↑ vapor pressure.
  • Clausius-Clapeyron Equation:
    • Relates vapor pressure to temperature.
    • ln(P) = (-ΔHvap/R) * (1/T) + C
    • Two-Point Form: ln(P2/P1) = (-ΔHvap/R) * (1/T2 - 1/T1)
  • Boiling:
    • Occurs when vapor pressure equals surrounding pressure.
    • Normal boiling point: Boiling point at 1.00 atm.
  • Dynamic Equilibrium: In a closed system, evaporation and condensation rates equalize.

12.5 Phase Diagrams

  • Phase Diagram: Shows the phase of a substance under different pressure-temperature conditions.
  • Triple Point: P and T where all three phases are in equilibrium.
  • Critical Point: P and T above which the substance is a supercritical fluid.
  • Water's phase diagram has a negative slope for the solid-liquid boundary line.