Liquids and Solids

Intermolecular Forces (IMFs)

• Weak, temporary attractive forces between particles in solids and liquids.

• Based on temporary electrostatic interactions.

• Includes atoms and ions.

Types of IMFs

• Ion-dipole attractions: Between ions and polar molecules.

• Hydrogen bonding: H bonded to N, O, or F (2nd strongest IMF).

• Dipole-dipole attractions: Between polar molecules (asymmetric electron distribution).

• London dispersion forces: Found in all substances, due to temporary dipoles.

Factors Affecting London Dispersion Forces

• Polarizability: More "squishy" electron cloud = greater London Dispersion Forces.

• Larger atoms/molecules have greater London Dispersion Forces.

• Shape: Linear molecules have stronger IMFs than branched.

Predicting dominant IMF

• Ion-dipole: Present between an ion and a polar molecule

• Hydrogen bonding: Present when H is bonded to N, O, or F

• Dipole-dipole: Present in polar molecules

• London dispersion forces: Present in all substances

Properties and IMFs

• Strong IMFs: high melting/boiling points, viscosity, low vapor pressure.

• Weak IMFs: low melting/boiling points, viscosity, high vapor pressure.

• Compatibility of IMFs affects solubility.

Liquid Properties

• Viscosity: Resistance to flow; stronger IMFs = greater viscosity.

• Surface Tension: Liquids minimizing surface area; stronger IMFs = greater surface tension.

• Capillary Action: Liquid flowing against gravity in a tube; stronger IMFs = taller liquid column.

Phase Changes

• State functions: Enthalpy, Entropy.

• Exothermic: Gas to liquid (condensation), liquid to solid (freezing), gas to solid (deposition).

• Endothermic: Solid to liquid (melting/fusion), liquid to gas (vaporization), solid to gas (sublimation).

Heating Curves

• Phase changes are isothermal (constant temperature).

• Horizontal portion: Phase change, use q = \Delta H_{transition} \times (amount of material).

• Sloped portion: Heating a single phase, use q = (amount of material) \times (constant) \times \Delta T.

Phase Diagrams

• Shows phases at different temperatures and pressures.

• Triple point: Point where all three phases coexist.

• Supercritical fluid: Temperatures and pressures above the critical point.

Vapor Pressure

• Pressure exerted by a vapor when the liquid and vapor are in dynamic equilibrium.

• Stronger IMFs = lower vapor pressure; Weaker IMFs = higher vapor pressure.

• Vapor pressure increases with temperature.

• Volatile substance: High vapor pressure, evaporates easily.

Boiling Point

• Temperature at which vapor pressure equals ambient pressure.

• Normal boiling point: Vapor pressure equals 1 atm.

Clausius-Clapeyron Equation

• \ln P = -\frac{\Delta H_{vap}}{R} \frac{1}{T} + C

• Allows experimental measurement of \Delta H_{vap}.

Solids

• Crystalline: atoms arranged in an orderly repeating pattern.

• Amorphous: atoms arranged more like a liquid.

Types of Crystalline Solids

• Metallic: Atoms held together by metallic bonding (e.g., Cu, Fe).

• Ionic: Ions held together by ion-ion interactions (e.g., NaCl, MgO).

• Covalent-network: Atoms held together by covalent bonds (e.g., C, SiO_2).

• Molecular: Discrete molecules held together by intermolecular forces (e.g., H2, H2O).

Unit Cells

• Lattice points at corners.

• Fraction of atoms within a unit cell:

  • Corner: 1/8

  • Edge: 1/4

  • Face: 1/2

  • Body: 1