intermolecular force

• States of matter


  The fundamental difference between states of matter is the strength of the intermolecular forces of attraction
• Condensed phases


  • Solids and liquids
• Intermolecular forces


  Attractions between molecules
• Types of intermolecular forces

  • Dispersion forces
  • Dipole-dipole forces
  • Hydrogen bonding
  • Ion-dipole forces

  
• London dispersion forces


  Attractions between an instantaneous dipole and an induced dipole
• Factors affecting the strength of London dispersion forces include size and polarizability of the molecules
• Dipole-dipole forces


  Attractions between polar molecules

  
• For molecules of approximately equal mass and size, the more polar the molecule, the higher its boiling point
• Hydrogen bonding


  Occurs when hydrogen is bonded to highly electronegative atoms like N, O, or F

  
• Ion-dipole forces


  Attractions between an ion and a polar molecule

  
• Liquid properties affected by intermolecular forces include boiling point, viscosity, surface tension, and vapor pressure
• Boiling point


  The temperature at which a liquid converts into a gas
• Strong intermolecular forces of attraction

  High boiling point
• Vapor pressure


  A measure of the tendency of a material to change into the gaseous or vapor state
• Strong intermolecular forces of attraction

  Low vapor pressure
• Crystalline solids


  • Held together by ionic, covalent, or metallic bonds
• Phase changes


  Transitions between the states of matter
• Types of phase changes

  • Melting
  • Boiling
  • Sublimation
  • Deposition
  • Condensation
• Endothermic process


  Absorption of heat/energy
• Heating curve


  A plot of temperature and heat added to a substance
• Cooling curve


  A plot of temperature and heat removed from a substance
• Solid-liquid equilibrium


  When solid and liquid phases of a substance coexist
• Liquid-vapor equilibrium


  When liquid and gas phases of a substance coexist
• Phase diagram


  Shows the physical states of a substance under different conditions of temperature and pressure

  
• Supercritical fluid


  A phase that occurs for a gas at a specific temperature and pressure such that the gas will no longer condense to a liquid regardless of how high the pressure is raised
• Critical temperature


  The temperature above which a gas cannot be liquefied by pressure alone
• The phase diagram of water shows the three areas of solid, liquid, and gas, as well as the critical temperature and pressure
• The phase diagram of CO2 shows the supercritical fluid region