Lecture 21: Intermolecular Forces

  • Intermolecular Forces - Electrostatic interactions between molecules
  • Atoms within a molecule make covalent and ionic bonds within each other, but molecules also have interactions with other molecules
  1. Ion \rightarrow Ion Interactions
  2. Strongest intermolecular forces because they involve formal charges
  3. Ion \rightarrow Dipole Interactions
  4. Dipole - A side of an electron with excess and a side of deficiency
  5. The partially negative side is attracted to positive ions and vice versa
  6. Each ion can make several of these interactions which store energy
  7. Dipole \rightarrow Dipole Interactions
  8. Hydrogen bonds - OH, FH, or NH bonds - interact with each other
  9. Very strong dipole - Dipole bond because they are the most electronegative elements \rightarrow Most polarized bonds \rightarrow Strong Dipole \rightarrow Strong Dipole (Dipole Interactions)
  10. The greater the partial charge - the greater the interaction, but never as strong as formerly charged particles
  11. Van-der Waals/London Dispersion Force
  12. Any substance can have this bond
  13. Ion - Ion > Ion - Dipole > Dipole - Dipole > Van der Waals
  14. The stronger the forces between the molecules, the more heat energy is needed to provide to melt and boil the sample
  15. Non-Polar Covalent = Van der Waals (CH4)(CH_4)
  16. Polar Covalent = *check geometry* (CO2,H2O)(CO_2,H_2O)
  17. (Bf3Bf_3 (non-polar - van der waals), NH3NH_3 (polar-dipole-dipole))
  18. (CF4CF_4 (non-polar - van der waals), CH3FCH_3F (polar-dipole-dipole))
  19. Ionic = Ion-Ion (NaCl)(NaCl) 
  • Polar Molecules
  • Non-Polar Molecules
  • H-H
  • In a moment the electrons can be close to one H (temporary) and gives one H
  • Temporary Dipole \rightarrow Has a force of attraction with other molecules \rightarrow London Dispersion Force (weak bands)
  • Ideal Gas - PV=nRTPV = nRT, no attraction between molecules and there is infinite amount of space between them
  • Real Gas - Attraction between molecules

dispersion forces

  • Fluctuations in the electron distribution in atoms and molecules result in a temporary dipole
  • Region with excess electron density has a partial (-) charge
  • Region with depleted electron density has partial (+) charge
  • The attraction forces caused by these temporary dipoles are called dispersion forces = London forces
  • Magnitude depends on
  • Polarity of the electron cloud
  • Large molar mass = more electrons = larger electron cloud = increased polarizability = stronger attractions
  • Shape of the molecule
  • More surface - surface contact = larger induced dipole = stronger attraction
  • Larger molecules have more electrons \rightarrow Increased polarizability

effect of molecular size on size of dispersion force

  • Noble gases are all non-polar atomic elements
  • As the molar mass increases, the number of electron increases
  • Strength of dispersion forces increases
  • The stronger the attractive forces between molecules the higher the boiling point will be
  • If intermolecular forces are weaker, vapour pressure is higher
  • Boiling point is lower
  • Vapour Pressure \rightarrow
  • Intermolecular Force \rightarrow
  • Boiling Point \rightarrow
  • Colligative properties depend only on the number of dissolved particles in solution and not on their identity
  • Non-colligative properties depend on the identity of the dissolved species and the solvent
  • Vapour pressure of solutions
  • The vapour pressure of a solvent above solution is lower than the vapour pressure of the pure solvent
  • The solute particles replace some of the solvent molecules at the surface

raoult's law

  • PP solvent solution = XX solvent
  • PP solution is lower than PP solvent because the mole fraction is always less than 1
  • PP : Lowering of vapour pressure
  • The vapour pressure of the solution is directly proportional to the amount of the solvent in the solution
  • Boiling point
  • Temperature at which the vapour pressure is equal to atmospheric pressure
  • Vapour pressure lowering occurs at all temperatures
  • Results in the temperature required to boil the solution being higher than the boiling point of the pure solvent
  • Also results in the temperature required to freeze the solution lower than the freezing point of the pure solvent
  • Vant Hoff Factors
  • ii is the ratio of moles of solute particles to moles of formula units dissolved
  • The measured Vant Hoff factors are generally less than the theoretical due to ion pairing in solution
  • Therefore, the measured Vant Hoff factor often causes the TT to be lower than expected
  • Vapour pressure of the solution is lower than that of the solvent
  • Boiling point of a solution is higher than that of a solvent

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