1. Intermolecular Forces Attraction forces that exist between molecules There are four types of intermolecular forces. Strongest to Weakest

2. Ion-ion or Ion-Dipole The forces holding ions together in ionic solids are electrostatic forces. Opposite charges attract each other. F = (q 1 q 2 ) / k r 2 Coulomb’s law of electrostatics

3. Dipole/Dipole One end (pole) of the molecule has a partial positive charge while the other end has a partial negative charge. The molecules will orientate themselves so that the opposite charges attract principle operates effectively.

4. Hydrogen Bonding The hydrogen on one molecule attached to O or N that is attracted to an O or N of a different molecule.

5. Application

7. London Forces: AKA Van Der Waals or Dispersion forces VERY, VERY weak Caused by the attraction of the electrons of one atom for the protons of another Molecules must be very, very close. Important only in gases

8. Solutions Composed of solute & solvent Water is the universal solvent Like dissolves like!! Polar solvents dissolve polar solutes Non-polar solvents dissolve non-polar solutes

9. Colligative Property – the property of a solution that is dependent on the number of particles dissolved in the solvent. Freezing point, boiling point, vapor pressure, and osmotic pressure But First, a new way to talk about concentration: Molality = n solute kg solvent What is m (molality) of 50g, C 6 H 12 O 11, of sugar in 117 g of water?

10. The van't Hoff factor, symbol i, expresses how may ions and particles are formed (on an average) in a solution from one formula unit of solute. If we put 1 mole of sugar in water, we get 1 mole of particles. HOWEVER, if we put 1 mole of NaCl in water, we get 2 mole of particles – 1 mole of Na + and 1 mole Cl -.

11. So, soluble compounds such as Group 1A salts, strong acids and strong bases produce a molality equal to the number of particles actually made. 1 mole of NaCl produces the same number of particles as 2 moles of sugar. We have to remember our solubility rules!!

12. Freezing Point Depression The freezing point of a pure solvent is higher than the freezing point of a solution made with that solvent. Example: salt on roads when there is danger of ice forming. To find the change in the freezing point when a solute is added:  m k f k f is a constant that is found experimentally for each solvent.

13. Boiling Point Elevation Adding a solute to a solvent cause the boiling point to increase. To find the change in boiling point for a particular solvent:  T = m k b where k b is an experimentally determined constant for a particular solvent

14. Table of Constants

15. Vapor Pressure Lowering

16. For this, we need Mole Fraction! X = n solute / n total where n total = n solute + n solvent  P = P o X This equation is referred to as Raoult’s Law which says simply that the vapor pressure above a solution is proportional to the mole fraction of the solute.

17. Osmotic Pressure

18. The movement of solvent across a semi-permeable membrane to establish equal concentration.

19.  = MRT molarity which is the amount of solute added Osmotic pressure,  increases because of solute The constants R and T are the ideal gas law constant and the system temperature