Intermolecular Forces

• Attraction forces that exist between molecules

• There are four types of intermolecular forces.

• Strongest to Weakest

Ion-ion or Ion-Dipole

• The forces holding ions together in ionic solids are electrostatic forces. Opposite charges attract each other.

• F = (q1 q2) / k r2

• Coulomb’s law of electrostatics

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.

Hydrogen Bonding

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

Application

• 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

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

• 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 = nsolute

• kgsolvent

• What is m (molality) of 50g, C6H12O11, of sugar in 117 g of water?

• 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-.

• 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!!

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 kf

• kf is a constant that is found experimentally for each solvent.

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 kb where kb is an experimentally determined constant for a particular solvent

Table of Constants

Vapor Pressure Lowering

Vapor Pressure Lowering

• For this, we need Mole Fraction!

• X = nsolute / ntotal where ntotal = nsolute + nsolvent

P = Po 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.

Osmotic Pressure

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

= 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