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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