Colligative Properties-number of particles is key, not the nature of the particles
Properties that only depend on the amount of solute dissolved and not on the identity of the solute are called colligative properties. The properties affected in this way include boiling point, melting point and osmotic pressure. The mechanism that causes the change is the lowering of the vapor pressure of the solvent.
Therefore, the magnitude of the effect will depend on the nature of the solvent but not on the nature of the solute. Solute particles interact with solvent particles. By the way, the solute has to be nonvolatile (does not become a gas easily) as this would seriously mess up the situation since it is vapor pressure lowering of the solvent we are talking about here. If the solute is interacting with the solvent, then fewer solvent particles are available to escape into the vapor phase and the vapor pressure is lowered at any given temperature.
Freezing Point Depression
∆tf = Kf m n
∆tf = the CHANGE in the freezing point not the new freezing point
Kf = the molal freezing point constant and is different for each solvent (i.e., look it up in a table)
m = molality (mol solute/Kg solvent)
n = number of solute particles
Boiling Point Elevation
∆tb = Kb m n
∆tf = the CHANGE in the boiling point not the new boiling point
Kf = the molal boiling point constant and is different for each solvent (i.e., look it up in a table)
m = molality (mol solute/Kg solvent)
n = number of solute particles
Osmotic Pressure
The best way to understand osmosis is remembering that nature does not like two different substances in contact with one another to remain unmixed especially if the substances are liquids (including solutions) and gases. Because solid particles are held in fixed positions, random mixing usually does not occur. Osmosis is a more biological phenomenon as all cell membranes are "semi-permeable" (allowing some particles to pass through while blocking others) and, therefore, osmosis can occur in biological systems and does. If you drink an excess of water, you can "swamp" your cells with too much "solvent" and dilute out your electrolytes enough to even cause death. This is an example of osmosis. Have you ever tried the experiment where you put a carrot in really salty water? What happens to the carrot is that because the solution in the cells of the carrot is more dilute than the salt water, the cells expel water to try to bring the salt concentration on both sides of the cell wall to about the same level. The result is a shriveled carrot.
The pressure required to stop osmosis, known as osmotic pressure, π, is
π = MRT where M is the molarity of the solution
parts per million (ppm)
parts per billion (ppb)
Solution Stoichiometry
Most chemistry is done using solutions. Therefore, most stoichiometry involving chemical reactions involves concentrations of solutions. There is really no difference in how you work these other than how you get to moles and the conversion at the end (instead of g, it may be volume or M).
rewrite the big M! Any time you see M in a problem, rewrite it to help you get into moles. For example, 2.5 M NaCl is 2.5 mol NaCl/L of solution. You are already in moles, multiply by volume given and you are in moles.
