1. Chapter 10 – Solutions & Their Properties
Molarity (M):
Moles of solute (mol) per liters (L) of solution:
Dilution Formula:
Used when preparing diluted solutions from concentrated ones.
Mole Fraction:
Used previously for gas problems.
XA = moles A = nA
total moles = ntot
Mole fraction of a component of the solution will equal moles of that component divided by the total moles present. The sum of the mole fractions should equal 1.

2. Mass Percent (%)
Mass percent is determined by the mass of the solute divided by the total mass of the solution, then multiplied by 100.
Example: If a solution is prepared by dissolving 24 g of NaCl in 152 g of water, calculate the percent, by mass, of NaCl.
Sometimes when the amounts are so small (trace amounts) we use parts per million (ppm) or parts per billion (ppb). For example, if we have 5 x 10-8 grams of arsenic in 1.00 grams of water that will be equivalent to 0.05 ppm of Arsenic in water.
For ppm we multiply by a factor of 106 and for ppb we multiply by a factor of 109.

3. Molality (m)
Molality is determined my moles of solute per kilograms of solvent.
m = moles of solute
kilograms of solvent
Conversions between concentration units:
Please reference example 10.4 on page 261.

4. Principles of Solubility
Like dissolves like…..meaning polar generally dissolves polar and nonpolar dissolves nonpolar.
Effects of Temperature:
Solids – the higher the temperature, generally the higher the solubility. An increase in temperature always shifts the position of an equilibrium to favor an endothermic process, which means H >0.
Gases – become less soluble as the temperature rises. Why?
Effects of Pressure:
Pressure only has an effect on gas-liquid solutions. Do pressure and the solubility of a gas have an inverse or direct relationship?
Henry’s Law:
Cg = kPg
Cg = concentration of the gas (M)
k = constant for the gas-liquid system – this will vary (M/atm)
Pg = Pressure of the gas (atm)
Example 10.5 – Page 264

5. Colligative Properties of Nonelectrolytes
Properties of solutions differ from the properties of the pure solvent.
The solution properties depend more so on the concentration of the solute particles.
Vapor Pressure Lowering
Boiling Point Elevation
Freezing Point Lowering

6. Vapor Pressure Lowering
Concentrated solutions evaporate more slowly than pure water, meaning they have a lower vapor pressure.
With an increase in concentration of solute the vapor pressure lowers – this is Raoult’s Law:
ΔP = X2 P10
ΔP = change in pressure (P10 - P1) - mmHg
X2 = mole fraction of solute
P10 = vapor pressure of the pure solvent - mmHg
Example 10.6 – page 265

7. Boling Point Elevation
A solution does not begin to boil until the temperature exceeds that of the solvent. The greater the concentration of the solute the higher the temperature needed to boil the solution:
ΔTb = kb x m
m = molality (m)
kb = Molal Boiling Point Constant (0C/m)
ΔTb = change in temp at which solution boils (Tb - Tb0)
-Tb = the temp at which the solution boils
-Tb0 = the temp at which the pure solvent boils
Example 10.7 – page 167

8. Freezing Point Depression
When a solution is cooled it does not begin to freeze until a temperature below the freezing point of the solvent is reached. The greater the concentration the lower the temperature required for freezing:
ΔTf = kf x m
m = molality
kf = Molal freezing point constant
ΔTf = change in temp at which solution freezes (Tf0 - Tf)
Tf = the temp at which the solution freezes
Tf0 = the temp at which the pure solvent freezes

9. Osmotic Pressure π = nRT = MRT V Example 10.8 – page 269
Osmosis is a process taking place through a membrane permeable to only the solvent is called osmosis. Water moves from a region where its vapor pressure or mole fraction is high to one in which it’s vapor pressure or mole fraction is low.
Osmotic pressure is directly proportional to the molarity (M) – this is why it is a colligative property:
π = nRT = MRT V
Example 10.8 – page 269