Chapter 12 Solutions
12.1
____________________ - component that is present in the greatest amount
- determines the state of matter in the solution
____________________- other component(s)
Types of solutions:
Example---------------------Type------------------------Solute--------------------Solvent
air
soda water
beer
sea water
dental filling
Au-Ag alloy
12.2 Solution Concentrations
molarity (M) =
EXAMPLE: 8.40 g of NaF is dissolved in 95.0 g of water to make 100 mL of solution. What is the molarity?
Percent by Mass
Percent by Volume
Mass/Volume Percent
sulfuric acid 35.7% by mass
95% ethanol by volume
saline 0.92% NaCl mass/volume
0.10% by volume - alcohol level
EXAMPLE: 10.0 mL of methanol (d = 0.720 g/mL) is dissolved in 97.0 g of water to make 100 mL of solution.
Parts per million (ppm)
Parts per billion (ppb)
Parts per trillion (ppt)
1 ppm fluoride in drinking water
What is the ppb of Li+ in 1.00 x 10-5 M LiCl?
Molality (m) =
8.40 g of NaF is dissolved to 100 mL of solution with water. Density of solution is 1.020 g/mL.
Mole Fraction (xi) =
Mole percent = mole fraction x 100%
8.40 g of NaF is dissolved to 100 mL of solution with water. Density of solution is 1.020 g/mL.
12.3 Energetics of Solution formation
Why do some substances dissolve in a given solvent, but others do not?
1. pure solvent ---> separated solvent molecules
2. pure solute ----> separated solute molecules
3. separated solvent and solute molecules ---->
net energy Hsol =
1. All IMF are of comparable strength
2. IMF between solute and solvent molecules are stronger than other IMF.
3. IMF between solute and solvent molecules are slightly weaker than other IMF.
Hsol > 0 , endothermic,
mix due to random motion (entropy)
4. IMF between solute and solvent molecules are much weaker than other IMF.
do not mix
(immiscible liquids)
Rule: Like dissolves like
Aqueous Solutions of Ionic Compounds
Interionic forces - _____________________________
ion-dipole forces - __________________________________
ions become surrounded by water molecules (hydration)
________________ force determines solubility
(solubilities rules in ch 4)
Examples: Predict solubility:
octane + heptane
octane + water
water + alcohol
water + NaCl water + AgCl
12.4 Equilibrium in Solution Formation
Dissolving process - ions go into solution
add more solid, keep dissolving, some collide and reform
eventually appears that no more will dissolve = ________________________
_________________: solution that is in equilibrium
(the concentration of the solute = SOLUBILITY),
_________________: solution not at equilibrium and in which more of the substance can dissolve
_________________: solution that contains more dissolved substance than a saturated solution, this system is not at equilibrium, unstable
12.5
EFFECTS OF TEMPERATURE
most gases _______________ solubility at higher temperature: explain based on KE?
the solubility of solids vary
ENDOTHERMIC dissolving
EXOTHERMIC dissolving
PRESSURE - Henry's law
double the pressure, __________________________ the solubility of a gas
12.6 Vapor Pressure of solutions
the presence of a solute _____________ the vapor pressure of the solvent in a solution
12.7 and 12.9
Freezing Point Depression and Boiling Point Elevation
The extent of the vapor pressure lowering depends on the __________________________ of the solute present in a solution.
In dilute solutions, use __________________
New BOILING POINT can be calculated by:
Tb =
Tb =
__________________ = Molal Boiling Point Elevation constant
__________________ = molality
__________________ = ionization factor (Van't Hoff factor)
(table 12.2 Constant Values)
for molecular substances i =
for dilute ionic substances, i =
Which of the following water solutions have the lowest boiling point?
0.100 m glucose solution
0.080 m NaCl
0.030 m Al(NO3)3
The FREEZING POINT of the solvent is also effected by the presence of a solute:
WHY?
Tf = Tf (solution) - Tf (solvent)
Tf =
Which of the following water solutions have the highest freezing point?
0.100 m glucose solution
0.080 m NaCl
0.030 m Al(NO3)3