Chapter 13 Notes- Properties of Solutions
13.1 The Solution Process [p.528]
1. Background
a. Solutions= homogeneous mixtures
i. Can be mixtures of solid, liquids, or gases
ii. Examples include saltwater, sterling silver, and air
b. Component- each substance of the mixtures
c. Solvent- component present in the greatest amount
d. Solutes – all lesser quantity components
e. Aqueous Solutions- solutions in which water is the solvent
2. 2 general factors determine the ability of substances to form solutions
a. Intermolecular interactions
b. Natural Tendency of substances to spread into larger volumes
3. The Effects of Intermolecular Forces
a. Must consider all forces involved between
i. Solute-solute
ii. Solute- solvent
iii. Solvent –solvent
b. Example is NaCL (aq)
i. Solute dissolves readily in water due to the ion-dipole attraction
ii. Solute- solvent- water molecules surround ions (solvation-hydration when water
is the solvent)
iii. Solvent- solvent hydration process the hydrogen bonding is weaken to make
space between water molecules for the ions
4. Energy Changes and Solution Formation (three aspects to consider)
a. ΔH soln = ΔH1 + ΔH2+ ΔH3
i. Solute –solute interactions must broken
1. Requires input of energy (endothermic)
2. ΔH1> 0
ii. Solvent – solvent interactions must be separated too
1. Endothermic
2. ΔH2> 0
iii. Solute – solvent interaction
1. Exothermic
2. ΔH3> 0
b. ΔH soln can be endothermic or exothermic
i. Depends on the amount of energy absorbed in separating the solute-solute
interactions and the solvent-solvent actions along with the quantity of heat
released by the formation of the solute-solvent interaction
ii. A solution will not form if ΔH soln is too endothermic. The solute-solvent be
great enough to counter the other two interactions
iii. General rule:
Polar substances dissolve in polar substances
Nonpolar in nonpolar
5. Solution Formation, Spontaneity, and Entropy
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Chapter 13 Notes- Properties of Solutions
6. Solution Formation and Chemical Reactions
a. Solutions that form spontaneously are usually processes in which the energy content if
the system decreases (exothermic)
i. Lower energy (enthalpy)
ii. Less energy required to maintain interactions
b. Principle #2 : Processes occurring at a constant temperature which result in greater
randomness or dispersal in space (entropy) are spontaneous
c. In most cases the formation of solutions is favored by the increase in entropy that
accompanies mixing
7. Solution Formation and Chemical Reactions
a. Note solutions allow for a recovering of the solute by physical means
b. Chemical reactions my result in a change of phase but the original reactants cannot be
recovered without a chemical change
13.2 Saturated Solutions and Solubility [p.534]
a. Saturated solution- contains the maximum amount of solute in a solvent based on the
temperature (solubility)
i.
Once max. amount is dissolved any extra solute will settle at the bottom of the
container
ii. Dynamic equilibrium is established between the dissolving and crystallizing of
the solute in the solution
b. Unsaturated solution contains less than the max amount
c. Supersaturated contains more than the max.
i.
Most prepare at high temperatures then cool slowly
ii. One crystal dropped on solution will cause the extra to crystallize
13.3 Factors Affecting Solubility [p.535]
1. Solute-Solvent Interactions
a. Gases larger in size and mass tend to increase in solubility to the increase in the London
dispersion forces.
b. When other factors are comparable, the stronger the attraction between solute and
solvent, the greater the solubility.
c. “Likes dissolve in likes”
d. Network solids have too great of intermolecular attractions to dissolve.
2. Pressure Effects
a. The solubility of the gas increases in direct proportion to its partial pressure above the
solution
b. Henry’s Law state the mathematical relationship
Sg = kPg
Sg= the solubility of the gas usually in molarity
k = Henry’s law constant (unique to each substance and temperature)
Pg = partial pressure of gas above the solution
3. Temperature Effects
a. As temperature of the water increases, solubility of a solid usually increases
b. Solubility of a gas decreases with an increase in water temperature
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Chapter 13 Notes- Properties of Solutions
c. Thermal pollution results from high water temperatures lowering the solubility of
oxygen
13.4 Ways of Expressing Concentrations [p.542]
1. Quantitative measurements indicate whether the solution is dilute or concentrated
2. Mass Percentage, ppm, ppb
1. Mass Percentage
mass _ solution
Mass % 
*100
total _ mass _ solution
2. ppm
mass _ solution
ppm 
*10 6
total _ mass _ solution
3. ppb
mass _ solution
ppb 
*10 9
total _ mass _ solution
3. Mole Fraction, Molarity, and Molality
1. Mole fraction (Xsub)
i. =
Moles of component
Total moles of all components
ii. No units
iii. Useful when dealing with gases
2. Molarity (M)
i. = Moles solute
Liters soln
ii. Useful for relating the volume of the solution to the amount of solute
3. Molality (m)
i. = moles of solute
Kilograms of solvent
ii. Unit = moles/ kg or m
iii. Molality does not change with temperature
4. Conversion of Concentration Units
1. May be necessary to move between types of concentrations
2. Density is frequently necessary in such conversions
13.5 Colligative Properties [p.546]
1. Colligative Properties- properties that are dependent of the concentration of the quantity of the
substance not the kind.
2. Lowering the Vapor Pressure
a. Past knowledge
i. Vapor pressure = when equilibrium is reached, the pressure of the vapor above
the liquid
ii. Involitale = substances without vapor pressure
iii. Volitale = substances with vapor pressure
b. The more parts in solution the lower the vapor pressure.
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Chapter 13 Notes- Properties of Solutions
c. Follows Raoult’s Law
PA = XAP˚A
PA= partial pressure of a solvent’ vapor above the solution
XA= mole fraction of the solvent
P˚A = vapor pressure of pure solvent
3. Boiling- Point Elevation
a. Boiling point of a solution is higher than the BP of a pure liquid
b. Equation use:
ΔTb= Kbm
ΔTb = increase in boiling point compared to pure solvent
Kb = molal boiling-elevation constant
m = Molality of the solution
c. ΔTb is proportional to the Molality of the solution
4. Freezing- Point Depression
a. Freezing point of the solution is lower than that of the pure liquid
b. Equation use:
ΔTf = Kfm
ΔTf = difference between the freezing point of eth solution and the freezing point
of the pure solvent
Kf = molal freezing-point constant
m = Molality of the solution
5. Osmosis
a. The net movement of the solvent always goes toward the higher solute concentration.
(tries to dilute a concentrated solution)
b. Osmotic pressure = the pressure needed to prevent osmosis of a pure solvent
n
    RT  MRT
V 
M= Molarity
Π = osmotic pressure
R = ideal gas constant
T = temperature
n = number of moles
V= volume
i. Isotonic = no osmosis occurs (equal osmotic pressure)
ii. Hypotonic = solution with lower osmotic pressure
iii. Hypertonic = more concentrated solution of two
iv. Crenation is the shrinking of a cell by placing a cell in a hypertonic solution
v. Hemolysis –the rupture of a cell due to placing it in a hypotonic solution
vi. IV solutions use to replenish nutrients must be isotonic to avoid cell damage
vii. Edema = retaining of water
viii. Active transport – transport from low concentration to high, not spontaneous,
energy must be expended
c. van’t Hoff, i is a factor unique to each solute and concentration
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Chapter 13 Notes- Properties of Solutions
i. represents the ratio of the ΔTf (measured) to the ΔTf ( calculated for
nonelectrolyte)
ii. value approaches the ideal value as concentration increases
6. Determination of Molar Mass- can be found by using any the colligative properties
13.6_Colloids [p.556]
1. Background
a. Colloidal dispersions or colloids – contain suspended particles that are larger than a
solute but smaller then components of a mixture
i. Represent the dividing line between solutions and heterogeneous mixtures
ii. Can be gases, liquids, or solids
b. Tyndall Effect- scattering of light by colloidal particles
2. Hydrophilic and Hydrophobic Colloids
a. Hydrophilic (water loving)- colloids that like aqueous medium keeps particles
suspended (polar groups)
b. Hydrophobic (water fearing) – must be folded inside molecule to stay suspended
c. Emulsifiers – have a hydrophilic and hydrophobic end to keep the particle suspended
3. Removal of Colloidal Particles
a. Coagulation – enlarges the size of the colloidal particles so they can be filtered out
b. Dialysis – uses a semipermeable membrane to remove the colloidal particles
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