GENERAL CHEMISTRY I
CHEM-1030
INSTRUCTOR’S LECTURE NOTES
CHANG, CHEMISTRY CHAPTER 12
Physical Properties of Solutions
Types of Solutions
Many types possible with gas, liquid or solid as the bulk solvent phase.
Examples: air, vodka, gasoline, H2 in Pt, dyed plastic,
Aqueous solutions are the commonest and most important
Although any liquid, e.g., oil, alcohol, etc. can act as a solvent.
All biochemistry and most inorganic chemistry done in water.
Solution Terms:
Unsaturated solution
Saturated solution
Supersaturated Solution
Crystallization
Precipitation (Can be due to chemical change or to exceeding solubility)
The Solution Process
Must consider solute-solute, solvent-solvent and solvent-solute forces.
“Like dissolves like”
How does water dissolve salts?
Solvation
Why does water dissolve lower alcohols?
Why doe water not dissolve the higher alcohols.
Solubility of hexane in CCl4 due to dispersion forces.
Concentration Units
Many units in use depending on type of problem or system
Percent by Mass
Expressed as %(w/w)
Units are grams solute per 100 g of solution (not just solvent)
%(w/w) can be used as conversion factor
Ex. How much NaBr is contained in 25.66 g of a 0.566 % (w/w) solution?
25.66 g solution x (0.556 g NaBr)/(100 g solution) = 0.143 g NaBr
Mole Fraction
Expressed as X
The mole fraction of any component of a solution is the number of moles of that component divided by
the total number of moles of all the components.
Ex. What is the mole fraction of ethanol in a solution made up of 0.875 mole of ethanol, 12.66 mol of
water and 0.100 mol of acetone?
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mole fraction of ethanol = mol ethanol/(mol ethanol + mol water + mol acetone)
= 0.875 mol ethanol/(0.875 mol ethanol + 12.66 mol water + 0.100 mol acetone)
= 0.875 mol/(13.64 mol) = 0.0642
Molarity
(Already studied in chapter 4)
Molarity expressed as M(mol solute/liter of solution)
Molarity is widely used because chemicals react on a mole basis
Also, it is easy to measure solution volumes in mL.
A disadvantage is that M changes with temperature as volume increases or decreases.
Molality
Expressed as m = moles solute/(1 kg solvent)
Ex. How many moles of KF are contained in 760.0 g of a 0.400 m KF solution?
760.0 g solution x (0.400 mol KF)/1 kg soln.) x (1 kg/1000 g) = 0.304 mol
Molality has the advantage that its value does not change with temperature.
Converting Between Concentration Units
Converting between M (mol/L or mol/1000 mL) and %(w/v) (g/100 mL) is easy;
(Two Examples, M → %(w/v) and %(w/v) → M.
To change between molarity (mol/L of solution) and molality (mol/kg of solvent) is harder.
#1) Must have solution density.
#2) Must convert from mass solution to mass solvent alone.
Ex: What is the molality of a 2.50 M glucose solution whose density is 1.255 g/mL?
(2.50 mol glucose/L soln) x (1 L soln/1000 mL soln) x (1 mL soln/1.255 g soln) =
0.00199 mol glucose/1 g soln = 1.99 mol glucose/1000g soln.
But, 1000 g of solution is not the same as 1000 g of solvent.
1.99 mol glucose = 358 g glucose.
The system 1.99 mol glucose/1000g of solution contains only 1000 g – 358 g = 642 g of water.
1.99 mol glucose/642. g water x (1000 g water/1 kg of water) = 3.10 mol glucose/kg of water.
Or 3.10 m glucose.
Temperature and Solubility
Solid Solubility
Most often, increasing temperature increases the solubility of solids (and liquids) in water.
(Ordinary pressure changes have no discernible effect.)
Increased thermal motion of solvent molecules keeps more solute molecules in hot solution.
Change of solubility with temperature varies widely.
Figure 12.3 Slide shows unpredictable nature.
If solubility changes rapidly with temperature, recrystallization can be used to purify an impure material
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Describe process
Fractional crystallization can also be use to separate two components of a mixture.
Gas Solubility
Gas solubility always decreases with increasing temperature
More gas molecules in a hot system will have enough energy to escape the liquid phase.
Examples: CO2 in beer, air in tap water, O2 in aquariums and in tropical seas.
Pressure and Gas Solubility
Pressure has a direct effect on the solubility of all gases.
Dynamic Equilibrium between a dissolved gas and its vapor over the solution.
Gases that do not react with water obey Henry's law relating concentration and pressure
At constant temperature, c  P
c units are mol/L, P is in atm.
c = kP
k has units mol/(L · atm). k decreases with increasing temperature.
Since k = c/P, then equation can also be used as c1/P1 = c2/P2
Qualitative examples: Opening a beer bottle, the bends.
Example: What is the concentration of argon in water at 30 ºC, given k = 1.24 x 10-4 mol/(L · atm) and
given the partial pressure of argon in air is 0.010 atm?
c = kP = 1.24 x 10-4 mol/(L · atm) x 0.010 atm = 1.2 x 10-6 mol/L
Colligative Properties of Nonelectrolyte Solutions
Colligative properties depend only on the concentration of solute particles, not on their nature.
(Colligative property equations work well only for concentrations below ~ 0.2 M.)
(Why are electrolyte solutions not included here?)
Vapor pressure Lowering
The solvent vapor pressure of solutions follows Raoult’s law.
The vapor pressure of any solution containing a nonvolatile solute is lower than the vp of the pure solvent.
In any solution, the concentration of solvent molecules at the surface is lower than for the pure solvent.
Fewer available solvent molecules always give a reduced vapor pressure.
Expressed as P1 = X1P1º
P values refer to the solvent vapor pressure and X1 refers to the mole fraction of the solvent.
Since the mole fraction of the solute (X2) is more often of interest, and X1 = 1 - X2, then:
If P1 = X1P1º, then P1 = (1 - X2) P1º
P1 = P1º - X2 P1º
P1º - P1 = X2 P1º = ΔP
Pressure change is vapor pressure of pure solvent times the solute mole fraction (whatever the solute).
Raoult's law also applies to solutions where both solvent and solute are volatile.
For chemically similar solvent and solute, Raoult's law is followed and vapor pressure vs X plot is linear.
Positive and negative deviations from Raoult's law occur when the Solvent-Solute interactions are either
weaker or stronger than Solvent-Solvent or Solute-Solute interactions.
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Slide
Stronger Solvent-Solute interactions retard the escape of molecules from the solution.
Weaker Solvent-Solute interactions increase escape rate of both kinds of molecules from solution surface.
Fractional Distillation
Separation of the volatile components of a solution through their different vapor pressures.
Examples: Liquor distillation, purification of organic synthesis mixtures.
Use of vacuum distillation to avoid temperature or air decomposition
Boiling Point Elevation
The vapor pressure of any solution is always lower than that of the pure solvent because the concentration
of solvent molecules in the solution is lower than in the pure solvent.
Therefore the boiling point of such a solution will be higher since a higher temperature is needed to reach
a vapor pressure of 1 atm.
Boiling point elevation Tb = Tb - Tbº (note order of T terms)
Tb magnitude is proportional to vapor pressure lowering and therefore to molality, m.
Tb  m
Tb = Kb m, where Kb is the boiling-point elevation constant.
Kb units are ºC/m or (ºC · kg solvent)/mol solute.
Kb values range for 0.52 for water to 2.79 for cyclohexane.
Boiling point elevations are small and not useful for determining GMW, but Kf values are.
The boiling point elevation of % ethylene glycol in water raises the b.p. of water in a car radiator.
Freezing Point Depression
Freezing point depressions larger and are easier to measure and useful for GMW determination.
Tf = Tf º - Tf (Note that Tf º  Tf so Tf is considered positive
As with boiling, Tf  m and Tf = Kf m
Dissolved solutes lower Tf because the presence of solute molecules produces disorder in the system
Qualitative examples are the lower freezing point of seawater and antifreeze.
Ex: What is the freezing point of a 50% (w/w) solution of ethylene glycol in water?
The Kf of ethylene glycol is 1.86 ºC/m. the molecular mass of ethylene glycol is 62.01 Daltons.
Tf = Kf m
Must calculate m of the 50% (w/w) solution
(500. g ethylene glycol/1 kg solution) x (1 mol ethylene glycol/62.01 g) = 8.06 mol/kg = 8.06 m
Tf = Kf m = 1.86 ºC/m x 8.06 m = 15 ºC.
Since Tf = 15ºC = Tf º - Tf, then Tf = -15ºC, sufficient to protect against moderately low temps.
Osmotic Pressure
Define Osmosis
Define Semipermeable Membrane
Osmotic Pressure, 
Net transfer of solvent in a closed system due to reduced solution vapor pressure.
The same equilibrium occurs in a system where solutions are separated by a semipermeable membrane.
The pressure needed to stop osmosis is defined as the osmotic pressure, 
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 = MRT, analogous to the ideal gas law equation PV = nRT if V is in L
How can a membrane be semipermeable to water?
Examples of raisins in Glugg, also hemolysis and crenation of RBCs.
What is the concentration of a glucose solution that has an osmotic pressure of 12.7 atm at 24.0 ºC?
 = MRT
M = /RT = 12.7 atm/(0.0821 L · atm · mol-1 · K-1 x 297.2 K) = 0.520 M
Using Colligative Properties to Determine Molecular Mass
Colligative Properties of Electrolyte Solutions
Colloids
Hydrophilic and Hydrophobic Colloids