June 17, 2009 – Class 45 and 46 Overview
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13.7 Osmotic Pressure
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Osmosis, osmotic pressure calculations including molar mass
determination, colligative properties, practical applications (red
blood cells, reverse osmosis, desalination of seawater)
13.8 Freezing-Point Depression and Boiling-Point
Elevation of Nonelectrolyte Solutions
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Calculations involving molality of solute, freezing-point
depression and boiling-point elevation, and the appropriate
constants, molar mass and molecular formula determination,
practical applications.
13.9 Solutions of Electrolytes
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Van't Hoff factor and colligative properties, interionic attractions
Osmotic Pressure
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Osmosis: the net flow of solvent molecules through a
semi-permeable membrane, from a more dilute solution
(or from the pure solvent) into a more concentrated
solution.
Osmotic Pressure
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Osmotic pressure (p): the pressure that would have to
be applied to a solution to stop the passage through a
semi-permeable membrane of solvent molecules from
the pure solvent.
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Osmotic pressure for dilute solutions of nonelectrolytes
can be calculated by modifying the ideal gas equation:
pV  nRT
n
p  RT  M  RT
V
Where p = osmotic pressure, R = 0.08206 L atm mol-1 K-1, M = molarity
Osmotic Pressure
Problem: What is the osmotic pressure at 25 oC of an aqueous
solution that is 0.0010 M C12H22O11 (sucrose)?
Problem: What mass of urea [CO(NH2)2] would you dissolve in
225 mL of solution to obtain an osmotic pressure of 0.015
atm at 25 oC?
Osmotic Pressure
Problem: Creatinine is a by-product of nitrogen metabolism and
can be used to provide an indication of renal function. A 4.04
g sample of creatinine is dissolved in enough water to make
100.0 mL of solution. The osmotic pressure of the solution is
8.73 mmHg at 298 K. What is the molar mass of creatinine?
Problem: What would be the osmotic pressure of a solution
containing 2.12 g of human serum albumin (a blood plasma
protein; molar mass = 6.86 x 104 g/mol) in 75.00 mL of water
at 37.0 oC?
Osmotic Pressure – Practical Applications
Desalination of saltwater by reverse
osmosis.
The membrane is permeable to water
but not to ions.
The normal flow of water is from side A
to side B.
If we exert a pressure on side B that
exceeds the osmotic pressure of
the saltwater, a net flow of water
occurs in the reverse directionfrom the saltwater to the pure
water.
The lengths of the arrows suggest the
magnitudes of the flow of water
molecules in each direction
Osmotic Pressure – Practical Applications
Suppose an animal or a plant cell is placed in a solution of sugar or salt in water:
Hypotonic: a dilute solution, with a higher water concentration than the cell; the cell
will gain water through osmosis.
Isotonic: a solution with exactly the same water concentration as the cell; there will
be no net movement of water across the cell membrane.
Hypertonic: a concentrated solution, with a lower water concentration than the cell;
the cell will lose water by osmosis.
Colligative properties
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Colligative properties: include vapor pressure lowering,
freezing point depression, boiling point elevation and
osmotic pressure. They have values that depend on
the number of solute particles in a solution but not on
their identity.
Vapor pressure is lowered when a solute is present.
– This results in boiling point elevation and freezing
point depression.
Freezing-Point Depression and Boiling-Point
Elevation of Nonelectrolyte Solutions
Freezing-Point Depression:
DT f   K f  m
DTf = freezing-point depression (oC)
(T – Tf, where T = freezing point of the solution, Tf = freezing
point of the pure solvent)
Kf = proportionality constant (oC m-1)
(depends on melting point, enthalpy of fusion, and molar
mass of the solvent)
m = molality =
amount of solute (moles)
mass of solvent (kg)
Freezing-Point Depression and Boiling-Point
Elevation of Nonelectrolyte Solutions
Boiling-Point Elevation:
DTb  K b  m
DTb = boiling-point elevation (oC)
(T – Tb, where T = boiling point of the solution, Tb = boiling
point of the pure solvent)
Kf = proportionality constant (oC m-1)
(depends on boiling point, enthalpy of vaporization, and
molar mass of the solvent)
m = molality =
amount of solute (moles)
mass of solvent (kg)
Freezing-Point Depression and Boiling-Point
Elevation of Nonelectrolyte Solutions
Freezing-Point Depression and Boiling-Point
Elevation of Nonelectrolyte Solutions
Problem: What is the molality of nicotine in an aqueous
solution that starts to freeze at -0.450 oC? If this
solution is obtained by dissolving 1.921 g of nicotine in
48.92 g of H2O, what must be the molar mass of
nicotine?
Solutions of Electrolytes
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Van't Hoff factor (i): is a measure of the effect of a solute
upon colligative properties, ie. it was observed that certain
solutes produce a greater effect on colligative properties.
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Example: A 0.0100 m solution of NaCl in H2O has an
experimental freezing point of -0.0361 oC.
measured DT f
 0.0361o C
i

 1.94
o
1
predictedDT f
 1.86 Cm  0.0100m
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Note that this is consistent with the formation of 2 ions from the
strong electrolyte, NaCl.
For MgCl2, i = 3.
For nonelectrolytes, i = 1
For weak electrolytes i has intermediate values. For example, for
acetic acid i is greater than 1, but significantly less than 2.
Solutions of Electrolytes
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Colligative property equations should now be re-written to
account for the Van’t Hoff factor:
n
p  i  RT  i  M  RT
V
DT f   i  K f  m
DTb  i  K b  m
Freezing-Point Depression and Boiling-Point
Elevation of Nonelectrolyte Solutions
Problem: Dr. Dawe’s mom adds salt to boiling water in order
to cook spaghetti faster. How much salt needs to be
added to 8 L of water at 1 atm in order to increase its
boiling point by 1.5 oC? Assume the density of water to
be 1 kg/L.
Problem: Predict the freezing point of aqueous 0.00145 m
MgCl2.
Problem: What is the expected osmotic pressure of a
0.0530 M MgCl2 solution at 25 oC?
Interionic Attractions
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Ions in solution do not behave independently of each
other.
The difference between the
expected i and that observed
is due to the electrostatic
interactions of the ions and
results in activities or effective
concentrations
Interionic Attractions
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Ions in solution do not behave independently of each
other.