Solutions Revisited
Kinetic-Molecular View of the
Solution Process
Solubilities of Solids
Effect of Temperature and Pressure
on Solubility
Molality and Mole Fraction
Raoult’s Law
Colligative Properties
Colloids
13 - 1
Physical states of solutions
Solutions can be made that exist in any of the
three states.
Solid solutions
dental fillings, 14K gold, sterling silver
Liquid solutions
saline, vodka, vinegar, sugar water
Gaseous solutions
the atmosphere, anesthesia gases
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Ideal solutions
A solution that forms without a change in
energy.
• The volume occupied by the solution is
equal to the sum of the volumes of the
components.
• The driving force for forming an ideal
solution is increased entropy.
+
Solvent
Solute
Solution
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Nonideal solutions
When a solute is added to a solvent, the
temperature of the resulting solution may
go up or down.
• There are changes in both entropy and
enthalpy.
• Volumes are not additive.
+
Solvent
or
Solute
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Predicting Solubilites
“Like dissolves like.”
Materials with similar polarity are soluble in
each other. Dissimilar ones are not.
Miscible
Liquids that are soluble in each other in all
proportions such as ethanol and water.
Immiscible
Liquids that are not soluble in each other
such as hexane and water.
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Solution of solids
When an ionic solid is placed in water, the
outer ions are exposed to the polar water
molecules.
Water will pull the ions from the solid and
surround them - solvate them.
Solvation of ions is an exothermic process
which helps overcome the lattice energy
that holds the crystal together.
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Solution of solids
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Solution of solids
While covalent compounds do not dissociate,
they are solvated in solution.
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Saturated Solutions
At saturation, the solute is in dynamic
equilibrium. The concentration is constant.
Solute species are
constantly in
motion, moving
in and out of
solution.
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Solubilities of solids
Ionic substances are not soluble in nonpolar
solvents like hexane.
• A large amount of energy is need to
separate the ions.
• A nonpolar solvent can’t solvate ions so
there is no solvation energy to offset the
lattice energy.
Predicting the solubility of ionic solids in
water is difficult because a number of
competing factors are involved.
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Solubilities of solids
Solids that exist as covalent networks are
very insoluble - glass and graphite.
Metals are also insoluble. The force that
holds them together is too strong.
Metals can only be ‘dissolved’ by chemical
reaction which converts them to soluble
compounds.
Zn(s) + 2HCl(aq)
ZnCl2(aq) + H2(g)
2Na(s) + 2H2O(l)
2NaOH(aq) + H2(g)
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Crystalline hydrates
Many compounds will crystallize from solution
with a definite proportion of water.
This water of hydration is an integral portion
of the crystal. Energy is required to remove it.
Example. CuSO . 5H O
4
2
Compounds with small, highly positive ions
such as Cu2+ and Mg2+ commonly form
hydrates.
Larger, less positive ions like K+ and Na+ do
not form hydrates.
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Temperature and solubility
Solubility
(g/100ml water)
300
SO2
KCl
glycine
NaBr
KNO3
sucrose
200
100
0
0
20
40
60
80
Temperature (oC)
100
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Pressure and solubility of gases
Increasing the pressure of a gas above a
liquid increases the concentration of the
gas.
This shifts the equilibrium, driving more gas
into the liquid.
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Pressure and solubility of gases
cg = kpgas
This law is accurate to
within 1-3% for slightly
soluble gases and
pressures up to one
atmosphere.
Solubility
(g/100g water)
Henry’s Law
At constant temperature, the solubility of a
gas is directly proportional to the pressure
of the gas above the solution.
0.010
O2
0.005
N2
He
0.000
0
1
2
Pressure (atm)
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Some new concentration units
Our discussion of the physical properties of
solutions requires the introduction of some
new concentration units.
Molality
The number of moles of solute dissolved in
one kilogram of solvent
Moles of solute
molality, m = Mass of solvent in kilograms
Unlike molarity, this unit does not change
with temperature.
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Some new concentration units
Mole fraction
The moles of solute, expressed as a fraction
of the total number of moles in the solution.
cA =
nA
nA + n B + nC + . . . .
Because the units in the numerator and
denominator are the same, mole fraction is
a unitless quantity.
The sum of all components must equal one.
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Mole fractions and partial pressure
For gases, we can relate the partial pressure
of a gas in a mixture to the mole fraction as:
pA = cA Ptotal
Example. What is the partial pressure of each gas
in a mixture of 2.43 mol N2 and 3.07 mol of O2 if the
total pressure is 26.9 atm?
c N2 =
nN2
nN2 + nO2 =
2.43 mol
2.43 mol + 3.07 mol
= 0.442
cO2 =
nO2
nN2 + nO2 =
3.07 mol
2.43 mol + 3.07 mol
= 0.558
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Mole fractions and partial pressure
pN2
= cN2 Ptotal
= 0.442 (26.9 atm)
= 11.9 atm
pO2
= cO2 Ptotal
= 0.558 (26.9 atm)
= 15.0 atm
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Raoult’s law
This law shows the vapor pressure
relationship for a volatile component in
solution and in its pure form.
PA = cA PoA
where
PA
cA
PoA
= vapor pressure of A in solution
= mole fraction of A
= vapor pressure of pure A
This relationship simply shows that as the amount
of A in a solution is reduced, its vapor pressure
will also go down. The material must be volatile.
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Distillation
The concentration dependent changes
in vapor pressure can be used to
separate mixtures of volatile solvents.
Distillation
This method relies on the the fact that
the mole fraction of the component
with the higher vapor pressure is
higher in the vapor above a liquid than
in the liquid.
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Distillation
We can determine boiling point of a mixture by
plotting of cA vs YA and cB vs YB.
BPB
vapor composition
T, oC
BPA
100%A
0%B
liquid composition
0%A
100%B
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Basic fractional
distillation equipment
Components
• Still pot with heat source
• Column
• Still head
• Receiver
The vapor is
initially produced in
the pot.
The separation
occurs in the
column.
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Distillation Equipment
Reflux condenser
Still head
Column
Valve
Receiver
Still pot
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The column
This is the heart of the distillation process.
The vapor will condense
in a series of zones called
plates. Each plate is an
equilibrium between the
gas and liquid phases.
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Colligative properties
“Bulk” properties that change when you add a
solute to make a solution.
• Based on how much you add but not
what the solute is.
• Effect of electrolytes is based on number
of ions produced.
Colligative properties
• vapor pressure lowering
• freezing point depression
• boiling point elevation
• osmotic pressure
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Vapor pressure lowering
The introduction of a nonvolatile solute will
reduce the vapor pressure of the solvent in
the resulting solution.
• The vapor pressure of a nonvolatile
component is essentially zero.
• It does not contribute to the vapor pressure
of the solution.
• However, the solution’s vapor pressure is
dependent on the solute mole fraction.
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Vapor pressure lowering
Water will end up in the ‘salt’ solution because it’s
vapor pressure is lower than the pure water.
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Boiling point elevation
When you add a nonvolatile solute to a
solvent, the boiling point goes up. This is
because the vapor pressure has been
lowered.
Dbp = Kbp x molality
The boiling point will continue to be elevated
as you add more solute until you reach
saturation.
Examples
Cooking pasta in salt water
Antifreeze
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Boiling point elevation
Example
Determine the boiling point for a 0.222 m
aqueous solution of sucrose.
Kbp = 0.512 oC m-1 for water.
Dbp = 0.512 oC m-1 (0.222 m)
= 0.114 oC
BP
= 100.00 oC + 0.114 oC = 100.11 oC
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Freezing point depression
When you add a solute to a solvent, the
freezing point goes down.
Dfp = Kfp x molality
The more you add, the lower it gets.
This will only work until you reach saturation.
Examples
“Salting” roads in winter
Making ice cream
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Freezing point depression
Example
Determine the freezing point for a 0.222 m
aqueous solution of sucrose.
Kbp = -1.86 oC m-1 for water.
Dbp = -1.86 oC m-1 (0.222 m)
= -0.413 oC
FP
= 0.00 oC - 0.413 oC = -0.41 oC
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Example constants
Kbp
Solvent
oC/m
Water
100.0
+0.512
0.0
-1.86
Benzene
80.1
+2.53
5.5
-5.12
Camphor
207.4
+5.61
178.8
-39.7
78.3
+1.22
-117.3
-1.99
Ethanol
Normal
fp, oC
Kfp
Normal
bp, oC
oC/m
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Ionic vs. covalent substances
Ionic substances have a greater effect per
mole than covalent.
• 1 mol/kg of water for glucose = 1 molal
• 1 mol/kg of water for NaCl
= 2 molal ions
• 1 mol/kg of water for CaCl2
= 3 molal ions
Effects are based on the number of particles!
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Osmosis
The movement of a solvent through a
semipermeable membrane from a dilute
solution to a more concentrated one.
Semipermeable membranes
• only allow small molecules to go through
• cell walls are semipermeable membranes
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Osmosis
13 - 36
Semipermeable membrane
ClNa+
ClNa+
ClNa+
13 - 37
Osmotic pressure
The pressure required to stop osmosis.
osmotic pressure = MRT
M = molar concentration
T = temperature in Kelvin
R = gas law constant
Since molarity is moles/liter, this equation is
just a modified form of the gas law equation.
nRT
P=
V
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Osmotic pressure
13 - 39
Osmotic Pressure
Three conditions can exist for cells.
• Concentration is the same on both sides.
isotonic
• Concentration is greater on the inside.
hypertonic cell
hypotonic solution
• Concentration is greater on the outside.
hypotonic cell
hypertonic solution
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Cell in isotonic solution
A red blood cell and
plasma have the same
osmotic pressure.
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Cells in hypertonic solution
If the level of salt in
the plasma is too high,
the cell collapses.
Crenation - water is
drawn out of the cell.
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Cells in hypotonic solution
If the level of salt
in the plasma is
too low, the cell
swells and ruptures.
Hemolysis - water is
drawn into the cell.
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Dialysis
The process where solvent and other small
molecules can pass through a membrane.
Similar to osmosis but the ‘holes’ in the
membrane are larger. As a result, even
hydrated ions can pass through.
The method relies on:
 diffusion
 osmosis
 ultrafiltration
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Dialysis
By passing large amounts of a pure solvent
past the membrane, we can flush out all but
the largest components.
pure
water in
water, ions
and small
molecule
out
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Colloids
Homogeneous mixtures of two or more
substances which are not solutions.
The substances are present as larger
particles than found in solution.
Dispersing medium - The substance in a
colloid found in the greater extent.
Dispersed phase - The substance found
in the lesser extent.
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Colloids
In colloidal suspensions, the
particles are much larger than the
solutes in a solution.
• For solutions, ions and molecules
have a size of about 10-7 cm.
• In colloids, the particles are larger,
with sizes from 10-7 to 10-5 cm.
• The colloidal particles are still too
small to settle out of solution due to
gravity.
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Tyndal effect
Unlike solutions, colloidal suspensions exhibit
light scattering.
1. purple gold sol
2. copper sulfate
solution
3. iron(III) hydroxide
colloid
1
2
3
13 - 48
Tyndal effect
13 - 49
Types of colloids
Dispersing
medium
Dispersed
phase
Name
Example
Gas
Gas
Liquid
Solid
Aerosol
Aerosol
Fog
Smoke
Liquid
Liquid
Liquid
Gas
Liquid
Solid
Foam
Emulsion
Sol
Whipped cream
Milk, mayo
Paint, ink
Solid
Solid
Solid
Gas
Liquid
Solid
Solid foam
Emulsion
Marshmallow
Butter
Pearls, opals
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Micelles
One important class of colloid is the micelle.
Molecules must have a polar and nonpolar
end to organize into this type of structure.
Examples
lipoproteins
soaps and detergents
Polar head
Nonpolar tail
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How soap works
Soaps and detergents
Work by forming
Micelles with oil.
-OOC
-OOC
COO
COO
-OOC
-OOC
-
COO
-
COO
Nonpolar tails
dissolve in oil.
-
Polar ‘heads’
are attracted
to the water.
13 - 52
Biological micelle example
cholesterol
protein
phospholipid
Lipids bound to
other molecules.
Combination results
in a micelle
structure.
13 - 53