Chapter 13
Properties of Solutions
Brown, LeMay, Bursten, Murphy
AP Edition 11e
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Ways of Expressing
Concentrations of Solutions
•
•
•
•
•
•
•
% by Mass
ppm
ppb
Mole Fraction
Molarity
Molality
Normality (antiquated)
You will have to interconvert concentration units
on free-response questions.
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Mass Percentage
Mass Percent =
mass of solute
 100
total mass of solution
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Parts per ?
Parts per Million (ppm)
mass of solute
 106
ppm =
total mass of solution
Parts per Billion (ppb)
mass of solute
 109
ppb =
total mass of solution
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Mole Fraction (X)
moles of A
XA =
total moles in solution
• In some applications, one needs the
mole fraction of solvent, not solute —
make sure you find the quantity you
need!
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Molarity (M)
M=
mol of solute
L of solution
• You will recall this concentration measure
from Chapter 4.3
• Since volume is temperature-dependent,
molarity can change with temperature.
Molality (m)
m=
mol of solute
kg of solvent
• Since both moles and mass do not change with
temperature, molality (unlike molarity) is not
temperature-dependent.
• Symbol is an italicized, lower case m
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Normality (N)
gram equivalent weight of solute
Normality 
liter of solution
• It is the only concentration unit that is reaction
dependent.
• It is an antiquated term and not used in most
of the more current literature.
• Ex. 1.0M Sulfuric acid is 2 N because each
mole of acid provides 2 moles of H+ ions Solutions
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Interconverting Units
Mass
Solvent
Molality
mol/kg solvent
Moles of
Solute
Mass
Solute
Molar Mass
If we know the density of
the solution, we can
calculate the molality from
the molarity and vice versa.
Mass
Solution
This is a good chart for
interconverting values
Density
Volume
of
Solution
Molarity
mol/L solvent
Practice p.520 #25-27
Solutions
Solutions
• Solutions are homogeneous mixtures of two
or more pure substances.
• In a solution, the solute is dispersed uniformly
throughout the solvent.
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Solutions
The intermolecular
forces between solute
and solvent particles
must be strong enough
to compete with those
between solute particles
and those between
solvent particles.
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How Does a Solution Form?
As a solution forms, the solvent pulls solute
particles apart and surrounds, or solvates,
them.
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How Does a Solution Form
If an ionic salt is
soluble in water, it is
because the ion-dipole
interactions are strong
enough to overcome
the lattice energy of
the salt crystal.
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Energy Changes in Solution
• Three processes affect the
energetics of solution:
1. separation of solute
particles
2. separation of solvent
particles
3. formation of solution by
solute & solvent interaction
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Energy Changes in Solution
The enthalpy change
of the overall
process depends on
H for each of these
steps.
This is a NET
Exothermic
Process
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Why Do Endothermic Processes Occur?
Things do not tend to occur
spontaneously (i.e., without
outside intervention) unless
the energy of the system is
lowered.
This is a NET Endothermic
Process
Yet we know in some
processes, like the
dissolution of NH4NO3 (an
ice pack) in water, heat is
absorbed, not released.
WHY?
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Enthalpy Is Only Part of the Picture
The reason is that increasing
the disorder or randomness
(known as entropy, S) of a
system tends to lower the
energy of the system.
So even though enthalpy may
increase, the overall energy
of the system can still
decrease if the system
becomes more disordered.
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Student, Beware!
• When a substance disappears, is it a solution or a chemical
reaction?
• Dissolution is a physical change — you can get back the
original solute by evaporating the solvent.
If you can’t, the substance didn’t dissolve, it reacted.
Types of Solutions
• Saturated
– In a saturated solution, the
solvent holds as much solute
as is possible at that
temperature.
– When the rate of dissolving
equals the rate of
crystallization a dynamic
equilibrium is attained.
– The amount of solute
needed to form a saturated
is known as the solubility of
that solute.
dissolve

 Solution
Solute  Solvent 

crystallize
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Types of Solutions
• Unsaturated
– If a solution is
unsaturated, less
solute than can
dissolve in the
solvent at that
temperature is
dissolved in the
solvent.
dissolved

 Solution
Solute  Solvent 

crystallization
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Types of Solutions
• Supersaturated
– In supersaturated solutions, the solvent holds
more solute than is normally possible at that
temperature.
– These solutions are unstable; crystallization can
usually be stimulated by adding a “seed crystal” or
scratching the side of the flask.
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Factors Affecting Solubility
Soluble means the amount of the solute will dissolve in the solvent at a given temperature.
• Depends on
nature of solute
and solvent:
• Depends on:
– Temperature
– Pressure in gases
– IMF
– Molecular Mass
– Polarity
Solutions
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Factors Affecting Solubility
• Chemists use the axiom “like dissolves like."
– Polar substances tend to dissolve in polar solvents.
– Nonpolar substances tend to dissolve in nonpolar solvents.
Solubilities of Some Alcohols in Water and Hexane
Explanation: As the number of C increase the effect of the OH group becomes increasingly small
part of the molecule and therefore the molecule becomes less polar and immiscible in polar
solvents.
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Factors Affecting Solubility: Polarity
The more similar the
intermolecular
attractions, the more
likely one substance
is to be soluble in
another.
Ethanol-Ethanol
Ethanol-Water
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Factors Affecting Solubility: Polarity
Cyclohexane (only 1 OH
Group) is not soluble in
water
Glucose (5 OH Groups)
is very soluble in water
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Factors Affecting Solubility: Polarity
• Vitamin A is soluble in nonpolar compounds,
like fats. A, D, E, & K are fat soluble
• Vitamin C is soluble in water.
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PRACTICE EXERCISE
Arrange the following substances in order of increasing solubility in water:
Answer: C5H12 < C5H11 Cl < C5H11 OH <
Answer
C5H10(OH)2 (in order of increasing polarity and
hydrogen-bonding ability)
Solutions
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Factors Affecting Solubility:
Pressure of Gas
• In general, the solubility of gases in water increases
with increasing mass.
• Larger molecules have stronger dispersion forces.
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Factors Affecting Solubility: Pressure
• The solubility of liquids &
solids does not change
appreciably with pressure.
• The solubility of a gas in a
liquid is directly
proportional to its
pressure.
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Pressure Effects
The gas is at equilibrium in
this solution. Gas
molecule enter and
leave the solution at the
same rate.
The equilibrium is
dynamic.
Solutions
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Pressure Effects
If you increase the
pressure the gas
molecules dissolve
faster.
The equilibrium is
disturbed.
Solutions
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Pressure Effects
The system reaches a new
equilibrium with more gas
dissolved.
The relationship is:
Henry’s Law C = kP
C = Concentration of
dissolved Gas
P = Pressure
k = Henry’s Law constant
Solutions
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Henry’s Law
You can use Henry’s Law to calculate the solubility of gas
C = kP
where
• C is the solubility of the
gas,
• k is the Henry’s Law
constant for that gas in
that solvent
• P is the partial
pressure of the gas
above the liquid.
Do #43-44 p. 521
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Factors Affecting Solubility: Temperature
Generally, the solubility
of solid solutes in liquid
solvents increases with
increasing temperature.
That means you can
dissolve more sugar in
hot tea than you can in
cold tea.
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Factors Affecting Solubility: Temperature
• The opposite is true of
gases.
– Carbonated soft drinks
are more “bubbly” if
stored in the refrigerator.
– Warm lakes have less O2
dissolved in them than
cool lakes.
– Lake Nyos Tragedy
8/21/1986 2000 Deaths
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Colligative Properties
• Changes in colligative properties depend
only on the number of solute particles
present, not on the kind or identity of the
solute particles.
• Colligative properties are
– Vapor pressure lowering
– Boiling point elevation (ethylene glycol)
– Melting point depression (salt on ice)
– Osmotic pressure (kidney function)
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Vapor Pressure
•
Because of solutesolvent intermolecular
attraction, higher
concentrations of
nonvolatile solutes make
it harder for solvent to
escape to the vapor
phase.
•
Therefore, the vapor
pressure of a solution is
lower than that of the
pure solvent.
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Raoult’s Law
Psoln = XA

P
A
where
– XA is the mole fraction of compound A, and
– PA is the normal vapor pressure of A at that
temperature.
Presence of a nonvolatile solute lowers the vapor
pressure of the solvent in a closed system. See
figure 11.10
Practice #45-46 p. 521
Study Sample Exercises 11-6 & 11.7
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Boiling Point Elevation and
Freezing Point Depression
Nonvolatile solutesolvent interactions
also cause solutions
to have higher boiling
points and lower
freezing points than
the pure solvent.
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Boiling Point Elevation
• The change in boiling
point is proportional to
the molality of the
solution:
Tb = Kb  m
Tb is added to the
normal boiling point of
the solvent.
where Kb is the molal
boiling point elevation
constant, a property of
the solvent.
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Boiling Point Elevation
• The change in freezing
point can be found
similarly:
Tf = Kf  m
• Here Kf is the molal
freezing point
depression constant of
Tf is subtracted from the the solvent.
normal freezing point of
the solvent.
Boiling Point Elevation and
Freezing Point Depression
Note that in both
equations, T does
not depend on what
the solute is, but
only on how many
particles are
dissolved.
Tb = Kb  m
Tf = Kf  m
Study Ex. 11.9 & 11.10
Practice #61-64 p. 522
Osmosis
• Some substances form semipermeable
membranes, allowing some smaller
particles to pass through, but blocking
other larger particles.
• In biological systems, most
semipermeable membranes allow water
to pass through, but solutes are not free
to do so.
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Osmosis
In osmosis, there is net movement of solvent
from the area of higher solvent
concentration (lower solute concentration) to
the are of lower solvent concentration
(higher solute concentration).
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Osmotic Pressure
The pressure required to stop osmosis,
known as osmotic pressure,  (capital
pi), is
= (
n
)
RT = MRT
V
where M is the molarity of the solution.
If the osmotic pressure is the same on both sides
of a membrane (i.e., the concentrations are the
same), the solutions are isotonic.
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Osmosis in Blood Cells
• If the solute
concentration outside
the cell is greater than
that inside the cell, the
solution is hypertonic.
• Water will flow out of
the cell, and crenation
results.
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Osmosis in Cells
• If the solute
concentration outside
the cell is less than
that inside the cell, the
solution is hypotonic.
• Water will flow into the
cell, and hemolysis
results.
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Colligative Properties of
Electrolytes
Since these properties depend on the number of
particles dissolved, solutions of electrolytes (which
dissociate in solution) should show greater changes
than those of nonelectrolytes.
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Colligative Properties of
Electrolytes
However, a 1M solution of NaCl does not show
twice the change in freezing point that a 1M
solution of methanol does.
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van’t Hoff Factor
One mole of NaCl in
water does not
really give rise to
two moles of ions.
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van’t Hoff Factor
Some Na+ and Clreassociate for a short
time, so the true
concentration of
particles is somewhat
less than two times the
concentration of NaCl.
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van’t Hoff Factor
• Reassociation is
more likely at higher
concentration.
• Therefore, the
number of particles
present is
concentrationdependent.
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van’t Hoff Factor
• We modify the
previous equations
by multiplying by the
van’t Hoff factor, i.
Tf = Kf  m  i
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Colloids
Suspensions of particles larger than
individual ions or molecules, but too small to
be settled out by gravity are called colloids.
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Tyndall Effect
• Colloidal suspensions
can scatter rays of light.
• This phenomenon is
known as the Tyndall
effect.
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Colloids in Biological Systems
Some molecules have
a polar, hydrophilic
(water-loving) end and
a non-polar,
hydrophobic (waterhating) end.
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Colloids in Biological Systems
Sodium stearate
is one example
of such a
molecule.
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Colloids in Biological Systems
These molecules
can aid in the
emulsification of fats
and oils in aqueous
solutions.
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