Solutions

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© 2009, Prentice-Hall, Inc.
I. What is a solution?
I. What is a solution?
A. Terms
A solution is a homogeneous
mixture of two or more
substances, consisting of
ions or molecules.
Solutions
I. What is a solution?
A. Terms
1. Solute/solvent
The solute is the substance that gets
The solvent does the
.
.
Solutions
I. What is a solution?
A. Terms
1. Solute/solvent
solute
solvent
The solute is the substance that gets dissolved.
Solutions
The solvent does the dissolving.
I. What is a solution?
A. Terms
1. Solute/solvent
2. Soluble/insoluble
A substance that is
can be dissolved.
I. What is a solution?
A. Terms
1. Solute/solvent
2. Soluble/insoluble
A substance that is soluble can be dissolved.
I. What is a solution?
A. Terms
1. Solute/solvent
2. Soluble/insoluble
3. Miscible/
immiscible
Miscible substances are
will dissolve in each other.
that
Solutions
I. What is a solution?
A. Terms
1. Solute/solvent
2. Soluble/insoluble
3. Miscible/
immiscible
Miscible substances are liquids that
will dissolve in each other.
I. What is a solution?
A. Terms
1. Solute/solvent
2. Soluble/insoluble
3. Miscible/ immiscible
4. Electrolytes/nonelectrolytes
Electrolytes are liquids that are
capable of conducting electricity.
© 2009, Prentice-Hall, Inc.
© 2009, Prentice-Hall, Inc.
I. What is a solution?
B. Types of solutions
I. What is a solution?
B. Types of solutions
1. Gaseous solutions
Air and other mixtures of gases
I. What is a solution?
B. Types of solutions
1. Gaseous solutions
2. Liquid solutions
In some cases water is the solvent. These are
called (
).
I. What is a solution?
B. Types of solutions
1. Gaseous solutions
2. Liquid solutions
In some cases water is the solvent. These are
called (aqueous).
I. What is a solution?
B. Types of solutions
1. Gaseous solutions
2. Liquid solutions
3. Solid solutions
Combinations of solids. Some are
homogenous mixtures of metals (
)
Solutions
I. What is a solution?
B. Types of solutions
1. Gaseous solutions
2. Liquid solutions
3. Solid solutions
Combinations of solids. Some are
homogenous mixtures of metals (alloys)
© 2009, Prentice-Hall, Inc.
Types of Solutions
• Solutions may exist as gases, liquids,
or solids.
– The solute is the dissolved substance. In the case
of a solution of a gas or solid in a liquid, it is the
gas or solid. Otherwise, it is the component of
lesser amount.
– The solvent is the dissolving medium. Generally it
is the component of greater amount.
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Presentation of Lecture
Outlines, 12–20
Gaseous Solutions
• Nonreactive gases can mix in all
proportions to give a gaseous solution.
– Fluids that dissolve in each other in all proportions
are said to be miscible fluids.
– If two fluids do not mix, they are said to be
immiscible.
– For example, air is a solution of oxygen, nitrogen,
and smaller amounts of other gases.
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Presentation of Lecture
Outlines, 12–21
Liquid Solutions
• Liquid solutions are the most common
types of solutions found in the chemistry
lab.
– Many inorganic compounds are soluble in water or
other suitable solvents.
– Rates of chemical reactions increase when the
likelihood of molecular collisions increases.
– This increase in molecular collisions is enhanced
when molecules move freely in solution.
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Presentation of Lecture
Outlines, 12–22
Solid Solutions
• Solid solutions of metals are referred
to as alloys.
– Brass is an alloy composed of copper and zinc.
– Bronze is an alloy of copper and tin.
– Pewter is an alloy of zinc and tin.
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Presentation of Lecture
Outlines, 12–23
I. What is a solution?
C. Making a solution
I. What is a solution?
C. Making a solution
1. Solute/solvent particles
are separated.
Solutions
I. What is a solution?
C. Making a solution
1. Solute/solvent particles
are separated.
2. Solvent particles are
attracted to solute
particles.
This is called solvation.
With the solvent
this is called hydration
I. What is a solution?
C. Making a solution
1. Solute/solvent particles
are separated.
2. Solvent particles are
attracted to solute
particles.
This is called solvation.
With the solvent water this is called hydration
I. What is a solution?
C. Making a solution
1. Solute/solvent particles
are separated.
2. Solvent particles are
attracted to solute
particles.
3. ‘Like dissolves like’
Factors in Explaining Solubility
• In most cases, “like dissolves like.”
– This means that polar solvents dissolve polar (or
ionic) solutes and nonpolar solvents dissolve
nonpolar solutes.
– The relative force of attraction of the solute for
the solvent is a major factor in their solubility.
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Outlines, 12–29
Polar/Non-polar Substances
Polar
Non-polar
Water
Carbon tetrachloride
Methanol
Toluene (paint thinner)
Ethanol
Hexane, heptane, octane
Acetic acid
Benzene
Acetone
Oil
Ionic compounds
Solutions
Solutions
Solutions
Molecular Solutions
• Polar molecules interact well with polar
solvents such as water.
– The dipole-dipole interactions of water with a polar
solvent can be easily explained as electrostatic
attraction.
H
O
d-
d+
H
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polar
d- solute d+
H
O
d-
d+
H
Presentation of Lecture
Outlines, 12–33
Molecular Solutions
• Nonpolar solutes interact with nonpolar
solvents primarily due to London forces.
– Heptane, C7H16, and octane, C8H18, are both
nonpolar components of gasoline and are
completely miscible liquids.
– However, for water to mix with gasoline,
hydrogen bonds must be broken and replaced
with weaker London forces between water and
the gasoline.
– Therefore gasoline and water are nearly
immiscible.
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Outlines, 12–34
Ionic Solutions
• Polar solvents, such as water, also
interact well with ionic solutes.
– Since ionic compounds are the extreme in polarity, we
can illustrate the electrostatic attractions of water for
cations and anions.
H
d+
H
O d- + dO-
H
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d+
H
H
O
d-
d+
H
H
-
d+
O d-
H
Presentation of Lecture
Outlines, 12–35
I. What is a solution?
D. Dissolving Ionic Compounds
in H2O
1. Water is polar because it has
two oppositely
ends
I. What is a solution?
D. Dissolving Ionic Compounds
in H2O
1. Water is polar because it has
two oppositely charged ends
I. What is a solution?
D. Dissolving Ionic Compounds
in H2O
1. Water is polar because it has
two oppositely charged ends
2. The
end of water is
attracted to the anion
Anion = negative ion
I. What is a solution?
D. Dissolving Ionic Compounds
in H2O
1. Water is polar because it has
two oppositely charged ends
2. The positive end of water is
attracted to the anion
Anion = negative ion
I. What is a solution?
D. Dissolving Ionic Compounds
in H2O
3. The combined
of water
molecules overcomes the
attractive force of solute.
I. What is a solution?
D. Dissolving Ionic Compounds
in H2O
3. The combined pull of water
molecules overcomes the
attractive force of solute.
I. What is a solution?
D. Dissolving Ionic Compounds
in H2O
3. The combined pull of water
molecules overcomes the
attractive force of solute.
4. The ionic compound is
dissociated into individual
ions.
Ex.
NaCl (s)
Na+ + Cl-
I. What is a solution?
E. Dissolving Molecular
Compounds in H2O.
I. What is a solution?
E. Dissolving Molecular
Compounds in H2O.
1. Water’s charged ends are
attracted to charged
portions of solute.
Solutions
I. What is a solution?
E. Dissolving Molecular
Compounds in H2O.
1. Water’s charged ends are
attracted to charged
portions of solute.
2. The solute is surrounded by
water molecules.
I. What is a solution?
E. Dissolving Molecular
Compounds in H2O.
1. Water’s charged ends are
attracted to charged
portions of solute.
2. The solute is surrounded by
water molecules.
3. The solute remains intact.
I. What is a solution?
F. Rate of Solvation
To speed up solvation. . .
I. What is a solution?
F. Rate of Solvation
To speed up solvation. . .
1. Stir/shake
I. What is a solution?
F. Rate of Solvation
To speed up solvation. . .
1. Stir/shake
2. Break apart solute
I. What is a solution?
F. Rate of Solvation
To speed up solvation. . .
1. Stir/shake
2. Break apart solute
3. Heat the solvent
II. Is it soluble?
II. Is it soluble?
A. Not all ionic compounds
dissolve in water.
II. Is it soluble?
A. Not all ionic compounds
dissolve in water.
1. Solubility Table
II. Is it soluble?
A. Not all ionic compounds
dissolve in water.
1. Solubility Table
2. Writing physical states
(in equations)
Insoluble = (s) for solid
Soluble = ______________
II. Is it soluble?
A. Not all ionic compounds
dissolve in water.
1. Solubility Table
2. Writing physical states
(in equations)
Insoluble = (s) for solid
Soluble = (aq) for aqueous
II. Is it soluble?
A. Not all ionic compounds
dissolve in water.
1. Solubility Table
2. Writing physical states
(in equations)
3. Solubility Rules
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Self Check – Ex. 1
Is copper (II) chloride
soluble?
Self Check – Ex. 2
Is zinc carbonate soluble?
Self Check – Ex. 3
Is calcium hydroxide
soluble?
II. Is it soluble?
A. Not all ionic compounds
dissolve in water.
B. In reality, ionic compounds
all dissolve in varying degrees.
Some compounds are very soluble while
others are slightly soluble
II. Is it soluble?
A. Not all ionic compounds
dissolve in water.
B. In reality, ionic compounds
all dissolve in varying degrees.
1. Solubility depends on
___________
Solubility
Curve
II. Is it soluble?
A. Not all ionic compounds
dissolve in water.
B. In reality, ionic compounds
all dissolve in varying degrees.
1. Solubility depends on
temperature
2. Solubility depends on
physical state
II. Is it soluble?
A. Not all ionic compounds
dissolve in water.
B. In reality, ionic compounds
all dissolve in varying degrees.
1. Solubility depends on
temperature
2. Solubility depends on
physical state
3. Solubility of gas depends
on pressure
III. Heat of solution
III. Heat of solution
A. Many reactions warm up when
you mix substances
Exothermic – heat given off
III. Heat of solution
A. Many reactions warm up
when you mix substances
B. Some reactions get colder
when you mix substances
Endothermic – heat absorbed
IV. Concentration
IV. Concentration
A. Terms
IV. Concentration
A. Terms
1. Concentrated
A solution that has a lot of
per solvent
IV. Concentration
A. Terms
1. Concentrated
A solution that has a lot of solute per solvent
IV. Concentration
A. Terms
1. Concentrated
2. Dilute
A solution that has a lot of
.
IV. Concentration
A. Terms
1. Concentrated
2. Dilute
A solution that has a lot of solvent.
IV. Concentration
A. Terms
1. Concentrated
2. Dilute
3. Saturated
A solution with
solute.
IV. Concentration
A. Terms
1. Concentrated
2. Dilute
3. Saturated
A solution with the maximum amount of solute.
IV. Concentration
A. Terms
1. Concentrated
2. Dilute
3. Saturated
4. Supersaturated
A solution with
.
IV. Concentration
A. Terms
1. Concentrated
2. Dilute
3. Saturated
4. Supersaturated
A solution with more solute than it can hold.
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Solubility and the Solution
Process
• The amount of a substance that will
dissolve in a solvent is referred to as its
solubility.
– Many factors affect solubility, such as
temperature and, in some cases, pressure.
– There is a limit as to how much of a given solute
will dissolve at a given temperature.
– A saturated solution is one holding as much
solute as is allowed at a stated temperature.
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Outlines, 12–80
Solubility: Saturated Solutions
• Sometimes it is possible to obtain a
supersaturated solution, that is, one
that contains more solute than is
allowed at a given temperature.
– Supersaturated solutions are unstable.
– If a small crystal of the solute is added to a
supersaturated solution, the excess immediately
crystallizes out.
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Presentation of Lecture
Outlines, 12–81
IV. Concentration
A. Terms
B. Measurements
1. Percentage
2. Molarity
3. Molality
4. Mole fraction
Colligative Properties
• Consider three beakers:
– 50.0 g of ice
– 50.0 g of ice + 0.15 moles NaCl
– 50.0 g of ice + 0.15 moles sugar (sucrose)
• What will the temperature of each
beaker be?
– Beaker 1:
– Beaker 2:
– Beaker 3:
Colligative Properties
• The reduction of the freezing point
of a substance is an example of a
colligative property:
–
A property of a solvent that depends
on the total number of solute particles
present
• There are four colligative properties
to consider:
–
–
–
–
Vapor pressure lowering (Raoult’s
Law)
Freezing point depression
Boiling point elevation
Osmotic pressure
Colligative: particles are particles
• Colligative comes from colligate – to tie
together
• Colligative properties have common origin
• Colligative properties depend on amount of
solute but do not depend on its chemical
identity
• Solute particles exert their effect merely by
being rather than doing
• The effect is the same for all solutes
Colligative properties for
nonvolatile solutes:
•
•
•
•
Vapour pressure is always lower
Boiling point is always higher
Freezing point is always lower
Osmotic pressure drives solvent from
lower concentration to higher
concentration
How Does a Solution Form
If an ionic salt is
soluble in water, it is
because the iondipole interactions
are strong enough
to overcome the
lattice energy of the
salt crystal.
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Physical vs Chemical
• Mixing is physical process; chemical
properties don’t change
• Properties of solutions are similar to
those of the pure substances
• Addition of a foreign substance to water
alters the properties slightly
Energy Changes in Solution
• Simply put, three
processes affect the
energetics of solution:
– separation of solute
particles,
– separation of solvent
particles,
– new interactions
between solute and
solvent.
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Energy Changes in Solution
The enthalpy
change of the
overall process
depends on H for
each of these steps.
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Endothermic Processes?
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.
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Factors Affecting Solubility
The more similar the
intermolecular
attractions, the more
likely one substance
is to be soluble in
another.
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Factors Affecting Solubility
Glucose (which has
hydrogen bonding)
is very soluble in
water, while
cyclohexane (which
only has dispersion
forces) is not.
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Factors Affecting Solubility
• Vitamin A is soluble in nonpolar compounds
(like fats).
• Vitamin C is soluble in water.
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Gases in Solution
• The solubility of
liquids and 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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Henry’s Law
Sg = kPg
where
• Sg is the solubility of
the gas,
• k is the Henry’s Law
constant for that gas in
that solvent, and
• Pg is the partial
pressure of the gas
above the liquid.
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Temperature
Generally, the
solubility of solid
solutes in liquid
solvents increases
with increasing
temperature.
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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.
© 2009, Prentice-Hall, Inc.
Colligative Properties
• Changes in colligative properties
depend only on the number of solute
particles present, not on the identity of
the solute particles.
• Among colligative properties are
– Vapor pressure lowering
– Boiling point elevation
– Melting point depression
– Osmotic pressure
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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.
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Colligative Properties – Vapor
Pressure
• A solvent in a closed container reaches a
state of dynamic equilibrium.
• The pressure exerted by the vapor in the
headspace is referred to as the vapor
pressure of the solvent.
• The addition of any nonvolatile solute
(one with no measurable vapor pressure)
to any solvent reduces the vapor
pressure of the solvent.
Colligative Properties – Vapor
Pressure
• Nonvolatile solutes reduce the ability
of the surface solvent molecules to
escape the liquid.
– Vapor pressure is reduced.
• The extent of vapor pressure
lowering depends on the amount of
solute.
– Raoult’s Law quantifies the amount of
vapor pressure lowering that is observed.
Non-volatile solutes and Raoult’s
law
• Vapor pressure of solvent in solution containing nonvolatile solute is always lower than vapor pressure of
pure solvent at same T
– At equilibrium rate of vaporization = rate of condensation
– Solute particles occupy volume reducing rate of evaporationthe
number of solvent molecules at the surface
– The rate of evaporation decreases and so the vapor pressure
above the solution must decrease to recover the equilibrium
Raoult’s Law
PA = XAPA
where
– XA is the mole fraction of compound A, and
– PA is the normal vapor pressure of A at
that temperature.
NOTE: This is one of those times when you
want to make sure you have the vapor
pressure of the solvent.
© 2009, Prentice-Hall, Inc.
Colligative Properties – Vapor
Pressure
Example: The vapor pressure of water at
20oC is 17.5 torr. Calculate the vapor
pressure of an aqueous solution prepared
by adding 36.0 g of glucose (C6H12O6) to
14.4 g of water.
Colligative Properties – Vapor
Pressure
Answer: 14.0 torr
Colligative Properties – Vapor
Pressure
Example: The vapor pressure of pure
water at 110oC is 1070 torr. A solution of
ethylene glycol and water has a vapor
pressure of 1.10 atm at the same
temperature. What is the mole fraction
of ethylene glycol in the solution?
Both ethylene glycol and water are liquids. How
do you know which one is the solvent and which
one is the solute?
Colligative Properties – Vapor
Pressure
Answer: XEG = 0.20
Colligative Properties – Vapor
Pressure
• Ideal solutions are those that obey
Raoult’s Law.
• Real solutions show approximately
ideal behavior when:
– The solution concentration is low
– The solute and solvent have similarly
sized molecules
– The solute and solvent have similar types
of intermolecular forces.
Colligative Properties – Vapor
Pressure
• Raoult’s Law breaks down when
solvent-solvent and solute-solute
intermolecular forces of attraction
are much stronger or weaker than
solute-solvent intermolecular forces.
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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Colligative Properties – BP
Elevation
• The addition of a
nonvolatile solute
causes solutions to
have higher boiling
points than the
pure solvent.

– Vapor pressure
decreases with
addition of non-
Higher temperature is
needed in order for vapor
pressure to equal 1 atm.
Molecular view of Raoult’s law:
Boiling point elevation
• In solution vapor pressure is reduced
compared to pure solvent
• Liquid boils when vapor pressure =
atmospheric pressure
• Must increase T to make vapor
pressure = atmospheric
Freezing Point Depression
• The change in freezing
point can be found
similarly:
Tf = Kf  m
• Here Kf is the molal
freezing point
depression constant of
the solvent.
Tf is subtracted from the normal
boiling point of the solvent.
© 2009, Prentice-Hall, Inc.
Colligative Properties - Freezing Pt Depression
• The addition of a
nonvolatile solute causes
solutions to have lower
freezing points than the
pure solvent.
• Solid-liquid equilibrium line
rises ~ vertically from the
triple point, which is lower
than that of pure solvent.

Freezing point of
the solution is lower
than that of the
pure solvent.
Molecular view of Raoult’s law:
Freezing point depression
• Depends on the solute only being in the liquid phase
–
–
–
–
Fewer water molecules at surface: rate of freezing drops
Ice turns into liquid
Lower temperature to regain balance
Depression of freezing point
Colligative Properties - Freezing Pt Depression
• Reminder: You should be able to do
the following as well:
– Calculate the freezing point of any
solution given enough information to
calculate the molality of the solution and
the value of Kf
– Calculate the molar mass of a solution
given the value of Kf and the freezing
point depression (or the freezing points
of the solution and the pure solvent).
Deviations from ideal
• Real solutions can deviate from the
ideal:
– Positive (Pvap > ideal) solute-solvent
interactions weaker
– Negative (Pvap < ideal) solute-solvent
interactions stronger
VI. Colligative Properties
A.
B.
C.
D.
Vapor Pressure: it’s lowered
Boiling Point: it’s elevated
Freezing Point: it’s lowered
Osmotic Pressure: it’s lowered
Colligative Properties - 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.
• If two solutions with identical
concentration (isotonic solutions) are
separated by a semipermeable membrane,
no net movement of solvent occurs.
Colligative Properties - Osmosis
• Osmosis: the net movement of a
solvent through a semipermeable
membrane toward the solution with
greater solute concentration.
• In osmosis, there is net movement of
solvent from the area of lower solute
concentration to the area of higher
solute concentration.
– Movement of solvent from high
solvent concentration to low solvent
concentration
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).
© 2009, Prentice-Hall, Inc.
Osmosis: molecular
discrimination
• A semi-permeable membrane
discriminates on the basis of molecular
type
– Solvent molecules pass through
– Large molecules or ions are blocked
• Solvent molecules will pass from a
place of lower solute concentration to
higher concentration to achieve
equilibrium
Osmotic pressure
• Solvent passes into more conc solution
increasing its volume
• The passage of the solvent can be
prevented by application of a pressure
• The pressure to prevent transport is the
osmotic pressure
Calculating osmotic pressure
• The ideal gas law states
PV  nRT
• But n/V = M and so
  MRT
• Where M is the molar concentration of
particles and Π is the osmotic pressure
• Note: molarity is used not molality
Colligative Properties Osmosis
• Osmosis plays an important role
in living systems:
– Membranes of red blood cells
are semipermeable.
• Placing a red blood cell in a
hypertonic solution (solute
concentration outside the cell is
greater than inside the cell)
causes water to flow out of the
cell in a process called
crenation.
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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Colligative Properties
• Placing a red blood cell in a hypotonic
solution (solute concentration outside
the cell is less than that inside the
cell) causes water to flow into the
cell.
– The cell ruptures in a process called
hemolysis.
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.
© 2009, Prentice-Hall, Inc.
© 2009, Prentice-Hall, Inc.
© 2009, Prentice-Hall, Inc.
Colligative Properties - Osmosis
• Other everyday examples of osmosis:
– A cucumber placed in brine solution loses
water and becomes a pickle.
– A limp carrot placed in water becomes
firm because water enters by osmosis.
– Eating large quantities of salty food
causes retention of water and swelling of
tissues (edema).
Osmotic pressure and molecular
mass
• Molar mass can be computed from any
of the colligative properties
• Osmotic pressure provides the most
accurate determination because of the
magnitude of Π
– 0.0200 M solution of glucose exerts an
osmotic pressure of 374.2 mm Hg but a
freezing point depression of only 0.02ºC
Osmotic Pressure
The pressure required to stop osmosis,
known as osmotic pressure, , 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.
© 2009, Prentice-Hall, Inc.
Determining molar mass
• A solution contains 20.0 mg insulin in 5.00
ml develops an osmotic pressure of 12.5
mm Hg at 300 K

M
12.5mmHg 1
RT
760mmHg
M
 6.68 10 - 4 M
L  atm
0.0821
300 K
mol  K
Mass Percentage
mass of A in solution
 100
Mass % of A =
total mass of solution
© 2009, Prentice-Hall, Inc.
Parts per Million and
Parts per Billion
Parts per Million (ppm)
mass of A in solution
 106
ppm =
total mass of solution
Parts per Billion (ppb)
mass of A in solution
 109
ppb =
total mass of solution
© 2009, Prentice-Hall, Inc.
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!
© 2009, Prentice-Hall, Inc.
Molarity (M)
M=
mol of solute
L of solution
• You will recall this concentration
measure from Chapter 4.
• Since volume is temperaturedependent, molarity can change with
temperature.
© 2009, Prentice-Hall, Inc.
Molality (m)
m=
mol of solute
kg of solvent
Since both moles and mass do not
change with temperature, molality
(unlike molarity) is not temperaturedependent.
© 2009, Prentice-Hall, Inc.
Changing Molarity to Molality
If we know the
density of the
solution, we can
calculate the
molality from the
molarity and vice
versa.
© 2009, Prentice-Hall, Inc.
© 2009, Prentice-Hall, Inc.
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