Chapter 12
Solutions
Chapter 12 Test – Review Section
• Two AP Problems
– PV = nRT
– Calculating an Empirical/Molecular
Formula
– Graham’s Law
– Real vs. Ideal Gas
– Some other simple concepts
3.8: Solutions
• Solute: The substance that dissolves (the
minor component of a solution).
KMnO4
3.8: Solutions
• Solvent: The substance in
which the solute dissolves
(the major component of a
solution).
Solution: A homogeneous mixture
of the solute and solvent.
KMnO4 solution
3.8: Solutions
Dilute
Concentrated
Solubility
• The solubility of a solute is the quantity
that will dissolve in a given amount of
solvent to produce a saturated solution.
• Expressed as grams of solute in 100
grams of solvent, usually water.
g of solute
100 g water
Unsaturated Solutions
• Contains less than
the maximum
amount of solute.
• No undissolved
solute present.
• Can dissolve more
solute.
Dissolved
solute
Saturated Solutions
• Contains the maximum
amount of solute that can
dissolve.
• Have undissolved solute
at the bottom of the
container.
• Solution is in equilibrium
with undissolved solute
Saturation:
Factors Affecting Solubility
Temperature Effects
Temperature and Solubility of
Solids
Temperature
0°
20°C
50°C
100°C
Solubility (g/100 g H2O)
KCl(s)
NaNO3(s)
27.6
74
34.0
88
42.6
114
57.6
182
Temperature and Solubility of
Solids
Temperature
0°
20°C
50°C
100°C
Solubility (g/100 g H2O)
KCl(s)
NaNO3(s)
27.6
74
34.0
88
42.6
114
57.6
182
The solubility of most solids increases with
an increase in the temperature.
Supersaturated
• A supersaturated solution contains more than the maximum
amount of solute and is unstable. It returns to a saturated
solution if disturbed in some way (i.e.: a seed crystal)
Supersaturated
Factors Affecting Solubility
Temperature Effects
Temperature and Solubility of
Gases
Temperature
0°C
20°C
50°C
Solubility (g/100 g H2O)
CO2(g)
O2(g)
0.34
0.17
0.076
0.0070
0.0043
0.0026
Temperature and Solubility of
Gases
Temperature
0°C
20°C
50°C
Solubility (g/100 g H2O)
CO2(g)
O2(g)
0.34
0.17
0.076
0.0070
0.0043
0.0026
The solubility of gases decreases with an
increase in temperature.
Factors That Affect Solubility
• Pressure
– If we increase the
pressure on a liquid/gas
solution, we will increase
the solubility of the gas
into the liquid solvent
• Example: soda pop
Solubility and Pressure
Henry’s Law states that the solubility of a
gas in a liquid is directly related to the
pressure of that gas above the liquid.
– Doubling pressure doubles gas solubility
What affects solubility?
What affects solubility?
• For solids in liquids as the temperature
goes up the solubility goes up.
What affects solubility?
• For solids in liquids as the temperature
goes up the solubility goes up.
• For gases in a liquid as the temperature
goes up the solubility goes down.
What affects solubility?
• For solids in liquids as the temperature
goes up the solubility goes up.
• For gases in a liquid as the temperature
goes up the solubility goes down.
• For gases in a liquid- as the pressure goes
up the solubility goes up.
GENERAL PROPERTIES OF SOLUTIONS
GENERAL PROPERTIES OF SOLUTIONS
1. A solution is a homogeneous.
GENERAL PROPERTIES OF SOLUTIONS
1. A solution is a homogeneous.
2. The solute remains uniformly distributed
throughout the solution and will not settle
out over time.
GENERAL PROPERTIES OF SOLUTIONS
1. A solution is a homogeneous.
2. The solute remains uniformly distributed
throughout the solution and will not settle
out over time.
3. The solute is dissolved as molecules or
ions.
GENERAL PROPERTIES OF SOLUTIONS
1. A solution is a homogeneous.
2. The solute remains uniformly distributed
throughout the solution and will not settle
out over time.
3. The solute is dissolved as molecules or
ions.
4. It has variable composition within limits
of its solubility.
Miscibility
• “like dissolves like”
• liquids with similar structures (similar type &
magnitude of intermolecular forces) will be soluble
in each other in all proportions (miscible).
• Liquids can also be partially miscible or
immiscible.
Miscible
Disorder and Energy Changes in
the Formation of Solutions
• Spontaneous changes occur on their own
without outside assistance once conditions are
favorable for them to occur.
• Whether a substance dissolves spontaneously
is dependent upon the two driving forces.
– Lower energy
– Increased disorder (entropy)
The solution process involves changes in energy among
the attractive forces within the solute and solvent.
i
Ideal Solutions
• An ideal solution forms when the attractive
forces between the solvent and solute are
so similar that there is no measurable
energy change that occurs during solution
formation.
• Example: Mixing of ideal gases
Generally solution formation increases
the amount of disorder (entropy)
• The dissolving of salts such as NaI in
water is just one example of a change
that occurs spontaneously even though
it is endothermic. It is a universal
phenomenon that something that
brings about randomness is more likely
to occur than something that brings
about order.
Dissolving of ionic compounds
Solution Conductivity
Video: Electrolytes
Types of solutes
high conductivity
Strong Electrolyte 100% dissociation,
all ions in solution
Na+
Cl-
Strong Electrolytes
• When an ionic compound, a strong electrolyte
dissolves it dissociates completely into ions.
• Strong electrolytes can also be molecular.
These substances also dissociate completely
into ions when they dissolve, no intact molecules
remain.
• This solution conducts electricity well.
• Strong electrolytes are either strong acids,
strong bases or soluble salts.
Strong Electrolytes
→
Here, the ions in the solid are being broken apart, being
pulled away from the other ions in the solid, and are being
surrounded by water molecules. An ion-dipole attraction
occurs between the polar water molecules and the
dissociated ions.
Types of solutes
slight conductivity
Weak Electrolyte partial dissociation,
molecules and ions in
solution
CH3COOH
H+
CH3COO-
Weak Electrolytes
• A weak electrolyte is a compound that when
dissolved in water only partially ionizes or
dissociates into ions. That is, the compound
exists in water as a mixture of individual ions
and intact molecules.
• This solution conducts electricity weakly.
• Weak electrolytes include weak acids and weak
bases.
Weak Electrolytes
→
Molecules which dissolve in water are polar molecules, so there
will be dipole-dipole, or possibly even hydrogen bonding,
between the solute and the solvent. For the molecules which do
dissociate into ions, you will have ion-dipole attractions.
Types of solutes
no conductivity
Non-electrolyte No dissociation,
all molecules in
solution
sugar
Nonelectrolytes
• A nonelectrolyte is a compound that when
dissolved in water does not ionize or dissociate
into ions at all. In water, this compound exists
entirely as intact molecules.
• The solution does not conduct electricity at all.
• Nonelectrolytes would be molecular compounds
other than the acids or bases.
• Examples: alcohol, sugar.
Nonelectrolytes
→
The molecules of solute mix with the water but stay together
as molecules.
Whether a compound is an electrolyte (strong or weak),
or a nonelectrolyte does not mean soluble or insoluble.
Bronsted Acids & Bronsted Bases
• A bronsted acid is a proton (H+) doner.
• A bronsted base is a proton (H+) acceptor.
Bronsted Acids & Bronsted Bases
Percent Solution by mass
• Represents the number of grams of solute
per 100 grams of solution.
mass of solute
% solution =
X 100
mass of solution
A solution contains 130g of HCl and 750g of water.
What is the percent by mass this HCl solution?
15 %
A concentrated HCl solution has a density of
1.19 g/mL and is 37.2% HCl by mass. What mass of
HCl is needed to make exactly 1L of this solution?
443 g HCl
What volume of a concentrated HCl solution that is
37.2% HCl by mass and has a density of
1.19 g/mL is needed to obtain 125g of HCl?
Molarity (M)
• Moles of solute per liter of solution.
moles of solute
Molarity =
liters of solution
What is the molarity of a solution prepared by
dissolving 10.0 g of NaCl in 250.0 mL of solution?
0.684M
Concentrated sulfuric acid has a density of
1.84 g/ml and is 98.3% H2SO4 by mass. What
is the molarity of this solution?
18.5M
How would you prepare 500. ml of a
6.0M NaOH solution?
How would you prepare 500. ml of a
6.0M NaOH solution?
120 g of NaOH dissolved in water and
diluted to a volume of 500. ml
I need to make 250 ml of a 0.50M NaOH solution
using the 6.0M NaOH prepared in the previous
example. How do I prepare this solution?
This is a dilution problem
Dilution Problems
Mi x Vi = Mf x Vf
I need to make 250 ml of a 0.50M NaOH solution
using the 6.0M NaOH prepared in the previous
example. How do I prepare this solution?
21 ml of the 6.0M NaOH diluted to a
total volume of 250 ml.
0.850 ml of a 5.00M Cu(NO3)2 solution is
diluted to a total volume of 1.80L. What is the
molarity of the resulting diluted solution?
0.00236M
How many milliliters of water are needed to dilute
12ml of a 0.44M acetic acid solution to 0.12M?
32 ml of water added
molality (m)
• Moles of solute per kilogram of solvent.
moles of solute
molality =
kg of solvent
What is the molality of a solution made from
dissolving 20.4 g KBr in 195 g of water?
0.879m
Mole Fraction (X)
• moles of component per total moles of
solution.
moles of component
mole fraction =
total moles in solution
A gas mixture contains 45.6g CO and 899g CO2.
What is the mole fraction of CO?
0.0741
What is the mole fraction of solute in a
3.0m NaCl aqueous solution?
0.051
Colligative properties
Properties determined by the number of
solute particles in solution rather than
the type of particles.
Vapor pressure lowering
Freezing point depression
Boiling point elevation
Osmotic pressure
How Vapor Pressure Lowering
Occurs
• Solute particles take up space in a
solution.
• Solute particles on surface decrease
number of solvent particles on the
surface.
• Less solvent particles can evaporate
which lowers the vapor pressure of
a liquid.
•
•
•
•
Consider two beakers in a bell jar.
Beaker “A” contains pure water.
Beaker “B” contains saltwater.
What change in volume would you
observe in beaker “A” as time
passes?
• What change in molarity would you
observe in beaker “B” as time
passes?
• Explain your answers.
A
B
• The volume in “A” decreases?
• The molarity in “B” decreases.
• The rate of evaporation (vapor
pressure) is greater in “A” than
in “B” however the rates of
condensation are equal.
A
B
B
Raoult’s Law
Vapor Pressure Lowering =
(mole fractionsolute)(Vapor Pressuresolvent)
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.
Answer: 14.0 torr
Boiling Point Elevation
• Boiling point elevation is a colligative property related to
vapor pressure lowering.
• The boiling point is defined as the temperature at which
the vapor pressure of a liquid equals the atmospheric
pressure.
• Due to the lowering of the vapor pressure of the solvent
when a solution forms, a solution will require a higher
temperature than the pure solvent to reach its boiling
point.
Freezing Point Depression
• Every liquid has a freezing point - the temperature at
which a liquid undergoes a phase change from liquid to
solid.
• When solutes are added to a liquid, forming a solution,
the solute molecules disrupt the formation of crystals of
the solvent.
• That disruption in the freezing process results in a
depression of the freezing point for the solution relative
to the pure solvent.
Freezing-Point Depression &
Boiling-Point Elevation
Boiling Point Elevation
∆Tb = Kbm
Freezing Point Depression
∆Tf = Kfm
Kb and Kf
Calculate the freezing point of a solution that contains
5.15 g of benzene (C6H6) dissolved in 50.0 g of CCl4.
Calculate the freezing point of a solution that contains
5.15 g of benzene (C6H6) dissolved in 50.0 g of CCl4.
-63°C
Find the boiling point of a solution containing
92.1g of iodine, I2, in 800.0g of chloroform.
Find the boiling point of a solution containing
92.1g of iodine, I2, in 800.0g of chloroform.
62.9°C
Semi-permeable Membrane
• a partition that allows solvent particles to pass
through but not solute particles
Osmosis
• Osmosis: the net
movement of a
solvent through a
semipermeable
membrane toward the
solution with greater
solute concentration.
Osmosis Examples
Other everyday examples of osmosis:
• A cucumber placed in brine solution loses
water and becomes a pickle.
• A wilting flower placed in water will perk up
because water enters by osmosis.
• Eating large quantities of salty food can
cause edema (swelling in the body caused
by the accumulation of abnormally large
amounts of fluid in the spaces between the
body's cells or in the circulatory system).
Osmotic Pressure
• Osmotic Pressure (π) – the pressure required to
prevent osmosis from occurring.
Osmotic Pressure (π)
π = MRT
R values
a
The average osmotic pressure of blood is 7.7 atm at 25 C.
What concentration of glucose will be isotonic with blood?
0.31M
Molar Mass Calculations
A sample of a human hormone weighing 0.546g was
dissolved in 15.0g benzene. The freezing point of the benzene
was lowered by 0.240 C. What’s the molecular weight of the
hormone? Kf(benzene) = 5.12 °C.kg/mol
776 g/mol
A solution containing 4.5 g of glycerol, a nonvolatile
nonelectrolyte, in 100.0 g of ethanol has a boiling point of
79.0oC. If the normal BP of ethanol is 78.4oC, calculate the
molar mass of glycerol. (Kb of ethanol = 1.22 °C/m)
92 g/mol
One liter of an aqueous solution containing 20.0g of
hemoglobin has an osmotic pressure of 5.9 torr at 22°C.
What is the molecular mass of hemoglobin?
62,500 g/mol
Colligative properties
Properties determined by the number
of solute particles in solution rather than
the type of particles.
Vapor pressure lowering
Freezing point depression
Boiling point elevation
Osmotic pressure
Ionic Solute vs. Molecular Solute
• Compare the properties of a 1.0 M aqueous sugar solution to a
0.5 M solution of table salt (NaCl) in water.
• Despite the concentrations, both solutions have precisely the
same number of dissolved particles because each sodium
chloride unit creates two particles upon dissolution - a sodium
ion, Na+, and a chloride ion, Cl-.
• Therefore their colligative properties, which depend on the
number of dissolved particles, would be identical.
• Both solutions have the same freezing point, boiling point,
vapor pressure, and osmotic pressure because those colligative
properties of a solution only depend on the number of dissolved
particles.
Formula Adjustment for Ionic Solutes
• The formulas for colligative properties
must take into account a multiplier to
account for the number of dissolved
particles for ionic compounds.
– Thus NaCl and NaNO3 would have multipliers
of two, while Na2SO4 and CaCl2 would have
multipliers of three.
Formula Adjustment for Ionic Solutes
(i) = # of ions in the ionic substance
• Vapor Pressure Lowering =
(mole fractionsolute)(Vapor Pressuresolvent)(i)
• ∆Tb = Kbm(i)
• ∆Tf = Kfm(i)
• π = MRT(i)
Calculate the boiling point of a 0.40m aqueous
solution of Ca(NO3)2.
Calculate the boiling point of a 0.40m aqueous
solution of Ca(NO3)2.
100.61°C
Colligative Properties and Ionic Solutes
• Colligative effects in ionic solutions are almost
always smaller than the effect predicted on the
basis of complete dissociation due to the fact
that ions in solution to not behave completely
independently of one another, especially at
higher concentrations. This is known as the
Van’t Hoff Factor.
Ion Pairs
Ion Pair
Van’t Hoff factor
• In dilute solutions, ionic compounds have simple
multiple effects, but as the solution concentration
increases the multiplier effect diminishes.
• As concentration increases, some of the ions
floating in solution find one another and form ion
pairs, in which two oppositely charged ions
briefly stick together and act as a single particle.
• Obviously, the higher the concentration, the
more likely it is that two ions will find one
another.
For example: While we might expect NaCl to lower the freezing
point of water twice as much as an equal amount of a
nonelectrolyte, we may actually get about 1.8 times the effect.
Ion Pair
15.3
Suspensions
•Suspensions are heterogeneous because at
least two substances can be clearly identified.
Suspensions can
be separated by
filtering.
15.3
Suspensions
•Suspensions
–What is the difference between a
suspension and a solution?
15.3
Suspensions
•A suspension is a heterogeneous
mixture with dispersed particles much
larger than the solute particles in a
solution. These dispersed particles settle
out upon standing.
15.3
Colloids
•In 1996, American
astronaut Shannon Lucid
shared a gelatin dessert
with her Russian
crewmates. Gelatin is a
heterogeneous mixture
called a colloid.
15.3
Colloids
•Colloids
–What distinguishes a colloid from a
suspension and a solution?
15.3
Colloids
–Colloids are heterogeneous mixtures
that have dispersed particles smaller than
those in suspensions and larger than
those in solutions.
15.3
15.3
Colloids
–The Tyndall Effect
•The scattering of visible light by colloidal
particles is called the Tyndall effect.
15.3
Colloids
•Particles in colloids and suspensions reflect or
scatter light in all directions. Solutions do not
scatter light.
1
15.3
Brownian Motion
• In 1827 the English botanist Robert
Brown noticed that pollen grains
suspended in water jiggled about under
the lens of the microscope, following a
zigzag path. Even more remarkable was
the fact that pollen grains that had been
stored for a century moved in the same
way.
Brownian Motion
15.3
Brownian Motion
• In 1827 the English botanist Robert Brown
noticed that pollen grains suspended in water
jiggled about under the lens of the microscope,
following a zigzag path. Even more remarkable
was the fact that pollen grains that had been
stored for a century moved in the same way.
– Brownian motion is caused by collisions of the
molecules of the water with the dispersed colloidal
pollen grains.
– This offered support for the kinetic molecular theory.