Solutions
Chapter 12
“Everyone has problems… but
chemists have solutions”
Solution
• Seawater is a solution (a homogeneous mixture)
• The majority component is the solvent (in this
case, water) and the minority component is the
solute (here, the salt)
The solvent DOES the dissolving…
the solute gets dissolved
Spontaneity
• Typically, solutions will mix together uniformly
on their own… they are in constant motion
• Spontaneous mixing – with no barrier, a solution
of equal, uniform concentrations will result
– This explains what happens when you drink sea
water… the sea water is more concentrated than
your solutions in your digestive tracts. So, it pulls
water from your cells as it passes by in an attempt to
make a more uniform mixture.
Types of Solutions
• Solid and liquid
– Like sea water, Kool-Aid, etc.
• Gas and liquid
– Soda
• Two different liquids
• Two different gases
• Two different solids
Aqueous Solutions
• Aqueous Solutions contain water as the
solvent and a solid, liquid, or gas as the solute
•
•
•
•
Sugar dissolves in water
Salt dissociates in water
Alcohol mixes (is miscible) with water
CO2 dissolves in water to make carbonated
water (club soda)
Solubility
• Water does not dissolve grease / oil
• We say it is not soluble in water or insoluble
• The solubility of a substance is the amount of
the substance that will dissolve in a given
amount of solvent
Solubility
• For example:
– The solubility of sodium chloride in water at 25⁰C is
36 grams NaCl per 100 g of water… the solubility of
grease in water is nearly zero
• The solubility of one substance in another
depends on both:
1. Nature's tendency toward mixing
2. The types of IMFs present!!
Nature’s Tendency
• Entropy – a measure of energy randomization
or energy dispersal in a system
• If I have two gases separated by a wall:
– As soon as I open them up to each other they will
spread out to the furthest points (dispersal of
energy) becoming more chaotic
– The tendency naturally ‘mixes’ the two
components uniformly
IMFs 
• Ideal Gases do not interact with one another
and will spontaneously mix naturally
– No interactions = no IMFs
• Solids and liquids do, however, have IMFs acting
between them
– Dispersion, dipole-dipole, H-bonding, ion-dipole
• Depending on the nature of these forces, the
solute and solvent may be impacted
The Effects of IMFs
• Solvent-Solute interactions
– The interactions between a solvent particle and a
solute particle
• Solvent-Solvent interactions
– The interactions between a solvent particle and
another solvent particle
• Solute-Solute interactions
– The interactions between a solute particle and
another solute particle
The Effects of IMFs
For Example
• Consider Water and Hexane
– The water is able to form strong hydrogen bonds and
thus has a strong affinity for other water molecules.
Water and hexane cannot form comparable H-bonds.
– The energy required to pull water molecules away from
one another is too great, and too little energy is
returned when the water interacts with the hexane.
– Although the tendency to mix is strong, they cannot
overcome the large energy disparity between the
powerful solvent-solvent interactions and the weak
solvent-solute interactions
In General…
“LIKES DISSOLVE LIKES”
• Polar solvents (water) dissolve many polar or
ionic solutes
• Nonpolar solvents (hexane) tend to dissolve
nonpolar solutes
• Similar kinds of solvents dissolve similar kinds of
solutes
Vitamins
• Fat soluble vs water soluble vitamins
• Look at page 524 for a medical application to
solution (solubility)
– GREAT test “type” question 
More Practice
• Determine whether each of the following
compounds is soluble in hexane
a)
b)
c)
d)
Water
Propane
Ammonia
Hydrogen chloride
How about these four ‘in water’ ??
Dynamic Equilibrium
• The same ‘equilibrium’ discussed last unit
– They go back and forth, back and forth, …
• Consider sodium chloride (an ionic salt)
Na+(aq) + Cl-(aq)
NaCl(s)
H2O
• Eventually the rates of dissolution and deposition
become equal – dynamic equilibrium has been
reached
Some Quick Vocab 
• Saturated Solution – a solution in equilibrium (eq)
with the solute (meaning its ‘full’) – if you add more
solute it will not dissolve
• Unsaturated Solution – if you can add more… This is
a solution containing less than the (eq) amount. If
you add more, it will be able to dissolve
• Supersaturated Solution – a soln containing more
than the (eq) amount of solute (under the right
conditions). -- not usually stable and most times the
excess solute will precipitate out
Factors that Affect Solubility
• What factors do you think may have an impact
on the solubility of a solution?
Temperature Dependence
(liquid-solid)
• Solubilities of solids in water can be highly
dependent on temperature.
• Think about hot coffee and extra sugar… think
about hot tea… or making iced tea (the south
way)
**The solubility of most solids in water
increases with increasing temperature (with
some exceptions, obviously) **
Temperature for Purification
• An easy way to purify a solid is by
recrystallization
• Create a saturated solution at an elevated
temperature… meaning there is more than what
can go in at room temp.
• Then the solution is allowed to cool… the excess
solute (in this now supersaturated solution) will
‘fall out’ / precipitate out.
• The recrystallized solid is now free of impurities
Temperature Dependence
(liquid-gas)
• Liquid-gas solutions are affected differently from
liquid-solid
• Think about heating up water on the stove… but
not to the B.P.
– Bubbles start to form – these are dissolved gases
coming out of solution
In general, the solubility of gases in liquids
decreases with increasing temperature
Temperature Dependence
(liquid-gas)
In general, the solubility of gases in liquids
decreases with increasing temperature
• Think of soda.. Warm soda vs cold soda…
which one bubbles / fizzes more?
Checkpoint
• A solution is saturated in both nitrogen gas and
potassium bromide at 75⁰C. When the solution is
cooled to room temperature, which of the
following is most likely to occur?
a) Some nitrogen gas bubbles out of solution
b) Some potassium bromide precipitates out of
solution
c) Both of the above occur
d) Nothing happens…
The Effects of Pressure
• Temp effects the solubility's so you should
assume pressure will to some degree as well
• Which will be more affected?
– Liquid-solid solution
– Liquid-gas solution
• Why ??
Pressure on Gas Solutions
• The solubility of a gas increases as
the pressure increases above the
liquid
• Soda in a sealed can is under
pressure… keeping the ‘bubbles’ in
the solution
• When you open it, you release the
pressure and the gas escapes,
resulting in the bubbles you see
Henry’s Law
Sgas = kHPgas
Sgas is the solubility of the gas (usually in M)
kH is a constant of proportionality (called the
Henry’s Law constant) – depends on the specific
solute and solvent and also on the Temperature
Pgas is the partial pressure of the gas (usually in atm)
Think about the equation…
• Look at the table
on the left. Why
do you suppose
that the constant
for ammonia is
bigger than the
others???
Practice
• What pressure of carbon
dioxide is required to keep
the carbon dioxide
concentration in a bottle of
club soda at 0.12 M at
25⁰C?
= 3.5 atm
More Practice
• Determine the solubility of oxygen in water at
25⁰C exposed to air at 1.0 atm. Assume a
partial pressure for oxygen of 0.21 atm.
Concentration
• The ratio of solute to solvent is an important way
to discuss or describe a solution
• How much solute is in how much solvent…
– Dilute Solution – contains small quantities of solute
relative to the amount of solvent
– Concentrated solution – contains large quantities of
solute relative to the amount of solvent
Some Units to Describe Concentration
Molarity
Molality
Parts by mass
Parts by volume
Mole fraction
Mole percent
Molarity
• The Molarity (M) is the amount of solute in
moles per Liter of solution
– Units are mol/L or simply “M”
𝑎𝑚𝑜𝑢𝑛𝑡 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑖𝑛 𝑚𝑜𝑙)
𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝑀 =
𝑣𝑜𝑙𝑢𝑚𝑒 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝑖𝑛 𝐿)
Molarity
• Note that it is “per Liter of solution not per
Liter of solvent…
• So, you would make a solution by adding the
needed amount of solute to the flask and then
adding “enough” water to reach the desired
volume.
Molality
• Molarity uses volume… and volume depends
on the temperature.. So molarity is
temperature dependent
• A concentration unit that is independent of
temperature is molality
𝑎𝑚𝑜𝑢𝑛𝑡 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑖𝑛 𝑚𝑜𝑙)
𝑀𝑜𝑙𝑎𝑙𝑖𝑙𝑡𝑦 (𝑚) =
𝑚𝑎𝑠𝑠 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 (𝑖𝑛 𝑘𝑔)
Parts by Mass / Parts by Volume
𝑚𝑎𝑠𝑠 𝑠𝑜𝑙𝑢𝑡𝑒
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝑏𝑦 𝑚𝑎𝑠𝑠 =
∗ 100%
𝑚𝑎𝑠𝑠 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
𝑝𝑝𝑚 =
𝑚𝑎𝑠𝑠 𝑠𝑜𝑙𝑢𝑡𝑒
𝑚𝑎𝑠𝑠 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
∗ 106
𝑝𝑝𝑏 =
𝑚𝑎𝑠𝑠 𝑠𝑜𝑙𝑢𝑡𝑒
𝑚𝑎𝑠𝑠 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
∗ 109
Practice
• What volume (in mL) of a soft that is 10.5%
sucrose (C12H22O11) by mass contains 78.5 g of
sucrose? (density of the solution 1.04 g/mL)
= 719 mL
More Practice
How much sucrose, in g, is contained in 355 mL of a soft drink
that is 11.5% sucrose by mass? (density = 1.04 g/mL)
A water sample is found to contain the pollutant
chlorobenzene with a concentration of 15 ppb (by mass).
What volume of this water contains 5.00x102mg of
chlorobenzene? (density = 1.00 g/mL)
Mole Fraction
• Mole Fraction (Xsolute) is the ratio of solute to
solvent
– The amount of solute (in moles) divided by the
total amount of solute and solvent (in moles)
𝑋𝑠𝑜𝑙𝑢𝑡𝑒
𝑎𝑚𝑜𝑢𝑛𝑡 𝑠𝑜𝑙𝑢𝑡𝑒 𝑖𝑛 𝑚𝑜𝑙
=
𝑡𝑜𝑡𝑎𝑙 𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑎𝑛𝑑 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 (𝑖𝑛 𝑚𝑜𝑙)
𝑋𝑠𝑜𝑙𝑢𝑡𝑒
𝑛𝑠𝑜𝑙𝑢𝑡𝑒
=
𝑛𝑠𝑜𝑙𝑢𝑡𝑒 + 𝑛𝑠𝑜𝑙𝑣𝑒𝑛𝑡
Mole Percent
• Mole Percent (mol %) (very similar to mol
fraction)
– is the mole fraction X 100%
𝑚𝑜𝑙 % = 𝑋𝑠𝑜𝑙𝑢𝑡𝑒 ∗ 100%
• So you must know the mole fraction first!!
Practice
• A solution is prepared by dissolving 17.2 g of
ethylene glycol (C2H6O2) in 0.500 kg of water.
The final volume of the solution is 515 mL.
Calculate the:
a) Molarity
b) Molality
c) Percent by mass
d) mole fraction
e) mole percent
More Practice
• A solution is prepared by dissolving 50.4 g
sucrose in 0.332 kg of water. The final volume
of the solution is 355 mL. Calculate each of
the following:
a) Molarity
b) Molality
c) Percent by mass
d) mole fraction
e) mole percent
Vapor Pressure of Solutions
• Recall, vapor pressure of a liquid is the
pressure of the gas above the liquid when the
two are in dynamic equilibrium (meaning
closed system)
– The rate of vaporization = rate of condensation
Vapor Pressure of Solutions
• What happens to the vapor pressure when
you make a solution?
• Typically, the vapor pressure of the solution is
lower than the vapor pressure of the pure
solvent
• Colligative Properties!!
Colligative Properties!!
• Colligative Properties are properties of
solutions that depend on the ratio of the
number of solute particles to the number of
solvent molecules in a solution, and not on
the type of chemical species present.
Why??
• Why is vapor pressure lower for the solution than
the pure solvent??
• Simply put:
–
–
–
–
–
the added solute makes room for less solvent
Less solvent means less evaporating
Less evaporation means less gas component
Less gas means lower vapor pressure 
Equilibrium is again reestablished but with fewer
molecules in the gas phase
Ionic Solutes and Vapor Pressure
• When the solute is ionic, it dissociates into its
ions (meaning more “particles” in solution)
• i.e. 1 mole of MgCl2 = 3 particles in soln
• i.e. 1 mole of NaCl = 2 particles in soln
• So, this means NaCl will effect the vapor pressure
twice as much as its covalent counterpart
(magnesium chloride would have 3x the impact)
Colligative Properties
• Freezing Point Depression
• Boiling Point Elevation
• Osmosis
• These ‘manipulations’ of the properties are
caused by the “lowering of the vapor pressure
by creating a solution – adding solute particles
to the solvent”
Lowering the Vapor Pressure
• The vapor pressure curve is shifted downward
Decrease in Vapor Pressure
• This downward shift of the vapor pressure
causes a lower melting point and a higher
boiling point
• These effects are called freezing point
depression and boiling point elevation
COLLIGATIVE PROPERTIES
Freezing Point Depression
• The freezing point of a solution containing a
nonvolatile solute is lower than the freezing
point of the pure solvent
• The amount that the freezing point is lowered
for solution is given by the following equation:
Freezing Point Depression
ΔTf = m x Kf
• ΔTf = the change in T of the F.P. of soln relative to
the F.P. of the pure solvent
• m = the molality of the solution (mol/kg)
• Kf = the freezing point depression constant (for
water it is 1.87⁰C/m
Practice, Page 550 for constants
• Calculate the freezing point of a 1.7 m
aqueous ethylene glycol solution
• Calculate the freezing point of a 2.6 m
aqueous sucrose solution
Boiling Point Elevation
• The boiling point of a solution containing a
nonvolatile solute is higher than the boiling
point of the pure solvent
• The amount that the boiling point is elevated
for solution is given by the following equation:
Boiling Point Elevation
ΔTb = m x Kb
• ΔTb = the change in T of the B.P. of soln relative to
the B.P. of the pure solvent
• m = the molality of the solution (mol/kg)
• Kb = the boiling point elevation constant (for water
it is 0.512⁰C/m
Practice, Page 550 for constants
• How much ethylene glycol (C2H6O2), in grams,
must be added to 1.0 kg of water to produce a
solution that boils at 105.0⁰C?
• Calculate the boiling point of a 3.60 m aqueous
sucrose solution
Osmosis
• Osmosis is why seawater causes dehydration
– The flow of solvent from a solution of lower solute
concentration to one of higher solute
concentration
Osmotic Pressure
• Osmotic Pressure, the pressure required to stop
the osmotic flow, can be calculated:
Π = 𝑀𝑅𝑇
M is the molarity of the solution
T is the temperature (in Kelvins)
R is the ideal gas constant (0.08206 L*atm/mol*K)
Practice
• The osmotic pressure of a solution containing
5.87mg of an unknown protein per 10.0 mL of
solution was 2.45 torr at 25C. Find the molar
mass of the unknown protein.
– Come up with a game plan!!
First, use the Π (convert to atm) to find the
molarity… which is mol/L!
= 4.45x103 g/mol
More Practice
• Calculate the osmotic pressure (in atm) of a
solution containing 1.50 grams ethylene glycol
(C2H6O2) in 50.0 mL of solution at 25⁰C
Practice… Great Test Quality Q!
• The osmotic pressure of a solution containing
22.15 mg of an unknown solute per 125.0 mL
of solution was 1.55 mmHg at 125⁰C. What is
the molar mass of the unknown solute?
Review Practice
• 22.5 g of pure sucrose is dissolved in 1.50 kg of
water at 25⁰C. The final volume of the solution
is775 mL.
Calculate the:
a) Molarity
b) Molality
c) Percent by mass
d) mole fraction
e) mole percent
• Now, calculate the osmotic pressure of this
solution
van’t Hoff Factor
• Like we said earlier, colligative properties
depend only on the number of particles in
solution, not the kind of particles
• Sugar will cause F.P.D. and B.P.E. but only
about half as much of an effect as NaCl…
why??
Practice
• Which of the following solution will have the
highest boiling point?
a) 0.50 M C12H22O11
b) 0.50 M NaCl
c) 0.50 M MgCl2
van’t Hoff Factor
• This “multiplier” is known as the van’t Hoff
Factor. Though this is ‘about’ twice as much…
about three times as much… look at table
12.9 (page 554). Not whole values
• van’t Hoff Factor (i)
• They can be calculated
van’t Hoff Factors
• How can they be calculated??
ΔTf = i m x Kf
ΔTb = i m x Kb
Π = i MRT
i is the van’t Hoff factor (the multiplier)
Remember, m is molality… M is molarity
Practice: van’t Hoff Factor and
Freezing Point Depression
The freezing point of an aqueous 0.050 m CaCl2 solution
is -27⁰C. What is the van’t Hoff factor (i) for CaCl2 at this
concentration? How does it compare to the predicted
value of i (look at table 12.9)?
ΔTf = im x Kf
What is ΔTf ??
More Practice
• Compute the freezing point of an aqueous
0.10 m FeCl3 solution using the van’t Hoff
Factor of 3.2
Colloids
• Colloidal Dispersion (or a colloid) is a mixture
in which a dispersed substance (which is
solute-like) is finely divided in a dispersing
medium (which is solvent-like)
• Examples: Water and soap, fog, smoke,
whipped cream, milk…
Colloids vs Heterogeneous Mixtures
• Colloidal particles (the solute pieces) are small
enough – 1 nanometer to 1,000 nanometers
(1 μm) – to be kept dispersed throughout the
dispersing medium by collisions with other
molecules or atoms
• If the particles are bigger than a μm, they are
not uniform and will be considered
heterogeneous
Tyndall Effect
• The particles are too small to be seen but DO
STILL scatter light, they are not dissolved tho
• The scattering of light by a colloidal dispersion
is known as the Tyndall Effect
– Think of fog or dusty air
• The Tyndall effect is often used to test
whether a mixture is a solution or a colloid
Check List
 Page 560 equations
 All vocab and new terms
 Calculation problems from your list… all concentrations and
units (M, m, ppt, ppm, ppb, parts by volume, mole
fraction), Mol %
 What is vapor pressure? How does it help determine
colligative properties?
 What is freezing point depression and boiling point
elevation? How do you calculate them?
 What is osmotic pressure? How do you calculate this?
 Know the equation and be able to calculate the van’t Hoff
Factor… what is this? How do you calculate this? Where is
this applicable?