Chapter 14:
Solutions and Their Properties
If you’re not part of the solution,
you’re part of the precipitate!
Properties of Solutions
Terminology
Solvation
Compositions
Solubility Rules
Heats of Solution
Structure/Intermolecular Forces
Henry's Law
Temperature
Vapor Pressures
Pressure
Raoult's Law
Colligative Properties
B. pt. Elevation
mS 
nS
kg A
Fr. pt. Depression
Tb , f  Kb , f  mS
PA   A  PA
A 
nA
nS  nA
Osomotic Pressure
  M S  R T
MS 
nS
V
o
Physical Properties of Solutions
Have you ever
wondered...
•Why antifreeze keeps
water from freezing?
•Why salt causes ice to
melt?
•Why cooks add salt to
boiling water?
•Why root beer foams
only when poured?
•What force opposes
gravity to allow water to
climb up a tree?
Solutions
• Solutions are homogeneous mixtures of two or
more pure substances.
• In a solution, the solute is dispersed uniformly
throughout the solvent.
An electrolyte is a substance that, when dissolved in
water, results in a solution that can conduct electricity.
A nonelectrolyte is a substance that, when dissolved,
results in a solution that does not conduct electricity.
nonelectrolyte
weak electrolyte
strong electrolyte
“like dissolves like”
Two substances with similar intermolecular forces are likely
to be soluble in each other.
•
non-polar molecules are soluble in non-polar solvents
CCl4 in C6H6
•
polar molecules are soluble in polar solvents
C2H5OH in H2O
•
ionic compounds are more soluble in polar solvents
NaCl in H2O or NH3 (l)
Energy Changes in Solution
To determine the enthalpy
change, we divide the
process into 3 steps.
1. Separation of solute
particles.
2. Separation of solvent
particles to make
‘holes’.
3. Formation of new
interactions between
solute and solvent.
Three types of interactions in the solution process:
• solvent-solvent interaction
• solute-solute interaction
• solvent-solute interaction
Hsoln = H1 + H2 + H3
Enthalpy Changes in Solution
The enthalpy
change of the
overall process
depends on H for
each of these steps.
Start
End
Start
End
Calculating Hsolution
(enthalpy of solution)
KF(s) -> K+(g) + F-(g)
 E = +821 kJ
K+(g) + F-(g) + H2O -> K+(aq) + F-(aq)
 E=-819 kJ
So Net KF(s) -> K+(aq) + F-(aq)
 Hsolution = +2kJ
Degree of saturation
• Saturated solution
 Solvent holds as much
solute as is possible at
that temperature.
 Undissolved solid
remains in flask.
 Dissolved solute is in
dynamic equilibrium
with solid solute
particles.
Degree of saturation
• Supersaturated
 Solvent holds more solute than is normally
possible at that temperature.
 These solutions are unstable; crystallization can
often be stimulated by adding a “seed crystal” or
scratching the side of the flask.
Gases in Solution
• In general, the
solubility of gases in
water increases with
increasing mass.
Why?
• Larger molecules
have stronger
dispersion forces.
Gases in Solution
Increasing
pressure
above
solution
forces
more gas
to dissolve.
• The solubility of
liquids and solids
does not change
appreciably with
pressure.
• But, the solubility of a
gas in a liquid is
directly proportional to
its pressure.
Henry’s Law
What happens to
the solubility of
carbon dioxide
in a bottle of
soda when the
pressure is
reduced?
Pressure and Solubility of Gases
The solubility of a gas in a liquid is proportional to the
pressure of the gas over the solution (Henry’s law).
c is the concentration (M) of the dissolved gas
c = kP
P is the pressure of the gas over the solution
k is a constant (mol/L•atm) that depends only
on temperature
low P
high P
low c
high c
Chemistry In Action: The Killer Lake
8/21/86
CO2 Cloud Released
1700 Casualties
Trigger?
•
earthquake
•
landslide
•
strong Winds
Lake Nyos, West Africa
Temperature
Generally, the solubility
of solid solutes in
liquid solvents
increases with
increasing
temperature.
Solubility is measured as
the mass of solute
dissolved in 100 g of
solvent at a given
temperature
Temperature
• The opposite is true of
gases. Higher
temperature drives
gases out of solution.
 Carbonated soft drinks
are more “bubbly” if
stored in the
refrigerator.
 Warm lakes have less
O2 dissolved in them
than cool lakes.
Mass Percentage
mass of A in solution
 100
Mass % of A =
total mass of solution
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!
Molarity (M)
M=
mol of solute
L of solution
• Because volume is temperature
dependent, molarity can change with
temperature.
Molality (m)
m=
mol of solute
kg of solvent
Because neither moles nor mass change
with temperature, molality (unlike molarity)
is not temperature dependent.
Changing Molarity to Molality
If we know the density
of the solution, we
can calculate the
molality from the
molarity, and vice
versa.
Colligative Properties
• 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
Vapor Pressures of Pure Water and a Water Solution
The vapor pressure of water over pure water is greater than the
vapor pressure of water over an aqueous solution containing a
nonvolatile solute.
Solute particles take up
surface area and lower
the vapor pressure
Vapor Pressure
As solute molecules are
added to a solution,
the solvent becomes
less volatile (has
decreased vapor
pressure).
Solute-solvent
interactions contribute
to this effect.
Vapor Pressure
Therefore, the vapor
pressure of a solution
is lower than that of
the pure solvent.
Colligative Properties
Lowering Vapor Pressure
• Raoult’s Law:
PA   A  P A
• Where: PA = vapor pressure with solute,
• PA = vapor pressure without solute (pure
solvent), and
• A = mole fraction of A (the pure solvent).
Boiling Point Elevation and
Freezing Point Depression
Solute-solvent
interactions also
cause solutions to
have higher boiling
points and lower
freezing points than
the pure solvent.
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.
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
freezing point of the solvent.
Colligative Properties
Boiling-Point Elevation
• Molal boiling-point-elevation constant, Kb, expresses
how much Tb changes with molality, mS :
Tb  K b  mS
• Decrease in freezing point (Tf) is directly proportional
to molality (Kf is the molal freezing-point-depression
constant):
T  K  m
f
f
S
In both equations, T does not depend on what
the solute is, but only on how many particles are
dissolved.
Colligative Properties of
Electrolytes
Because these properties depend on the number of
particles dissolved, solutions of electrolytes (which
dissociate in solution) show greater changes than those
of nonelectrolytes.
e.g. NaCl dissociates to form 2 ion particles; its limiting
van’t Hoff factor is 2.
Colligative Properties of
Electrolytes
However, a 1 M solution of NaCl does not show
twice the change in freezing point that a 1 M
solution of methanol does.
It doesn’t act like there are really 2 particles.
van’t Hoff Factor
One mole of NaCl in
water does not really
give rise to two moles
of ions.
van’t Hoff Factor
Some Na+ and Cl−
reassociate as
hydrated ion pairs, so
the true concentration
of particles is
somewhat less than
two times the
concentration of
NaCl.
The van’t Hoff Factor
• Reassociation is more
likely at higher
concentration.
• Therefore, the
number of particles
present is
concentration
dependent.
The van’t Hoff Factor
We modify the
previous equations
by multiplying by the
van’t Hoff factor, i
Tf = Kf  m  i
i = 1 for non-elecrtolytes
Osmosis
• Semipermeable membranes allow some
particles to pass through while blocking
others.
• In biological systems, most
semipermeable membranes (such as
cell walls) allow water to pass through,
but block solutes.
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).
Water tries to equalize the concentration on
both sides until pressure is too high.
Colligative Properties
Osmosis
• Osmotic pressure, , is the pressure required to stop
osmosis:
 V  n  R  T
n
     R T
V 
  M S  R T
Molar Mass from
Colligative Properties
We can use the
effects of a colligative
property such as
osmotic pressure to
determine the molar
mass of a compound.
Properties of Solutions
Terminology
Solvation
Compositions
Solubility Rules
Heats of Solution
Structure/Intermolecular Forces
Henry's Law
Temperature
Vapor Pressures
Pressure
Raoult's Law
Colligative Properties
B. pt. Elevation
mS 
nS
kg A
Fr. pt. Depression
Tb , f  Kb , f  mS
PA   A  PA
A 
nA
nS  nA
Osomotic Pressure
  M S  R T
MS 
nS
V
o