Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Solutions
Adapted by SA Green from:
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.
Solutions
• Solutions are homogeneous mixtures of two
or more pure substances.
• In a solution, the solute is dispersed uniformly
throughout the solvent.
How Does a Solution Form?
1. Solvent molecules attracted to surface ions.
2. Each ion is surrounded by solvent molecules.
3. Enthalpy (DH) changes with each interaction broken
and formed.
Ionic solid dissolving in water
How Does a Solution Form
The ions are solvated
(surrounded by
solvent).
If the solvent is water,
the ions are
hydrated.
The intermolecular
force here is iondipole.
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.
Enthalpy Changes in Solution
The enthalpy
change of the
overall process
depends on DH for
each of these steps.
Start
End
Start
End
Enthalpy changes during dissolution
DHsoln = DH1 + DH2 + DH3
The enthalpy of
solution, DHsoln, can
be either positive or
negative.
DHsoln (MgSO4)= -91.2 kJ/mol --> exothermic
DHsoln (NH4NO3)= 26.4 kJ/mol --> endothermic
Enthalpy Is Only Part of the Picture
Entropy is a measure of:
• Dispersal of energy in
the system.
• Number of microstates
(arrangements) in the
system.
b. has greater entropy,
 is the favored state
Entropy changes during dissolution
Each step also involves a
change in entropy.
1. Separation of solute
particles.
2. Separation of solvent
particles to make
„holes‟.
3. Formation of new
interactions between
solute and solvent.
Dissolution vs reaction
Ni(s) + HCl(aq)
NiCl2(aq) + H2(g)
dry
NiCl2(s)
• 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.
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
• Unsaturated Solution
 Less than the
maximum amount of
solute for that
temperature is
dissolved in the
solvent.
 No solid remains in
flask.
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.
Temperature
Generally, the
solubility of solid
solutes in liquid
solvents increases
with increasing
temperature.
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.
Factors Affecting Solubility
• Vitamin A is soluble in nonpolar compounds
(like fats).
• Vitamin C is soluble in water.
Factors Affecting Solubility
Example: ethanol in water
The stronger the
intermolecular
attractions between
solute and solvent,
the more likely the
solute will dissolve.
Ethanol = CH3CH2OH
Intermolecular forces = H-bonds; dipole-dipole; dispersion
Ions in water also have ion-dipole forces.
Factors Affecting Solubility
Glucose (which has
hydrogen bonding)
is very soluble in
water.
Cyclohexane (which
only has dispersion
forces) is not watersoluble.
Which
vitamin is
water-soluble
and which is
fat-soluble?
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
Sg = kPg
where
• Sg is the solubility of
the gas;
• k is the Henry‟s law
constant for that gas in
that solvent;
• Pg is the partial
pressure of the gas
above the liquid.
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.
Concentrations of
Solutions
Mass Percentage
mass of A in solution
 100
Mass % of A =
total mass of solution
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
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
• You will recall this concentration
measure from Chapter 4.
• 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.
Mass/Mass
Moles/Moles
Moles/L
Moles/Mass
SAMPLE EXERCISE 13.4 Calculation of Mass-Related Concentrations
(a) A solution is made by dissolving 13.5 g of glucose (C6H12O6)
in 0.100 kg of water. What is the mass percentage of solute in this
solution? (b) A 2.5-g sample of groundwater was found to
contain 5.4g of Zn2+ What is the concentration of Zn2+ in parts
per million?
PRACTICE EXERCISE
(a) Calculate the mass percentage of NaCl in a solution containing 1.50 g of NaCl in 50.0 g of water. (b) A
commercial bleaching solution contains 3.62 mass % sodium hypochlorite, NaOCl. What is the mass of NaOCl
in a bottle containing 2500 g of bleaching solution?
PRACTICE EXERCISE
A commercial bleach solution contains 3.62 mass % NaOCl in water. Calculate (a) the molality and (b) the mole
fraction of NaOCl in the solution.
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 Pressure
As solute molecules are
added to a solution,
the solvent become
less volatile
(=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.
Raoult‟s Law
PA = XAPA
where
• XA is the mole fraction of compound A
• 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.
SAMPLE EXERCISE 13.8 Calculation of Vapor-Pressure Lowering
Glycerin (C3H8O3) is a nonvolatile nonelectrolyte with a
density of 1.26 g/mL at 25 0C. Calculate the vapor
pressure at 25 0C of a solution made by adding 50.0 mL
of glycerin to 500.0 mL of water. The vapor pressure of
pure water at 25 0C is 23.8 torr (Appendix B).
PRACTICE EXERCISE
The vapor pressure of pure water at 110 0C is 1070 torr. A solution of ethylene glycol and water has a vapor
pressure of 1.00 atm at 110 0C. Assuming that Raoult’s law is obeyed, what is the mole fraction of ethylene
glycol in the solution?
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:
DTb = Kb  m
DTb 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:
DTf = Kf  m
• Here Kf is the molal
freezing point
depression constant of
the solvent.
DTf is subtracted from the normal
freezing point of the solvent.
Boiling Point Elevation and
Freezing Point Depression
In both equations,
DT does not depend
on what the solute
is, but only on how
many particles are
dissolved.
DTb = Kb  m
DTf = Kf  m
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
 DTf = 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.
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.
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.
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.
Colloids:
Suspensions of particles larger than
individual ions or molecules, but too small to
be settled out by gravity.
Tyndall Effect
• Colloidal suspensions
can scatter rays of light.
• This phenomenon is
known as the Tyndall
effect.
Colloids in Biological Systems
Sodium stearate
is one example
of such a
molecule.
Colloids in Biological Systems
These molecules
can aid in the
emulsification of fats
and oils in aqueous
solutions.