2. Solutions
Solutions
• Solutions are homogeneous mixtures of two
or more pure substances.
• In a solution, the solute is dispersed uniformly
throughout the solvent.
3. Solutions
How Does a Solution Form?
1. Solvent molecules attracted to surface ions.
2. Each ion is surrounded by solvent molecules.
3. Enthalpy (∆H) changes with each interaction broken or
formed.
Ionic solid dissolving in water
4. Solutions
How Does a Solution Form?
1. Solvent molecules attracted to surface ions.
2. Each ion is surrounded by solvent molecules.
3. Enthalpy (∆H) changes with each interaction broken or
formed.
5. Solutions
How Does a Solution Form
The ions are solvated
(surrounded by
solvent).
If the solvent is water,
the ions are
hydrated.
6. Solutions
Dissolution vs reaction
• 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.
Ni(s) + HCl(aq) NiCl2(aq) + H2(g) NiCl2(s)
dry
7. Solutions
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.
8. Solutions
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.
9. Solutions
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 scratching the side of the
flask.
10. Solutions
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.
11. Solutions
Factors Affecting Solubility
The stronger the
intermolecular
attractions between
solute and solvent,
the more likely the
solute will dissolve.Example: ethanol in water
Ethanol = CH3CH2OH
Intermolecular forces = H-bonds; dipole-dipole; dispersion
Ions in water also have ion-dipole forces.
15. Solutions
Gases in Solution
• In general, the
solubility of gases in
water increases with
increasing mass.
Why?
• Larger molecules
have stronger
dispersion forces.
16. Solutions
Gases in Solution
• 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.
Increasing
pressure
above
solution
forces
more gas
to dissolve.
17. Solutions
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.
19. Solutions
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.
22. Solutions
Parts per Million and
Parts per Billion
ppm =
mass of A in solution
total mass of solution
× 106
Parts per Million (ppm)
Parts per Billion (ppb)
ppb =
mass of A in solution
total mass of solution
× 109
23. Solutions
moles of A
total moles in solution
XA =
Mole Fraction (X)
• In some applications, one needs the
mole fraction of solvent, not solute—
make sure you find the quantity you
need!
24. Solutions
mol of solute
L of solution
M =
Molarity (M)
• Because volume is temperature
dependent, molarity can change with
temperature.
27. Solutions
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 µg of Zn2+
What is the concentration of Zn2+
in parts per million?
PRACTICE EXERCISE
Calculate the mass percentage of NaCl in a solution containing 1.50 g of NaCl in 50.0 g of water.
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.
28. Solutions
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
29. Solutions
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.
31. Solutions
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.
32. Solutions
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°C. Calculate the vapor
pressure at 25°C 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°C is 23.8 torr (Appendix B).
PRACTICE EXERCISE
The vapor pressure of pure water at 110°C is 1070 torr. A solution of ethylene glycol and water has a vapor
pressure of 1.00 atm at 110°C. Assuming that Raoult’s law is obeyed, what is the mole fraction of ethylene
glycol in the solution?
51. Solutions
Boiling Point Elevation
The change in boiling
point is proportional to
the molality of the
solution:
∆Tb = Kb m
where Kb is the molal
boiling point elevation
constant, a property of
the solvent.∆Tb is added to the normal
boiling point of the solvent.
52. Solutions
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.
53. Solutions
Boiling Point Elevation and
Freezing Point Depression
In both equations,
∆T does not depend
on what the solute
is, but only on how
many particles are
dissolved.
∆Tb = Kb m
∆Tf = Kf m
54. Solutions
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.
57. Solutions
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.
58. Solutions
The van’t Hoff Factor
• Reassociation is
more likely at higher
concentration.
• Therefore, the
number of particles
present is
concentration
dependent.
59. Solutions
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
60. Solutions
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.
61. Solutions
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.
63. Solutions
Pressures of a solution
• Osmotic pressure (the pulling pressure) of
a solution is the measure of tendency of a
solution to pull water into it by osmosis
because of the relative concentration of non
penetrating solute and water
• Hydrostatic pressure of a solution (the
pushing pressure) is the pressure exerted
by a stationary fluidic part of the solution on
an object (semi permeable membrane in case
of osmosis)
64. Solutions
Dr. Kashif Rahim 68
Osmotic Pressure
• Laws of Osmotic Pressure :
1. The Osmotic Pressure is directly
proportional to the concentration of
the solute.
2. The Osmotic Pressure is directly
proportional to the absolute
temperature.
65. Solutions
Osmotic Pressure
• The pressure required to stop osmosis,
known as osmotic pressure, π, is
n
V
π = ( )RT = MRT
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.
66. Solutions
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.
67. Solutions
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.
68. Solutions
Reverse Osmosis
• Reverse osmosis is a membrane based filtration method that removes
many types of large molecules and ions from solutions by applying
pressure to the solution when it is on one side of a selective
membrane.
• If an external pressure is applied on a concentrated solution, this
pressure is distributed evenly throughout the solution
• If the applied pressure is higher than the osmotic pressure water will
flow towards the other side of the membrane leaving solute behind
• This technique is used for purification of water
70. Solutions
Dr. Kashif Rahim 74
Reverse Osmosis is a water treatment process whereby dissolved salts, such as
sodium, chloride, calcium carbonate, and calcium sulfate may be separated from water
by forcing the water through a semi-permeable membrane under high pressure. The
water diffuses through the membrane and the dissolved salts remain behind on the
surface of the membrane.
71. Solutions
Importance of Osmosis and
Osmotic Pressure
• Oncotic pressure of blood plasma
• Formation of tissue fluid
• Regulation of cell volume
