Chapter 12
Remember that a solution is any homogeneous mixture.
There are many types of solutions:
Solute
Examples
gas
Solvent Resulting
Solution
gas
gas
gas
gas
liquid
solid
liquid
solid
liquid
liquid
liquid
solid
solid
liquid
solid
liquid
solid
soda water
H2 gas in
palladium
whiskey
NaCl in water
brass
air
Some important terms to learn:
Solvent – The substance in a solution that causes the
solution to be made. Usually it is the most abundant
substance in the mixture.
Solute – The substance that is dissolved. There can be
more than one solute in a solution. In the most common
solution, the solute is a solid or a liquid and the solvent is
a liquid.
Solubility - Usually expressed as g solute/ 100 g solvent
(sometimes for aqueous solutions [any solution in which
H2O is the solvent] g solute/ 100 mL solvent)
Saturated Solution – A solution in which the maximum
amount of solute that will dissolve is present. Each
solute-solvent system is unique
Unsaturated Solution – A solution in which less than the
maximum amount of solute is dissolved.
Supersaturated Solution – Occasionally, during the
dissolving process, more than the theoretical maximum
amount of solute has dissolved. This is an unstable
situation. In time, some of the solute crystallizes out of the
solution. Frequent stirring during the solution process can
help prevent this.
Crystallization is basically the same as precipitation,
except that the processes result in solids of different
appearance. Precipitation usually results in small
particles while crystallization frequently results in larger
crystals.
In order for a solute to dissolve in a solvent, the solute
particles have to fit in between the solvent particles. This
usually can help us predict which solutes will dissolve in which
solvents. This is usually accomplished with solute-solvent
attractions. There is a general Rule of Thumb for solubility:
Like Dissolves Like. We will try to explain each case briefly
Some solutions have complete miscibility (gas in gas, many
liquids in liquids). Miscibility means complete solubility in
any proportion.
Let’s learn some more concentration terms:
1. Wt. % =
Sometimes this is written as g solute/ 100 g solution
Moles of A
2. Mole Fraction of A = XA =
total moles of all components
3. Molality (m) = The # of moles of solute dissolved in
1 kg of solvent
moles of solute
m
mass of solvent in kg
Let’s do examples 12.2 and 12.3 on pages 507-509
All the concentration terms (including molarity) can be
converted into each other.
Let’s do example 12.4 on page 510 in your text (we will also
calculate weight %).
Solubility varies with temperature. For a gas dissolved in a
liquid, the solubility always decreases as temperature
increases. For solids dissolved in liquids, there is no general
rule about the variation of solubility with temperature, except
that most often solubility increases as temperature increases.
See figure 12.3 on page 512 in your text:
External pressure has no effect on the solubilities of liquids
and solids in a liquid solvent. It does effect greatly the
solubility of gases in liquids. The effect can be calculated
using Henry’s Law, which says that the solubility of a gas
in a liquid is directly proportional to the partial pressure of
that gas over the solution.
cA = kPA , where cA is the molar concentration of gas A, PA
is the partial pressure of A over the solution and k is a
constant that depends only on T. This law is dramatically
illustrated with carbonated beverages. CO2 is dissolved in
water in a sealed container and CO2 is added to the container,
thus increasing the PCO2 and thus increasing the amount of
CO2 dissolved. When the container is opened, the CO2 above
the solution escapes, thus the solubility of the CO2 decreases
and CO2 has to escape from the solution, hence the bubbling
effect. If left open long enough all the CO2 escapes and the
beverage becomes flat.
Most gases obey Henry’s Law with some notable exceptions,
especially if the gas reacts with water. In those cases much
higher solubilities can be obtained. A good example is
ammonia in water.
There are a whole class of physical properties of solutions that
are interrelated. These are called
Colligative Properties – Depend only on the number of solute
particles in a solution but not on the nature of the solute
particles.
1. Raoult’s Law – When a non-volatile solute (basically
any solid), the partial pressure of a solvent, PS, over a
solution is directly proportional to the mole fraction of the
solvent, XS, and the VP of the pure solvent, P.
PS = XS P
The net effect is that the vapor pressure of the solution is
always lower than the vapor pressure of the pure solvent.
Raout’s Law can be rewritten as:
P = Xsolute P S0
If both components are volatile, Raoult’s Law still works. It
has to be applied for each component, solute and solvent:
PA = XAP 0A
and
PB = XBP 0B
Using Dalton’s Law of Partial Pressures
PT = PA + PB
PT = XAP
0
A
or
+ XBP
0
B
Raoult’s Law works for an ideal solution. Most solutions are
not ideal, but we will consider all solutions to be ideal.
Let’s do example 12.7 on page 517 in your text.
All the other colligative properties can be derived from
Raoult’s Law:
2. A non-volatile solute will raise the B.P. of a solution
above that of the pure solvent:
Tb = kbm,
where m is the molality and kb is a constant dependent on
the solvent, not the solute, and is called the Molal BoilingPoint Elevation Constant
Similarly, a non-volatile solute will lower the freezing
point of a solution below that of the pure solvent:
Tf = kfm,
where m is the molality and kf is a constant dependent
on the solvent, not the solute, and is called the Molal
Freezing-Point Depression Constant
Let’s do example 12.8 on page 522 of your text.
Osmosis – When 2 solutions are separated by a
semipermeable membrane (one that allows solvent
molecules to pass through but not solute particles [usually
we are dealing with small solvent molecules like water and
much larger solute particles]), solvent will flow from the
lower concentration to the higher concentration. This is
due to the attempt to achieve equilibrium (same
concentration) between the 2 solutions. This process can
be stopped if pressure is applied to the membrane from the
high concentration side. The pressure needed to stop the
flow is called the Osmotic Pressure ()
We compare the concentrations of the 2 solutions with new
terms:
hypertonic – The solution of higher concentration
hypotonic – The solution of lower concentration
isotonic - 2 solutions of equal concentration
The flow of solvent is always from hypotonic to hypertonic
until they are isotonic.
Animal cell walls are semipermeable membranes and cells
can be made to burst or shrivel if placed in hypotonic
solutions or hypertonic solutions respectively.
 = MRT
where M is the molarity (this is the only colligative
property that uses molarity), R is the Universal Gas
Law Constant and T is the Kelvin temperature. Note
that this is very similar to the Universal Gas Law
Equation. Helpful for remembering.
Colligative properties can be used to find the molecular
weights of solutes in non-electrolyte solutions. You will do this
in the last experiment in lab, making use of the Freezing Point
Depression. Boiling Point Elevation can also be used. If we
know Kf and measure the change in freezing point or boiling
point, then we can determine the molarity. If we know the
mass of solvent and unknown solute, we can then determine
the MW.
Let’s do example 12.10 and 12.11 on pages 526 and 527 in
your text.
What happens if the solute is an electrolyte? Remember that
colligative properties depend on the number of particles only.
NaCl, when dissolved in water produces Na+1 and Cl-1 ions.
Thus one molecule of NaCl produces 2 particles. A 1.0 molal
NaCl solution really has a 2.0 molal effect. Likewise, BaCl2
produces 3 particles and a 1.0 molal solution has a 3.0 molal
effect
For electrolytic solutions the colligative property equations
all have the format of:
Tf = i(kfm)
where i is the actual number of particles produced per
formula unit. This i is inserted in the same way in the
other colligative property formulas. It is called the van’t
Hoff factor. In reality the van’t Hoff factor is usually less
than expected because even a strong electrolyte has
some ions pairing up as one particle. Unless otherwise
stated, we will ignore this and simply assume complete
ionization.