Colligative Properties
Now let's go back to a question asked at the very beginning of this lesson,
and that is - how are solutions different than pure liquids? One of the ways in
which they are different, is that when you add a solute to a
liquid both the freezing point and boiling point of the solution change.
Water, of course, is the liquid we will be dealing with. The freezing point of
pure water is 0°C. The normal boiling point of water is 100°C. But if you make
a solution using water as the solvent, the freezing point of that solution will not
be 0°C nor will the boiling point be 100°C.
In addition the vapor pressure of the liquid changes. Also, something called
the osmotic pressure of the liquid changes and that is related to the process
of osmosis. In this section we will look at each of these in turn.
The freezing point depression, boiling point elevation, vapor pressure
lowering, and osmotic pressure are all related to one another, because the
magnitude of the change depends on the concentration of solute particles. It
is also dependent on the nature of the solvent. These properties are not so
much dependent on the nature of, or the chemical properties of the solute that
is dissolved, but simply on the number of solute particles present, whether
they are ions or molecules doesn't make too much difference except in the
number. It is the concentration that make the difference, not the nature. Of
course, that concentration does have to take into account whether that solute
dissociates and if so, how much. Because of this they are all grouped together
as a set of properties, and they are called the colligative properties.
Freezing Point Depression | Boiling Point Elevation |
Vapor Pressure Change | Osmosis
Freezing Point Depression
Freezing point depression is not just another way of referring to the early
February blues. Instead, it has to do with the change that occurs in the
temperature at which a liquid freezes (or a solid melts) when a solute is
dissolved in it.
For example, consider the effect that salt has on the melting point of ice
(which is the same as the freezing point of water). In the Midwest, salting of
roads is very common during the winter; it melts the ice and snow that's
present on the roads. Here in the Northwest, sanding is more common than
salting for roads but some people will use salt on sidewalks. If you've ever
made ice cream in a cranked churn, you probably layered ice and rock salt to
lower the freezing point enough to freeze the ice cream mixture in the
canister.
Another example of this is the use of antifreeze in car radiators. By using a
solution instead of pure water in the radiator, the liquid will not freeze until you
get to some temperature below 0°C (which is 32°F), rather than freezing right
at 0°C.
The next time you see a
container of antifreeze, look
on the label and it will show
you that as you increase the
amount (concentration,
actually) of the antifreeze in
the car's cooling system the
freezing point of the solution
decreases. There is a limit to
that, and the instructions
usually say that you should
not exceed a certain
percentage. If you get to the
point where you have water
in the antifreeze rather than
antifreeze in the water, the
freezing point starts going
back up--although it is still
very low.
In general, for dilute solutions, the amount of change in the freezing point is
proportional to the concentration of the solute in the solution.
Boiling Point Elevation
The effect on the boiling point is just the opposite; that is, the boiling point of
a liquid is increased if something is dissolved in it.
Again, the antifreeze in a car
radiator is an example of
this. In the advertisements
it's called summer protection
against boil-overs.
Essentially what they have
done is raised the boiling
point of the liquid, of the
water, by making it a
solution.
One way of summarizing both
freezing point depression and
boiling point elevation is to
say that the addition of solute
extends the temperature
range over which the liquid
can exist. It will boil at a
higher temperature and it will
freeze at a lower
temperature. Note that as the
concentration of the solute
increases, so does the effect
on both the freezing point and
the boiling point.
A very practical use of this phenomenon was mentioned earlier - antifreeze.
Almost anything can be used to lower the freezing point of water, or to raise
the boiling point, as long as it dissolves in water. But there are some very
important, practical limitations. For example, if you have a very volatile
material such as ethanol, it would not work well in a car's radiator, because as
the engine got hot the ethanol could vaporize and leave; it would no longer be
dissolved in the solution, therefore it would not be able to lower the freezing
point of the solution. On the other hand an ionic material would work but it
would probably enhance the corrosion of the interior of the car's engine. You
don't want that to happen either. So you have to be very careful about the
other physical and chemical properties of the material that you choose when
you want to lower the freezing point or raise the boiling point for some
particular purpose.
In general, the extent of the boiling or freezing point change is
proportional to the number of moles of solute added to a certain amount
of water, regardless of what specific chemical is used.
Electrolyte Effect
However, you have to take into account the degree of dissociation of the
solute. If a strong electrolyte dissociates into two ions, it will be twice as
effective at changing the boiling or freezing point. If a strong electrolyte
dissociates into three ions, then it is three times as effective as the same
number of moles of a nonelectrolyte. If a weak electrolyte dissociates 5% then
it would be 5% more effective than an equal number of moles of a
nonelectrolyte.
Vapor Pressure Change
Another property of a liquid that changes when a solute is added to it, is
the vapor pressure.
First let's review what vapor pressure is. As you know, liquids will evaporate.
The rate and extent to which it evaporates depends on the temperature. If you
put a liquid at a certain temperature into a closed, evacuated container it will
evaporate until the vapor exerts a certain amount of pressure. That pressure
is called the vapor pressure.
An evacuated container is not necessary. If the liquid evaporates into air, the
vapor is mixed with air and the pressure that it exerts cannot be measured
directly. However, the vapor does still exert pressure and it is still called the
vapor pressure of the liquid.
The amount of evaporation increases when the temperature increases. When
the temperature is such that the vapor pressure is equal to the atmospheric
pressure, the liquid boils. That temperature is called the boiling point. (This is
why cooking times and techniques have to be adjusted for high altitude
cooking. Since, at high altitudes, the atmospheric pressure is lower than at
sea level, the temperature needed to raise the vapor pressure equal to the
atmospheric pressure is lower. A pot of water set to boil pasta, for example,
will boil at a temperature lower than 100 oC.)
This graph shows the vapor
pressure of pure water
increasing as the temperature
increases. When the
temperature reaches 100°C
the vapor pressure reaches 1
atm and the water boils.
When a solute is added to a
liquid, it will decrease the
vapor pressure of the liquid,
as long as the solute itself is
not volatile.
In fact this is the way in which
the boiling point is affected.
Remember, the boiling point
is the temperature at which
the vapor pressure equals the
atmospheric pressure. So if
you add something to the
solution that decreases the
vapor pressure - then you will
have to heat the liquid to
even higher temperatures
before the vapor pressure
reaches atmospheric
pressure. That's how the
boiling point is increased.
Osmosis
A phenomenon that is related somewhat to the change in freezing point, the
change in boiling point, and the change in vapor pressure of solutions when
compared to pure solvents, is the process of osmosis. I'm sure you've heard
of it. I'd like you to take a look at it now (using the pictures on this pages) and
later when you are in the lab. I also recommend that you record your
observations in exercise 18 in your workbook.
This osmosis apparatus
contains a pure water separated
from a solution by a thin
membrane. So we have three
things: a solution, a membrane,
and pure solvent. Water is the
pure solvent in this case.
By looking at the
change in liquid
levels, you can see
that something has
moved from the
water, through the
membrane, into the
solution. If something
came from pure water
it had to be water
molecules because
pointer marks initial
there was nothing
position
else there.
2 minutes
-------------->
So the water passed through the membrane from
the water side to the solution side. Actually, water
passes through the membrane in both directions
but it moves faster into the solution than out of it.
Water can go in either direction, but the solute
particles within the solution cannot pass through
the membrane. What seems to happen is that the
top of tape marks
initial position
presence of the solute particles restricts the flow of
water molecules from the solution into the pure
solvent. But they are not able to restrict the motion
of the pure solvent into the solution. Consequently,
the water passes from the pure solvent to the
solution in greater amounts than water molecules
from the solution pass into the pure solvent.
So, the water passes through the membrane in both directions, but it passes
through at a higher rate from the pure solvent into the solution, than it passes
from the solution into the pure solvent. That process is called osmosis.
Another aspect of osmosis is something called osmotic pressure. As solvent
molecules pass through the membrane into the solution, they build up
pressure. If the side with the solution were closed off and had a pressure
gauge mounted on it, you would be able to read the osmotic pressure
generated by the flow of water into the solution.
Here is another way of describing osmotic pressure. The flow of water into the
solution can be stopped by applying pressure to the side where the solution
is. The amount of pressure needed to stop the flow is the osmotic pressure.
Actually the flow of molecules is not stopped, but the flow is equal in both
directions and cancels out.
It is not necessary for one of the liquids to be pure water in order for osmosis
to occur. All that is necessary is for the concentrations on the two sides of the
membrane to be different.
Isotonic Solutions
In the diagram shown here the
dotted line represents a
semipermeable membrane
through which water molecules
(but not solute particles) can
pass. The small dots represent
water molecules and the larger
red dots represent solute
particles. Note that the solute
concentration is the same on
both sides of the membrane.
The solutions are said to
be isotonic compared to one
another.
Because the solute
concentrations are the same on
both sides of the membrane,
water molecules move through
the membrane equally well in
both directions. There is no net
flow of water in either direction.
Hypertonic Solutions
In this diagram the solution on
the left side of the membrane
has a higher solute
concentration than the
solution on the right side of
the membrane. The solution
on the left is said to
be hypertonic compared to
the one on the right.
The higher solute
concentration on the left
reduces the flow of water
molecules from left to right,
causing a net flow of water
from the right to the left. If the
right side represented a cell
placed in a hypertonic
solution, water would leave
the cell causing it to
dehydrate and collapse.
Hypotonic Solutions
In this diagram the solution on the
left side of the membrane has a
lower solute concentration than
the solution on the right side of the
membrane. The solution on the
left is said to be hypotonic
compared to the one on the right.
The lower solute concentration on
the left allows for increased flow of
water molecules from left to right,
causing a net flow of water from
the left to the right. If the right side
represented a cell placed in a
hypotonic solution, water would
enter the cell causing it to swell
and perhaps burst.