Ice Cream Chemistry: Colligative Properties and Freezing
Point Depression
Introduction:
Solution properties that depend only on the concentration of solute particles in solution are called colligative properties.
The identity of the solute does not matter. Vapor pressure, boiling point, and freezing point are colligative properties that
depend on concentration of solute present.
Vapor Pressure: Substances such as gasoline and alcohol have high vapor pressures and evaporate quickly from open
containers. Water, on the other hand, has a much lower vapor pressure at equivalent temperatures and evaporates at a
much lower rate. In order to bring water to a boil at standard atmospheric pressure, its temperature must be raised to 100
°C. At this temperature, the molecules have sufficient energy to overcome the forces of attraction that keep them together
and go into the vapor phase. The vapor pressure of water has been raised to a point where it is now equal to the prevailing
atmospheric pressure.
Throughout our discussion of vapor pressure and its relationship to boiling points, we have been careful to emphasize we
were dealing with pure substances. Pure water boils at 100°C and freezes at 0°C when the pressure is 760 torr (1 atm), but
what happens when molecules or ions of a solute are added to our solvent? In essence, you are diluting the water
molecules, actually diluting the solvent itself. The solute particles get in the way of the H2O’s. They can’t act like pure
water molecules anymore. Hydrogen bonding and dipole attractions to each other are altered.
So, for example, when you try to boil water that has a solute in it, the H2O molecules can’t escape to vapor as easily.
There are fewer H2O molecules at the surface because some of the surface spaces have been taken up by the solute
particles. Many of the H2O molecules on their way to the surface are also colliding with the solute particles and bouncing
into the depths of the solution. Because of this, the water solution’s vapor pressure is lower than that of pure water.
Boiling Point: You can get the number of escaping water molecules in a solution to increase by speeding them up. So to
get a solution to boil, you have to heat it hotter than if it were just pure water. In other words, the boiling point of the
solution will be higher than 100°C, and the more solute particles there are, the higher the boiling point will be. Water’s
boiling point goes up 0.52 Co for every mole of solute particles it contains per kilogram of water. So we have a regular
conversion factor (Kb) between the concentration (in molality = m) and temperature rise (oC):
Boiling point elevation constant for water = Kb = 0.52 oC/m
The relationship between the change in boiling point and the concentration of a solution is expressed as:
∆𝑇𝑏 = 𝑖𝐾𝑏 𝑚
Tb is the boiling point increase as compared to pure solvent
Kb is the boiling point elevation constant for the solvent
i is the van’t off factor that explains the number of molecules that are formed when the substance is dissolved in solution
m is the molality of the solution:
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝑚=
𝑘𝑔 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡
Freezing Point: Water with a solute in it should have a harder time getting its molecules all lined up neatly into the solid
crystal arrangement of ice. So we would have to cool it down (to slow down the molecules) an extra amount below 0°C, if
we want it to freeze. Water’s freezing point actually goes down 1.86 C° for every molal’s worth of solute particles it
contains. In fact, any solute will affect (lower) the freezing point of a solvent by a more significant amount than it will
affect (raise) the boiling point.
The relationship between the freezing point and the concentration of a solution is expressed as:
∆𝑇𝑓 = 𝑖𝐾𝑓 𝑚
Tf is the boiling point increase as compared to pure solvent
Kf is the freezing point elevation constant for the solvent
i is the van’t off factor
m is the molality of the solution
Calculations of expected freezing points are done very similarly to the examples we saw for prediction of boiling points.
The only differences are 1) we use the Kf rather than the Kb, and 2) we must remember that our final freezing point is
below that of the pure solvent rather than above.
Pre-Lab
1. What are colligative properties? Give an example of a colligative property.
2. Draw a picture (at the molecular level) of how adding solute impacts the freezing point of a solution. Be sure to
indicate how solute-solvent interactions.
3. You dissolve 50 g of NaCl in 0.75 kg of water. Calculate the molality of this solution.
4. Using the Kf from the pre-lab, calculate the freezing point depression of the solution made in question 3.
Materials Available:
Materials
Beaker, 600 mL
Test Tubes
Thermometer
Sodium Chloride, NaCl
Sucrose
Ice
Water
Quantity
1
6
1
~100 g
~100 g
As Needed
As Needed
Procedure:
Part I: Test of Freezing Point Depression
1. You will be designing and running a lab based on the information that you have researched. Based on the
introduction and pre-lab material, formulate a research question that will allow for you to study how the amount of
salt and sugar dissolved in water impacts both the freezing point.
Our Research Question: ________________________________________
___________________________________________________________.
2. Propose a hypothesis that answers your question above:
Our Hypothesis: _____________________________________________
___________________________________________________________.
3. Using the materials available and the information that you have gained, outline a procedure below that will test the
following:
 How the amount of salt dissolved alters the freezing point of water.
 How the amount of sugar dissolved alters the freezing point of water
 NOTE – You will have a 0.1 m and 0.01 m of each solution.
Be sure to clearly identify which pieces of equipment (listed in the Materials section) you plan to use during your
investigation.
4. Make a simple sketch showing how you plan to set up your equipment
5. Create data tables for what you will record data on and where you will put this data. Put this is the data collection
section of this handout.
6. Consult with your instructor for approval before beginning your experiment.
7. As you conduct your experiment, keep detailed written records. Be sure to list all steps taken as you perform your
experiment, and all measurements and observations made during the experiment.
Analysis:
1. Calculate the theoretical freezing point depression for each of the solutions that you tested. Remember you can
calculate this by:
∆𝑇𝑏 = 𝐾𝑏 𝑚
You should have had 0.01 and 0.1 m solutions of both salt and sugar.
2. Using the actual values and the theoretical values calculate the percent error for each trial and record in the
following table:
𝑀𝑒𝑎𝑠𝑢𝑟𝑒𝑑 𝑉𝑎𝑙𝑢𝑒 − 𝐶𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝐸𝑟𝑟𝑜𝑟 =
× 100%
𝐶𝑎𝑙𝑐𝑢𝑎𝑙𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒
Solution
Percent Error
0.1 m NaCl
0.01 m NaCl
0.1 m Sugar
0.01 m Sugar
3. Formulate a hypothesis about the percent error was higher for some solutions than it was for others. Discuss your
reasons with your instructor before moving on.
Part II: Making Ice Cream (EXTRA CREDIT)
We use the principle of freezing point depression in numerous ways. Solute molecules (called antifreeze) are added
to the water in car radiators. Antifreeze both lowers the freezing point and raises the boiling point so it doesn’t freeze
in the winter and is less likely to boil over in the summer. Any solute molecules or ions would do the job but we use
ethylene glycol because it is non-corrosive and non-volatile. It doesn’t corrode the cooling system or evaporate away
when the engine is hot.
Dissolving salt in water will lower its freezing point. For example, to melt ice on streets and sidewalks, salt is often
sprinkled on it. Unless it is extremely cold, there is always a slight film of liquid water on the surface of ice, (at least
there is as soon as someone steps on it or a car drives over it). The salt dissolves in this film of water, lowering its
freezing point (and that of the ice surface it is in contact with) below the air’s temperature, and the surface layer of
the ice melts. This puts salt water in contact with the next layer of ice molecules, and they melt, and so on until either
all the ice is melted or the salt water gets too dilute to melt more, considering the temperature of the weather.
Anything that dissolves in water could be used. Salt is used because it is cheap. Unfortunately, it’s corrosive to steel.
Many northern states are facing a problem where their bridges may start falling down because of the tons of salt that
have been spread on their roadbeds.
When there is a lot of ice melting going on in a short time, there is a lot of heat of fusion being used up. Because of
this, near an ice/salt mixture, it gets pretty cold. This rapid absorption of heat is the reason we can enjoy ice cream
during the summer months.
Pre-Lab
1. How is homemade ice cream created? Be sure to list the ingredients that are required.
2. One of the keys of making ice cream is immersing the mixture in ice in order to cool it. The problem with this step
is that the solution needs to be below the normal freezing point of ice. How do you propose you could lower the
freezing point of this ice bath that the contents will be immersed in?
3. What are some differences between the homemade ice cream procedure that you discovered and the process we
use to mass-produce ice cream? Provide at least 5 differences.
Materials Available:
Materials
Milk
Eggs
Sugar
Cream (½ and ½)
Gallon Sized Ziplock Back
Large Container
Ice
Salt
Quantity
As Needed
As Needed
As Needed
As Needed
1–2
1
As Needed
As Needed
Procedure:
Part I: Test of Freezing Point Depression
1. You will be designing and running a lab based on the information that you have researched. Based on the
introduction and pre-lab material, formulate a research question about how to make the best ice cream.
Our Research Question: ________________________________________
___________________________________________________________.
2. Propose a hypothesis that answers your question above:
Our Hypothesis: _____________________________________________
___________________________________________________________.
3. Using the materials available and the information that you have gained, outline a procedure below that will test the
following:
 How do you make the ideal ice cream from scratch?
Be sure to clearly identify which pieces of equipment (listed in the Materials section) you plan to use during your
investigation.
Data Collection
Data Tables:
Discussion:
1. Does your data indicate that the initial hypothesis should be accepted or rejected? Cite specific data to support your
acceptance or rejection!
2. What were the inherent structural errors that were introduced as a result of your experimental design? How did these impact
your final data? How can this be fixed in order to improve your results should you repeat the lab?
3. What next steps could you take in order to further investigate this topic? Cite specific areas that you believe you could study
and how your data allows you to better study these areas.
4. In the early days, methanol was used universally as antifreeze. In 1937, antifreeze solution based on ethylene glycol were
introduced and quickly took the place of methanol, despite the fact that it costs more than methanol. Research into the
properties of both substances and propose why ethylene glycol is a better antifreeze.