Freezing Point Depression
How Low Can you Go?
Introduction: People who live in the northern states are familiar with winter and the
snowy, icy roads that go with the season. Road crews spread salt (sodium chloride, calcium
chloride, or a salt mixture on the roads in order to lower the temperature at which freezing
occurs. If the road already has ice on it, the salt helps to melt the ice, forming a solution with
a lower freezing point than that of pure water.
Background: The freezing point of a liquid is the temperature at which the forces of
attraction among molecules are just great enough to cause a phase change from the liquid
state to the solid state; the phases are in equilibrium.
During the freezing process of water, the water molecules come together to form the more
orderly, crystalline pattern of ice molecules. When a solute is added to a pure solvent (such
as water), the crystalline pattern is interrupted by the presence of salt “impurity.” Salt and
other dissolved solutes interfere with the ability of the solvent to crystallize, and the solution
remains liquid even at a temperature below the freezing point of the pure solvent. Thus the
solution has a lower freezing point than its pure solvent; this is called freezing point
depression.
Objective: The purpose of this experiment is to measure the freezing point of pure water
and the freezing point depression for various solutions. The effect of the concentration and
number of dissolved solute particles on the freezing point of water will be determined.
Pre-Lab Questions: Read the procedure carefully, and answer the following questions.
1. What factors will be held constant when determining the freezing point depression for
each solute in the experiment?
2. Why is it important to keep the thermometer off the bottom of the beaker?
3. Why is it necessary to measure the temperature of the pure ice-water mixture,
instead of assuming it to be 0.00C?
Materials: Aluminum chloride hexahydrate
Balance
Calcium chloride dehydrate
250 mL beakers (4)
Sodium chloride
Thermometer
Sucrose
Stirring rods (4)
Crushed Ice
Weighing dishes (4)
Water
Graduated Cylinder
Safety Precautions: Avoid contact of all chemicals with eyes and skin. Wash hands
thoroughly with soap and water before leaving the lab.
Procedure: Form a working group of 4 students. Each pair will measure the freezing point of
water with their thermometers (steps 2-6), and then determine the freezing point depression
for two of the solutes. Both pairs of students in each group will share their data and complete
the data table and questions.
1. Label four 250 mL beakers #1-4.
2. Add approximately 100.0 grams of ice-water mixture to the beaker by (a) first adding 70-80
grams of crushed ice and (b) enough water so the total mass of the ice-water mixture is
approximately 100.0 grams. Note: If the balance does not have a large enough capacity,
weigh the ice in a weighing dish; measure the remaining water using a balance or graduated
cylinder.
3. Stir the ice-water mixture. Wait for the temperature to stabilize and record the
temperature to the nearest tenth of a degree.
Beaker #1 – Sodium Chloride
4. Weigh approximately 30.0 grams of sodium chloride and record the mass.
5. Add the sodium chloride to the ice water mixture in beaker #1; stir until slushy.
6. Stir the contents of the mixture while recording the lowest temperature obtained before
the temperature of the mixture starts to rise again.
Beaker #2 – Sucrose
7. Repeat steps 2 – 6 for Beaker #2, using 100.0 grams of ice water and 30.0 grams of sucrose.
Beaker #3 – Calcium Chloride Dihydrate
8. Repeat steps 2 – 6 for beaker #3.
Beaker #4 – Aluminum Chloride Hexahydrate
9. Repeat steps 2 – 6 for beaker #4.
Data Table:
Beaker
Number
Solute
Mass of Ice
and Water
Mass of
Solute
Freezing
Point of
Water
Lowest
Temperature
of Mixture
Post-Lab Calculations: Show All Work
1. Determine the molar mass of each solute and calculate the number of moles of each
solute.
2. Calculate the experimental value of the freezing point depression for each solute.
3. Calculate the ∆Tf per mole of solute.
4. Write balanced dissociation reactions for all that apply, and determine the ΐ for each
solute.
5. Calculate the concentration, molality, of each solution.
6. Determine the theoretical value of the freezing point depression for each solute based
on the molality of the solution and the ideal number of solute particles (ΐ).
7. Compare the experimental and theoretical values of ∆Tf for each solute. Discuss
possible sources of error in your conclusion.
8. Which solute has the greatest freezing point depression per mole? Which had the
least? Is this what would be expected?
9. Prepare a graph of∆Tf / mole on the y-axis versus ΐ on the x-axis. Describe the
relationship shown by the graph.
Results Table:
Beaker Number
Solute
Molar Mass
Moles of Solute
∆Tf (exp)
[∆Tf(exp)]/ mole
Ϊ (ideal)
m
∆Tf (theor)
1
NaCl
2
C12H22O11
3
CaCl22H2O
4
AlCl36H2O