PURPOSE:  To investigate the freezing point depression caused by adding a salt to a solvent.

MATERIALS (for two students):  1000 ml beaker, 500 ml beaker, 2 new test tubes (15-18 mm diameter), thermometer, 2 flexible plastic drinking straws, salt (NaCl or CaCl₂), ice cream mixture, pails to collect snow, cheesecloth square or fine-mesh strainer.

INTRODUCTION: It is 5 AM on a snowy, winter morning and the temperature is -5^oC.  Suddenly a snow plow rumbles down the road and sprays a salt over the road.  The ice and snow slowly begin to melt making it safer for motorists.  How did the application of a salt cause the snow to turn into a liquid at -5^oC, well below the freezing point of water?

Dissolving a nonvolatile solute, such as a salt, affects several properties of the solvent.  The vapor pressure is lowered, the boiling point is raised, and the freezing point is lowered.

For dilute solutions, the decrease in freezing point and the increase in boiling point are directly proportional to the number of moles of solute particles dissolved per 1000 grams of solvent (molality).  To be more exact, the change in temperature ( _ T) can be determined by multiplying molality (m) by the molal freezing point constant for the solvent (K_f) using the equation:

D T = m ^. K_f

The accepted K_f value for water is 1.86 ^oC/m.

People made use of this property of solutes in making ice cream long before automobiles existed.  Because a cream mixture has a lower freezing point than pure water, it is not possible to make ice cream by packing it in plain ice.  Therefore, ice and salt were packed around a canister containing the cream mixture.  Many layers of snow and salt were needed as family members took turns cranking the ice cream machine until the mixture finally froze!

You will measure the freezing point depression of a solution of salt and melted snow while making your own ice cream.  Using this information, you will calculate the apparent molality of the salt solution (based on the freezing point depression you measure) and compare it to the theoretical value (based on the mass of salt and snow in the solution).

PROCEDURE:

A.  Working with a partner, weigh out 50.0 grams of a salt and set aside.

B.  Pack 2-3 cm of snow in the bottom of the large beaker.  Record the temperature of the snow.

C.  Sprinkle about a quarter of the salt over the snow.  Continue to add layers of snow and salt until the beaker is nearly full. Use all the salt.

D. Fill a new test tube about two-thirds full of the ice cream mixture and insert a straw, crimped end down.

E. Place the test tubes into the beaker so that the liquid inside the test tube is completely surrounded by snow.  Do not allow the snow-salt mixture to overflow into the cream mixture as it spoils the taste.

F.  Place the thermometer into the snow-salt mixture and begin recording the temperature every minute until the ice cream has frozen completely.  Use the straw to stir the cream mixture to speed solidification.  Add more snow as needed so that the cream mixture in the test tube is always surrounded by snow. Record the temperature at which the ice cream mixture freezes.  Record the minimum temperature reached by the snow-salt mixture.

G.  When the ice cream is frozen warm the test tube slightly in your hand and pull gently on the straw to remove the ice cream.  Enjoy!

H.  Mass the smaller beaker.

I.  Pour the snow-salt mixture through cheesecloth or a strainer into the smaller beaker to filter out the unmelted ice.  Mass the solution in the beaker and calculate the mass of the solution.

DATA

Chemical formula of salt ……….….________

Mass of salt....................…….……________

Initial snow temperature……...…… ________

Ice cream freezing temperature….…________

Minimum temperature of solution…..________

Mass of small beaker...........……....________

Mass of beaker and solution…........________

Total mass of solution..........……....________

CONCLUSIONS:  The equation D T = m ^. K_f makes the assumption that the salt ionizes completely in solution.  This is true only for very dilute solutions.  At higher concentrations, negative ions are attracted to positive ions an ionization is incomplete.  The data you collected in this experiment -- which had the side benefit of a useful product, the ice cream -- will allow you to determine if incomplete ionization occurred in the salt solution.  To find out, do the following calculations to compare the theoretical molality of the salt solution (based on the mass of salt and snow in the solution) to the apparent molality of the salt solution (based on the freezing point depression you recorded).

1. Calculate the number of moles of salt added to the snow.

2. Calculate the molality of the salt solution.

3. Calculate the theoretical molality of the ions, assuming that the salt was 100% ionized.

4. Calculate the apparent molality of the solution, using the coldest temperature that you measured and the equation given above.

5. Calculate the percent ionization of the salt.  Did incomplete ionization occur in the salt-snow mixture as you made ice cream?

DISCUSSION:

6. Is there any limitation to the decrease in the freezing point of a solvent with any one solute?  Explain your answer.

7. If different salts were used by the students in the class, compare the results.  Explain the differences obtained.

8. To prevent freezing, a 50:50 mixture of ethylene glycol and water (antifreeze) should be used in car radiators.  Pure ethylene glycol will freeze at a higher temperature than the antifreeze solution.  Explain this observation.

9. This activity was conducted using snow at, or near, its freezing point and a cream mixture a few degrees above 0^oC.  Both the freezing of the ice cream and the decrease in temperature of the snow-salt mixtures represent a loss of heat.  Explain how this decrease in heat energy could occur in a warm environment such as your classroom.