Abstract

Objective

The goal of this project is to investigate how dissolving substances in water affects the freezing point of the solution.

Introduction

To make ice cream with an old-fashioned hand-crank machine, you need ice and rock salt to make the cream mixture cold enough to freeze. If you live in a cold climate, you've seen the trucks that salt and sand the streets after a snowfall to prevent ice from building up on the roads. In both of these instances, salt is acting to lower the freezing point of water.

For the ice cream maker, because the rock salt lowers the freezing point of the ice, the temperature of the ice/rock salt mixture can go below the normal freezing point of water. This makes it possible to freeze the ice cream mixture in the inner container of the ice cream machine. For the salt spread on streets in wintertime, the lowered freezing point means that snow and ice can melt even when the weather is below the normal freezing point of water. Both the ice cream maker and road salt are examples of freezing point depression.

Other substances that dissolve in water also lower the freezing point of the solution. In this project, you'll investigate how the freezing point of a diluted solution changes with the amount of solute added.

Terms, Concepts, and Questions to Start Background Research

To do this project, you should do research that enables you to understand the following terms and concepts:

• Colligative properties
• Molecular weight
• Moles
• Molality
• Freezing point depression

Questions

• When sodium chloride dissolves in water, how many solute molecules result from each molecule of solid dissolved?
• When sucrose dissolves in water, how many solute molecules result from each molecule of solid dissolved?

Bibliography

• For more information on colligative properties, see:
  • Eli, Todd & Keith, date unknown. "Colligative Properties," Chemworld, ThinkQuest Library, Oracle Education Foundation [accessed March 28, 2007] http://library.thinkquest.org/C006669/data/Chem/colligative/colligative.html?tqskip1=1.
  • Nave, C.R., 2006. "Colligative Properties of Solutions," HyperPhysics, Department of Physics and Astronomy, Georgia State University [accessed March 28, 2007] http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/collig.html.
• For information on Avogadro's number and molecular weight, see:
  
  • Lachish, U., 2000. "Avogadro's Number, Atomic and Molecular Weight," [accessed March 28, 2007] http://urila.tripod.com/mole.htm.
  • Furtsch, T.A., date unknown. "Avogadro's Number," Tennessee Technological University [accessed March 28, 2007] http://gemini.tntech.edu/~tfurtsch/scihist/avogadro.htm.
• To try a simulated experiment on freezing point depression or boiling point elevation, see (Flash animation, requires browser plug-in):
  Greenbowe, T.J., 2005. "Boiling-Point Elevation and Freezing-Point Depression," Department of Chemistry, Iowa State University [accessed March 28, 2007] http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/propOfSoln/colligative.html.

Materials and Equipment

To do this experiment you will need the following materials and equipment:

• Thermometer capable of reading at least −10°C
• Large styrofoam cup (12 oz. or more, or 400 mL Pyrex® beaker)
• 100 ml graduated cylinder
• Gram balance (accurate to 0.1 gram), such as the Fast Weigh MS-500-BLK Digital Pocket Scale, 500 by 0.1 G, available from Amazon.com
• Test tubes
• Salt
• Sugar
• Water
• Ice
• Spoon

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Experimental Procedure

Preparation of Ice Bath

1. Fill the styrofoam cup (or beaker) 3/4 full with ice.
2. Cover the ice with 1/4 to 1/2 inches of table salt.
3. Stir this ice-salt mixture with a spoon and make sure the temperature drops to at least −10°C.
4. You will be using this ice bath to freeze many samples of test solutions. During the course of your experiments, you may need to pour off melted water from the ice bath and to replenish the ice and salt. Wait until the temperature drops to at least −10°C before continuing.
5. Always carefully rinse and dry the thermometer before using it to measure the freezing point of your test solutions. You don't want the thermometer to carry salt water into your test solutions!

Determination of Freezing Points of Solutions

1. You will measure the freezing point for seven different liquids:
   1. Pure water
   2. NaCl in water (three different concentrations)
   3. Sucrose in water (three different concentrations)
2. Preparing NaCl solutions (label each solution clearly):
   1. Prepare solution #1 of NaCl by adding 5.8 g of NaCl to 100 mL of water. Mix until all crystals dissolve.
   2. Prepare solution #2 of NaCl by adding 4.35 g of NaCl to 100 mL of water. Mix until all crystals dissolve.
   3. Prepare solution #3 of NaCl by adding 2.9 g of NaCl to 100 mL of water. Mix until all crystals dissolve.
3. Preparing sucrose solutions (label each solution clearly):
   1. Prepare solution #1 of sucrose by adding 34 g of sucrose to 100 mL of water. Mix until all crystals dissolve.
   2. Prepare solution #2 of sucrose by adding 25.5 g of sucrose to 100 mL of water. Mix until all crystals dissolve. Prepare solution #3 of sucrose by adding 17 g of sucrose to 100 mL of water. Mix until all crystals dissolve.
4. Fill a test tube half-way with the liquid to be tested, and place it in the cup with the ice and salt.
   1. The liquid in the test tube should be below the level of the ice and salt in the cup.
   2. Do not allow any ice or salt from the cup to get into the test tube.
5. Stir the solution in the test tube gently with a thermometer while keeping track of the temperature.
6. When the first ice crystals appear on the inside wall of the test tube, record the temperature. This is the freezing point of the liquid.
7. Repeat steps 4 and 5 with each of the three liquids. To assure yourself that your results are consistent, you should do at least three separate freezing-point determinations for each liquid.
8. Calculate the molalities of the NaCl and sucrose solutions.
   1. Molality is defined as the number of moles of a substance, divided by the weight (in kg) of the solvent.
   2. The number of moles of a substance is defined as the weight of the substance (in g) divided by the gram molecular weight of the substance.
   3. The gram molecular weight of NaCl is 58.443.
   4. The gram molecular weight of sucrose is 342.3.
   5. 100 ml of water weighs 0.1 kg.
9. Make a table of your results, like the following:

   solution	g substance	molecular weight substance	amount water (kg)	molality	freezing point (Tf, °C)	ΔT
   NaCl #1	5.8	58.443	0.1	1	?	?
   NaCl #2	4.35	58.443	0.1	0.74	?	?
   etc.

10. Calculate the freezing point depression, ΔT, by subtracting the freezing point of the test solution from your measured freezing point for the plain water solution.
11. For each of your NaCl and sucrose solutions, graph the average amount of freezing point depression (ΔT, y-axis) vs. molality.

Variations

• Can you explain your results using the equation: where i is the number of particles produced in solution per molecule of the solid, and K_f for water = 1.86°C/m.
• For some other chilly experiments, see the Science Buddies projects:
  • Supercooling Water and Snap Freezing, and
  • Investigating the 'Mpemba Effect': Can Hot Water Freeze Faster than Cold Water?

• For more science project ideas in this area of science, see Cooking & Food Science Project Ideas.

Credits

Andrew Olson, Ph.D., Science Buddies

Sources

This project is based on:

• MacQuade, J., et al., 1986. "It's Getting Colder (Freezing Point Depression)," Woodrow Wilson Summer Chemistry Institute, Woodrow Wilson National Fellowship Foundation [accessed March 28, 2007] http://www.woodrow.org/teachers/chemistry/institutes/1986/exp9.html.

Last edit date: 2011-11-18 12:00:00