[WW HOME] [CHEMISTRY] [CONSUMER] [SEARCH] [DISCLAIMER]

IT'S GETTING COLDER (FREEZING POINT DEPRESSION)

* PURPOSE

The purpose of this experiment is to demonstrate the effect of solutes on the freezing point of water.

* DESCRIPTION

This experiment is suitable for a first-year course. It could also be included in a second-year course as part of a unit on colligative properties. The calculations require the use of the equation, [delta ]T = (Kf ) (m) (i). In the first part of the experiment, various solutes are added to water and the resultant freezing points of the solutions are determined. The value of i, dissolved particles per formula unit, is calculated. In the second part, the molar mass of commercial antifreeze is determined.

* TIME REQUIRED

Two lab periods.

* MATERIALS

Chemicals:
ice C12H22O11 (sucrose)*
NaCl
commercial automotive antifreeze*
Equipment:
test tubes
thermometer
400-mL beaker*
100-mL graduated cylinder
stirring rod
*See Modifications/Substitutions

* HAZARDS

Since commercial antifreeze is primarily ethylene glycol, it is highly toxic and should not be ingested. The ice used in the experiment could become contaminated with antifreeze by accident; students should be warned not to eat the ice. Goggles must be worn throughout the experiment.

* MODIFICATIONS/SUBSTITUTIONS

  1. Table sugar may be substituted for reagent grade sucrose.
  2. Table salt can be substituted for laboratory grade NaCl.
  3. Since antifreeze is mostly ethylene glycol, the chemical reagent itself could be used.
  4. A-10 ounce Styrofoam cup can be substituted for the 400-mL beaker.

* PROCEDURE

Preparation of Ice Bath

  1. Fill the large 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 stirring rod and make sure the temperature drops to at least -10°C.
Determination of Freezing Points of Solutions:
  1. Prepare a solution of NaCl by adding 5.8 grams of NaCl to 100 mL of water. Mix until all crystals dissolve.
  2. Prepare a solution of sucrose by adding 34 grams of sucrose to 100 mL of water. Mix until all crystals dissolve.
  3. Place a test tube that is 1/2 full of water in the ice bath.
  4. Stir the water in the test tube gently with a thermometer while keeping track of the temperature.
  5. When the first ice crystals appear on the inside wall of the test tube, record the temperature. This should be the freezing point of the liquid. (In this step water is the pure solvent).
  6. Repeat steps 3-5 with the prepared NaCl and sucrose solutions.
  7. Calculate the molalities of the NaCl and sucrose solutions.
  8. Using the equation, [delta ]T = (Kf)(m)(i), determine the value of i, where i is the number of particles produced per formula unit and Kf for water = 1.86°C/m.

Molecular Mass Determination From Freezing Point Depression

  1. Dissolve 6.2 grams of commercial antifreeze in 100 mL of water.
  2. Freeze this solution in the same manner as in the previous experiment. Be sure to record the freezing point temperature.
  3. Calculate the molecular mass of this solute based on the freezing point depression.
    Molecular mass of solute = [(Kf) (grams of solute)] ÷ [([delta ]T) (kg of solvent)]

* DISPOSAL

All solutions may be flushed down the drain with plenty of water.

* DISCUSSION

Colligative properties of solutions depend upon the concentration of solute particles. The freezing points of water solutions are always lower than that of pure water. The change in freezing point caused by the presence of a solute dissolved in water can be calculated from the equation,
[delta ]T = (Kf)(m)(i),
where Kf is the molal freezing point depression constant (1.86°C/m for water), m is the molality of the solution, and i is the number of particles produced per formula unit.
Molality = moles of solute/kg solvent
Since colligative properties depend upon the number of particles in solution, a one molal solution of an electrolyte (NaCl), which dissociates in water, lowers the freezing point more than a one molal solution of a non-electrolyte (sucrose). The freezing point of a one molal solution of NaCl is actually -3.37°C, only 1.81 times that of a non-electrolyte, not the -3.62°C that would be expected if NaCl were completely dissociated. This difference is believed to be due to the interionic attractions that prevent the ions from behaving as totally independent particles. The activity or effective concentration of the ions is less than would be indicated by the actual concentration. Some of the ions may exist as solvated units called an ion pairs. The more dilute the solution of an electrolyte, the more widely separated the ions, the less the interionic attractions, and the closer the effective concentration of the ions approaches the actual concentration.

* TIPS

  1. This activity is easily accomplished in two parts and can be completed on different days.
  2. The freezing point depressions of the salt and sucrose solutions could serve as the starting points of a study of colligative properties.

* REFERENCES

Holtzclaw, H.F., Robinson, W.R., and Nebergall, W.H.,College Chemistry with Qualitative Analysis, D. C. Heath and Company, Lexington, MA, 1984, p. 359. This work discusses colligative properties.

Submitted by Joe MacQuade, Gina Monks, Doug Rickard, Irene Walsh, and Joe Wilkins


[WW HOME] [CONSUMER] [SEARCH] [FEEDBACK] [DISCLAIMER]

Woodrow Wilson Leadership Program in Chemistry * lpt@www.woodrow.org
The Woodrow Wilson National Fellowship Foundation * webmaster@woodrow.org
CN 5281, Princeton NJ 08543-5281 * Tel:(609)452-7007 * Fax:(609)452-0066