IT'S GETTING COLDER (FREEZING POINT DEPRESSION)
The purpose of this experiment is to
demonstrate the effect of solutes on the freezing point of water.
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, 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.
Two lab periods.
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
- ice C12H22O11 (sucrose)*
- commercial automotive antifreeze*
- test tubes
- 100-mL graduated cylinder
- stirring rod
- Table sugar may be substituted for reagent grade sucrose.
- Table salt can be substituted for laboratory grade NaCl.
- Since antifreeze is mostly ethylene glycol, the chemical reagent itself
could be used.
- A-10 ounce Styrofoam cup can be substituted for the 400-mL beaker.
Preparation of Ice Bath
Determination of Freezing Points of Solutions:
- Fill the large beaker 3/4 full with ice.
- Cover the ice with 1/4 to 1/2 inches of table salt.
- Stir this ice-salt mixture with a stirring rod and make sure the
temperature drops to at least -10°C.
- Prepare a solution of NaCl by adding 5.8 grams of NaCl to 100 mL of
water. Mix until all crystals dissolve.
- Prepare a solution of sucrose by adding 34 grams of sucrose to 100 mL
of water. Mix until all crystals dissolve.
- Place a test tube that is 1/2 full of water in the ice bath.
- Stir the water in the test tube gently with a thermometer while keeping
track of the temperature.
- 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).
- Repeat steps 3-5 with the prepared NaCl and sucrose solutions.
- Calculate the molalities of the NaCl and sucrose solutions.
- Using the equation, 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
All solutions may be flushed down the
drain with plenty of water.
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
- Dissolve 6.2 grams of commercial antifreeze in 100 mL of water.
- Freeze this solution in the same manner as in the previous
experiment. Be sure to record the freezing point temperature.
- Calculate the molecular mass of this solute based on the freezing
Molecular mass of solute = [(Kf) (grams of
solute)] ÷ [(T) (kg of solvent)]
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.
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
- This activity is easily accomplished in two
parts and can be completed on different days.
- The freezing point depressions
of the salt and sucrose solutions could serve as the starting points of a study
of colligative properties.
Submitted by Joe MacQuade, Gina Monks, Doug Rickard, Irene Walsh, and Joe Wilkins
Woodrow Wilson Leadership Program in Chemistry
The Woodrow Wilson National Fellowship Foundation
CN 5281, Princeton NJ 08543-5281