Experiment 18 Colligative Properties
When a pure solid is dissolved in a solvent, several physical properties of that solvent change in a way that depend only on the relative amounts of the solute and the solvent present. Such properties are termed colligative properties. There are four common colligative properties: boiling point elevation, freezing point depression, vapor pressure reduction, and osmosis. The most familiar of these is the freezing point depression. If salt is placed on icy roadway, the freezing point of the water will be lowered, and the ice will melt, depending only on the amount of salt that has been added.
Pure water freezes at 0 °C. If 1 mol of dissolved particles, molecules, or ions is added to 1 kg of water, the freezing point of water drops from 0 oC to -1.86 oC. This drop is called the freezing point depression constant for water. Other liquids will have different constants. The relationship between the concentration of solute particles and the freezing point depression is expressed as
DT = kfp m
where DT is the change in the freezing point, kfp is the freezing point depression constant, and m is the molality of the solution. Molality is defined as the number of moles of solute in one kilogram of solvent.
Different solutes affect the freezing point depression in different ways. One mol of a molecular substance will lower the freezing point of 1 kg of water by 1.86 oC. One mol of NaCl, on the other hand, will lower the freezing point of 1 kg of water by 3.72 oC, twice as much. This is because NaCl will dissociate in water to form Na+(aq)and Cl-(aq) ions. Therefore twice as many particles are present in solution.
In Part A of this experiment, the freezing point of a pure solvent, cyclohexane will be determined. In Part B, a molecular solid, naphthalene, will be added to the cyclohexane, and the freezing point of the mixture will be measured. A plot of temperature (y axis) versus time (x axis) is used for this purpose. Figure Ia shows a typical cooling curve for pure cyclohexane. The temperature will level off, reaching a constant value. The freezing point is the constant temperature that occurs while cyclohexane is freezing.
Unlike pure samples, which have a sharp freezing point, solutions have freezing points that occur over a wider range. The cooling curve for a solution is similar, except that the temperature will continue to drop during crystallization (Figure Ib). As the cyclohexane in the solution freezes, the solution becomes more concentrated because less liquid remains. The molal concentration of the solute increases and the freezing point further decreases. The result is a steadily decreasing freezing temperature. The freezing point that you want is the initial freezing point because that is the only point at which you will know how much liquid solvent remains (Figure Ib).
A typical cooling curve for a pure substance. The freezing point is tf, the constant temperature in the middle of the cooling curve.
A typical cooling curve for a solution. Two straight lines have been drawn to describe the data immediately before and during freezing. The intersection of these lines occurs at tf’, the initial freezing point of the solution.
Materials and Equipment
unknown sample 3 large test tubes
stop watch 250 mL beaker
125 mL Erlenmeyer flask
rubber stopper & thermometer assembly (see Figure II)
A. Freezing Point of Pure Cyclohexane
Prepare an ice-water bath in a 250 mL beaker.
Obtain a large test tube from your drawer. Make sure it is clean and dry. Place the test tube inside an empty 125 mL Erlenmeyer flask. Tare the test tube and flask on the analytical balance and record their mass. Pipet 10.00 mL (±0.01 mL) of cyclohexane (density = 0.779 g/mL; MW = 84.1g/mol) in the test tube and determine the mass of the cyclohexane you are using for this experiment.
Remove the test tube from the flask and insert a stopper fitted with a thermometer and copper-wire stirrer (Figure II). Heat the test tube for about 30 seconds by holding it in your hands, then place it in the ice-water bath.
While one partner gently and constantly agitates the solution, the other should take temperature readings every 15 seconds until the cyclohexane is completely frozen. Remember to occasionally move the test tube around in the ice bath so that no warm water is allowed to pool around the test tube. Note the temperature at which the sample begins to crystallize. While stirring with the wire stirrer, continue to collect data for about two minutes.
Show the data to your Instructor. If it is necessary to repeat the procedure, use the same sample. Melt the cyclohexane and repeat steps if necessary
Clean and dry the stopper/thermometer assembly for the next procedure.
B. Freezing Point Depression Constant of Cyclohexane
Obtain another large test tube from your drawer. This time, weigh approximately 150 mg of naphthalene (MW = 128 g/mol) to the nearest mg. Follow step 1 from Part A to determine the exact mass of your sample.
Pipet 10.0 mL of cyclohexane into the test tube. If necessary, warm the solution in warm tap water in order to allow the naphthalene to dissolve completely. Remove the test tube from the Erlenmeyer flask and insert a stopper fitted with a thermometer and wire stirrer.
Repeat steps from part A to determine the freezing point of this solution. Use a salt-ice water bath instead of a ice water bath. Stirring the solution is very important to obtain a good cooling curve. Note the temperature at which the mixture begins to crystallize and the temperature at which it is fully crystallized.
Calculate the freezing point depression constant.
Complete a second run. Warm the mixture you have just used, and then proceed as before.
C. Molar Mass Determination of an Unknown
Clean the test tube, and obtain a fresh 10.00 mL (±0.01 mL) portion of cyclohexane.
Obtain an unknown sample and record its number on your data sheet. Add 100 mg of the unknown to the test tube, weighed to the nearest milligram.
Determine the freezing point of the resulting solution in the same way as you have done in parts A and B.
Plot a cooling curve for Parts A and B as shown in Figure I. Obtain the freezing points for the pure cyclohexane and for the naphthalene/cyclohexane solution. Determine the change in the freezing point, DT. Calculate the molality of the naphthalene/cyclohexane solution, and determine the freezing point depression constant.
For the unknown, use the equation given in the introduction, plus the kfp value determined from the cyclohexane/naphthalene data to calculate the molality, m. Since the number of kg of solvent is known, the number of moles of solute can be determined. Since the mass of the solute is known, the molar mass can be calculated.
Dispose of the cyclohexane, naphthalene, and the unknown in the waste beakers located under the hood.
Determine the freezing point for pure cyclohexane first. Record the mass and then set up a table to record the freezing point in 15-second intervals. Record the freezing point for cyclohexane.
Repeat the same procedure for parts B and C of the experiment.
Perform two trials and take the average for your freezing point of the mixture.
1. What would be the effect of each of the following on the calculated molar mass of the solute? Please explain.
(a) Some cyclohexane evaporated while the freezing point of pure cyclohexane was measured.
(b) Some cyclohexane evaporated after the solute was added.
(c) The thermometer was not properly calibrated. It gave a reading that was 1.5 °C lower during the course of the experiment.
2.(a) Why does the temperature of the solution during freezing not remain constant after the addition of naphthalene to cyclohexane?
(b) Why does the temperature of pure cyclohexane remain constant during freezing?