Chem 254
Experiment 6
Determination of Molar Mass and van’t Hoff Factor by Freezing Point
Depression
Colligative Properties
The properties of dilute solutions that only depend on the number of solute molecules present per unit
volume of solution, but not on the nature of the solute are called colligative properties, which are:
1. the vapor pressure lowering
2. the freezing point depression
3. the boiling point elevation
4. the osmotic pressure
For an ideal dilute solution, the chemical potential of a pure solvent μ*A is reduced to μ*A+RT lnxA. The
liquid vapour equilibrium occurs at a higher temperature, T’b, and solid-liquid equilibrium occurs at a lower
temperature T’f (Figure 1). The effect is greater on the freezing point because of the slopes of the lines.
chemical potential, μ
Solid
Pur
e liq
uid
So
luti
on
Vapour
T’f
Tf
Freezing point
depression
T’b Tb
Boiling point
elevation
Figure 1. The chemical potential of pure solvent and solvent of a dilute solution
When the solution frezees, the chemical potential of the liquid should be in equilibrium with the solid.
=
Since
Solution of the differential equation assuming that fusH is independent of T.
Using the approximations
and
When the solution is dilute, xB is proportional to molality, b and thus,
where
where Kf is freezing point constant. Kf values for some common solvents are given in Table 1.
Table 1. Kf values for some common solvents
-1
Kf (kg mol )
Benzene
5.12
Carbon tetrachloride
29.8
Ethanol
1.99
Water
1.86
Colligative properties of solutions depend on the total amount of solute particles. Thus, the extend of
dissociation of electrolytes, van’t Hoff factor, i affects colligative properties.
The ideal value of i for a strong electrolyte is the number of ions per unit formula. The measured value is
smaller than the ideal value due to the electrostatic attractions between the ions.
B
Tb
A
D
Tf
temperature, °C
temperature, °C
For pure substances, during a cooling process the temperature stays constant at the boiling and freezing
points until all the substances evaporate or freeze as shown in Figure 2.a. However, unlike with the pure
substances, most systems consisting of two or more components exhibit a temperature range over which
the solid and liquid phases are in equilibrium (Figure 2. b).
C
Time
a)
b)
Time
Figure 2. Cooling curve for a) a pure substance b) solution
Purpose: The aim of this experiment is to determine the molar mass of an unkown solute and van’t
Hoff factor of an electrolyte using colligative property, freezing point depression.
Apparatus and Chemicals
Apparatus: Constant temperature bath, appropriate containers, hand stirrer, thermocouple.
Chemicals: Solvents: water, benzene or similar; Solutes: urea, naphthalene or similar, NaCl or similar.
Procedure:
I. Determination of freezing point and freezing point constant, Kf of a solvent
A. Freezing point of the pure solvent.
1. Fill the jacket with 75 mL ethanol.
2. Put 20 mL of distilled water into the inner vessel and add magnetic stirrer.
3. Insert the protective sleeve for temperature sensor.
4. Drop ethanol onto the protective sleeve.
5. Insert temperature probe (T2) and connect it to the digital temperature meter.
6. Fix the device to support rod and position it as low as possible into the beaker with ice/salt
freezing mixture.
7. Set the magnetic stirrer to a medium stirring speed.
8. Immerse second temperature probe (T 1) and connect to the digital temperature meter.
9. Turn on the digital temperature meter.
10. Push T 1…4 button to choice measuring probes.
11. Push C/K button to choose the unit ˚C. Red light emitting diode indicates the upper digital display
and green light emitting diodes the lower part of the digital display.
12. Adjust the temperature at the freezing mixture to around -10˚C.
13. When the temperature in the inner vessel has reached approximately 0˚C, start recording
temperature every 30 second.
14. After the determination of freezing point of distilled water, lift the apparatus out of the freezing
mixture and wait until the water in the inner vessel has completely liquefied or warm the solution
gently.
15. Draw cooling curve of water
B. Freezing point constant, Kf of a solvent
1. Weight 0,3 g of urea and introduce the substance to the inner vessel and dissolve it completely.
2. Determine the freezing point of the mixture as described part 1, A,13 but take measurement when
temperature of the system drops to 0˚C.
3. Repeat the procedure for two more times by adding 0,3 g of urea to the solution at each time.
II. Determination of molar mass of an unknown solute
1. Take out the system, pour the water and fill it with 20 mL fresh distilled water.
2. Weight 0,3 g of unknown substance. Introduce substance to the inner vessel and dissolve it
completely.
3. Determine the freezing point of the mixture as described part A, 13 but take measurement when
temperature of the system drops to 0˚C.
III. Determination of Van’t Hoff factor of a strong electrolyte
1. Take out the system, pour the water and fill it with 20 mL fresh distilled water.
2. Weight 0, 3 g of NaCl. Introduce the substance to the inner vessel and dissolve it completely.
3. Determine the freezing point of the mixture as described part A, 13 but take measurement when
temperature of the system drops to 0˚C.
NOTES:

Don’t pour the ethanol in the jacket

Please clean the temperature probe

Clean apparatus at the end of the experiment

Be careful that ethanol is flammable in outer jacket
Questions:
1. Compare the Kf and Van’t Hoff factor (i) values with the theoretical ones.
2. Explain the reason why supercooling is observed in the experiment
3. Why the ethanol is put into the jacket and protective sleeve?
4. Why the decrease in the freezing point is observed, explain it in detail.
5. Why the temperature decrease in NaCl is more than urea?
DATA SHEET
Experiment 6-Determination of Molar Mass and van’t
Factor by Freezing Point Depression
Group Number:
I.
Hoff
Date:
Assistant name and signature:
Determination of freezing point constant, Kf of the solvent.
Table 2. Variation of temperature with time
Pure water
Time(s)
T(˚C)
water + 0.30 g urea
water + 0.60 g urea
water + 0.90 g urea
Time(s)
Time(s)
Time(s)
T(˚C)
T(˚C)
30
30
30
30
60
60
60
60
90
90
90
90
120
120
120
120
150
150
150
150
180
180
180
180
210
210
210
210
240
240
240
240
270
270
270
270
300
300
300
300
330
330
330
330
360
360
360
360
390
390
390
390
420
420
420
420
1. Mass of solvent =
2. Plot cooling curves.
T(˚C)
3. Molality of solution water + 0.30 g of urea =
Molality of solution water + 0.60 g of urea =
Molality of solution water + 0.90 g of urea =
4. Fill the following Table 3.
Table 3
Tf
Tf
molality
Pure water
water + 0.30 g of urea
water + 0.60 g of urea
water + 0.90 g of urea
5. Plot Tf versus molality. Determine the cryoscopic constant Kf of the solvent.
Kf =
II. Determination of molar mass of an unknown solute
Table 4. Variation of temperature with time
water + 0.3 g
unknown
Time(s)
1. Plot cooling curve.
T(˚C)
30
60
90
120
150
180
210
240
270
300
330
360
390
420
2. Fill the following Table 5.
Table 5
Tf
Pure water
water + 0.30 g of unknown solute
Molar mass=
Tf
(KfmB) /mA
III. Determination of van’t Hoff factor of a strong electrolyte
Table 3. Variation of temperature with time
water + 0.30 g
NaCl
Time(s)
o
T( C)
30
60
90
120
150
180
210
240
270
300
330
360
390
420
1. Plot cooling curve.
2. Fill the following Table
Tf
Pure water
water + 0.30 g of NaCl
Tf
molality
van’t Hoff factor