Introduction:

Solutions are homogeneous mixtures that contain two or more substances. The major component is called the solvent, and the minor component(s) are called the solute(s). Since the solution is primarily composed of solvent, the physical properties of the solution resemble those of the pure solvent. However, some of these physical properties, called colligative properties, are independent of the nature of the solute and depend only upon the concentration of solute particles. Examples of colligative properties include vapor pressure reduction, boiling point elevation, freezing point depression, and osmotic pressure.

Pure water freezes at 0.0 °C (273 K), boils at 100.0 °C (373 K), and exerts a vapor pressure of 23.76 mm Hg at 25.0 °C (298 K). These values are altered by the presence of a solute. Your are probably familiar with some common examples of these effects: Antifreeze is used to lower the freezing point and raise the boiling point of your engine coolant (water); and salt is used to melt ice on the roads.

These effects are expressed quantitatively by the colligative-property law, which states that the freezing point and boiling point of a solution differ from those of the pure solvent by amounts that are directly proportional to the molal concentration of the solute. In its general form this equation is written:

where DT is the boiling point elevation or freezing point depression, K is a constant that is specific for each solvent, and m is the molality of the solution and is expressed as the number of moles of solute per kilogram of solvent. Some representative constants, boiling points, and freezing points are given in the table below.  

The equations relating the boiling point elevation and freezing point depression to molality are:

where DT_bp and DT_fp are the change in boiling point and freezing point respectively, K_bp and K_fp are the boiling and freezing constants for a particular solvent, and i is the van't Hoff factor which is used to account for the number of particles a given solute generates in solution. Substances are usually classified as ionic or covalent, or more precisely, electrolytes and non electrolytes. Electrolytes dissolve in a solvent to give ions (charged particles) in solution while non electrolytes dissolve to give molecules. A mole of non electrolyte such as ethanol would dissolve to produce a mole of ethanol molecules. However, a mole of sodium chloride (NaCl) would dissolve to form two moles of ions (Na⁺ and Cl^-). Because colligative properties are related to the number of solute particles, we would expect a mole of NaCl to have twice the effect as a mole of ethanol. Therefore, (i =1) for ethanol and other non electrolytes, (i =2) for NaCl, and (i = 5) for Al₂(CO₃)₃.

As an example, lets say that we dissolve 10g of ethanol, C₂H₅OH, in 100 g of pure water. What would be the boiling point, and freezing point of this solution? First we must determine the molality of the solution. To do this we need to know the moles of solute and kilograms (Kg) of solvent:

For water the K_bp = 0.52°C/m, and K_fp = 1.86°C/m. So for this ethanol solution,

the boiling point is increased by 1.13°C to 101.13°C. Likewise, the freezing point

is decreased by 4.04°C to –4.04°C.  It is very important to not confuse the actual freezing point with the freezing point depression!

In the above example, we assumed that the van't Hoff factor for ethanol had a value of one (i=1).  However, we know this is only true for very dilute solutions.  Lets assume that the same ethanol/water solution had an experimentally determined freezing point of -3.75°C.  We could calculate the van't Hoff factor as follows:

You can also see that if the molal freezing point depression constant is known, it is possible to calculate the molecular weight of a solute by measuring the freezing point of a solution that contains a known amount of solvent and solute. Lets say you were attempting to determine the molecular weight of urea (CH₄N₂O) using the freezing point depression of naphthalene (C₁₀H₈). During the experiment, you found that a solution containing 15 g of urea in 100 g of naphthalene had a freezing point of 63.3°C. From the table above, we see that the freezing point of pure naphthalene is 80.6°C.  Therefore, DT = 17.3°C. We will also assume, since urea is a non-ionic solid, that i = 1.   From the above table, we know that the K_fp for naphthalene is 6.9°C/m.  Therefore, we can calculate the molality of our solution:

Remember that molality is defined as the number of moles of solute per 1000 g of solvent. However, in this example, we only used 100 g of solvent (naphthalene), and so the moles of urea are:

Since moles can be calculated by dividing the grams of a substance by the molecular weight of that substance, we can now calculate the molecular weight of urea:

Thus the molecular weight of urea is 59.8 g/mol.

Purpose:

The first purpose of this lab is to experimentally determine the van't Hoff (i) factor for two different substances, sucrose (C₁₂H₂₂O₁₁) and sodium chloride (NaCl). From the discussion above you should be able to ascertain these values. However, these theoretical values are only valid for very dilute solutions (<0.01 m). At higher concentrations, the ions do not behave independently and so the values of (i) will be less than theoretical.

The second purpose of this lab is to determine the molecular weight of commercial antifreeze.

Method:

In this experiment you will need to determine the freezing point depression for several solutions. However, the temperature at which a solution freezes is difficult to determine by direct visual observation because of a phenomenon called super cooling and also because solidification of solutions usually occurs over a broad temperature range. To overcome these problems, temperature-time graphs, called cooling curves, are used which readily reveal the true freezing temperature. The major features of a general cooling curves are shown below:

Procedure:

Preparation of an Ice Bath:

1. Fill a large Syrofoam cup with ice.
2. Add approximately 150 mL of distilled water to the cup.
3. Add approximately 25 g of salt to the cup.
4. Stir this ice-salt mixture with a scoopula until a temperature range of -5 °C to -10 °C is obtained.  Note:  colder temperatures are easily obtained by adding more salt, however, it makes the solutions cool too fast to obtain good results.
5. Be sure to check the temperature of the ice bath BEFORE each experiment.  If the temperature of the bath rises above -5 °C, then drain the water and add more ice and salt.

1. Dissolve approximately 3.1 g of commercial antifreeze in 50 g of water (weigh both accurately).  Mix until both liquids have completely dissolved.
2. Be sure to use a clean and dry the test tube and add approximately 15 ml of the antifreeze solution to it.
3. Place the test tube, thermometer and agitator into the ice bath so that the level of the water in the test tube is below the level of the ice. Make sure you can still read the thermometer.
4. Record the temperature of the antifreeze solution every 30 seconds while vigorously agitating.
5. Continue to record the temperature until the solution solidifies.  Note:  these solutions tend to form a slush and is difficult to observe the exact freezing point.  Keep collecting data until you obtain 5 - 10 readings below the freezing point of water.
6. Remove the test tube and put the contents in the marked antifreeze waste bottle in the hood.  You will need to rinse, and dry the test tube between trails.
7. Repeat the cooling curve a second time.  Note:  Dispose of any left over antifreeze solution in the marked antifreeze waste bottle in the hood.

1. Prepare a solution of sucrose by adding approximately 17 g of sucrose (C₁₂H₂₂O₁₁) to 50 g of water (weigh both accurately). Mix until all crystals have dissolved.
2. Be sure your large test tube is clean and dry and add approximately 15 ml of the sucrose solution to it.
3. Place the test tube, thermometer and agitator into the ice bath so that the level of the water in the test tube is below the level of the ice. Make sure you can still read the thermometer.
4. Record the temperature of the sucrose solution every 30 seconds while vigorously agitating.
5. Continue to record the temperature until the solution solidifies.  Note:  these solutions tend to form a slush and is difficult to observe the exact freezing point.  Keep collecting data until you obtain 5 - 10 readings below the freezing point of water.
6. Remove the test tube and dump the contents down the drain.  You will need to rinse, and dry the test tube between trails.
7. Repeat the cooling curve a second time.

1. Prepare a solution of NaCl by adding approximately 2.9 g of table salt to 50 g of water (weigh both accurately). Mix until all crystals have dissolved.
2. Be sure to use a clean and dry the test tube and add approximately 15 ml of the sodium chloride solution to it.
3. Place the test tube, thermometer and agitator into the ice bath so that the level of the water in the test tube is below the level of the ice. Make sure you can still read the thermometer.
4. Record the temperature of the saline solution every 30 seconds while vigorously agitating.
5. Continue to record the temperature until the solution solidifies.  Note:  these solutions tend to form a slush and is difficult to observe the exact freezing point.  Keep collecting data until you obtain 5 - 10 readings below the freezing point of water.
6. Remove the test tube and dump the contents down the drain.  You will need to rinse, and dry the test tube between trails.
7. Repeat the cooling curve a second time.

1. Be sure to use a clean and dry the test tube and add approximately 15 ml of distilled water to it.
2. Place the test tube, thermometer and agitator into the ice bath so that the level of the water in the test tube is below the level of the ice.  Make sure you can still read the thermometer.  You need the agitator to keep ice from forming on the edges of the test tube.
3. Record the temperature of the water every 30 seconds while vigorously agitating.
4. Continue to record the temperature until the water solidifies.  Note:  it is difficult to determine the freezing point of pure water.  Keep collecting data until you have at least 5 reading at the same temperature.
5. Remove the test tube and dump the ice down the drain.  You will need to rinse and dry the test tube between trails.
6. Repeat the cooling curve a second time.

Results:

Use the following table as a guide for your data collection:  

Remember, to allow enough space in you lab notebooks to record all of the time and temperature data form these 8 experiments.

Calculations:

From your cooling curves, calculate:

• The freezing point of pure water
• The freezing point of the sucrose solution
• The freezing point of the sodium chloride solution
• The freezing point of the antifreeze solution

• The effective (i) for the sucrose solution
• The effective (i) for the sodium chloride solution
• The molecular weight of the antifreeze

• Report the average (i) value obtained for the sucrose solution
• Report the average (i) value obtained for the sodium chloride solution
• Report the average molecular weight for the commercial antifreeze
• What values of (i) did you expect for sucrose and sodium chloride? Did your experiment confirm this? Why?
• What was the effect of more or less stirring on the super cooling portion of the cooling curve?

(Updated 6/5/07 by C.R. Snelling)