Chapter 8. Laboratory: Colligative Properties of Solutions

Vapor pressure reduction occurs as the number of solute particles in a solution increases. In effect, some solute particles displace solvent particles at the surface of the solution (the liquid-gas boundary), making fewer solvent particles available for vaporization and thereby reducing the vapor pressure.
Boiling point elevation is sometimes considered a separate colligative property, but is in fact a special case of vapor pressure reduction. The boiling point is reached at the temperature where the vapor pressure of the liquid matches the pressure of the gas at the liquid-gas boundary. Because the presence of solute particles reduces the vapor pressure of the solution, it also increases the boiling point of the solution. Boiling point elevation (ΔT_b or ΔT_bp) is the product of the van’t Hoff factor (i) of the solute, the boiling point elevation constant (also called the ebullioscopic constant) of the solvent (K_b or K_bp), and the molality (m) of the solution, and can be expressed as the formula ΔT = iK_bm.
One familiar application of vapor pressure reduction is the use of antifreeze (which could just as easily be called anti-boil) in automobile radiators. The addition of high-boiling ethylene glycol and similar chemicals to the water in the radiator raises the boiling point of the solution above that of pure water, reducing the likelihood of the radiator boiling over.

Freezing point depression occurs as the number of solute particles in a solution increases. One familiar application of freezing point depression is the application of road salt (usually sodium chloride or calcium chloride) to melt ice on streets and sidewalks. As the ice dissolves the salt (yes, one solid can dissolve another solid) the solid ice is converted to a liquid solution of the salt, because the solution has a lower freezing point than the essentially pure water that makes up the ice. Freezing point depression is calculated in the same way as boiling point elevation, but substituting the freezing point depression constant (K_f or K_fp) for the boiling point elevation constant. Because the freezing point depression constant is conventionally expressed as an unsigned value, the formula for calculating freezing point depression, ΔT = –(iK_fm), adds a negative sign to indicate the reduction in freezing point.

Osmotic pressure occurs when a differential concentration of solute particles causes pressure to be exerted across a semipermeable membrane. The phenomenon of osmotic pressure is exploited in many chemical and industrial processes as well as in medical procedures such as kidney dialysis, not to mention such routine bodily functions as your kidneys extracting waste products from your bloodstream. The behavior of solutions under osmotic pressure resembles the behavior of gases. In fact, the formula for calculating osmotic pressure uses the ideal gas constant: π = (niRT)/V, where π is the osmotic pressure, n the moles of solute, i the van’t Hoff constant, R the ideal gas constant, T the absolute temperate in kelvins, and V the volume.

Colligative properties change with changing concentration of a solution. For most procedures, chemists use molarity (moles per liter of solution or mol/L) to specify concentration, but when working with colligative properties, molality (moles per kilogram of solvent or mol/kg) is a more useful metric, because molality can be determined with extreme accuracy using only an analytical balance and because molality does not change with increasing or decreasing volume as the solution is heated or cooled.

In this laboratory, we’ll examine all three colligative properties of solutions.

Laboratory 8.1: Determine Molar Mass by Boiling Point Elevation

Dissolving a nonvolatile solute in a solvent increases the boiling point of that solvent by an amount proportional to the quantity of the solute. Although this phenomenon occurs with any solvent and solute, for reasons of safety and economy we’ll use water as the solvent for this session, and sodium chloride (table salt) and sucrose (table sugar) as the solutes.

Part II: Prepare Molal Solutions of Sodium Chloride and Sucrose

To test the effect of molality and dissociation on boiling point, we need to prepare solutions of ionic and molecular (covalent) compounds of known molality. I chose sodium chloride and sucrose because both of these chemicals are inexpensive, readily available, and extremely soluble in water. Sodium chloride is ionic. In solution, sodium chloride dissociates into sodium ions (Na⁺) and chloride ions (Cl^–), and should therefore have a van’t Hoff factor of 2. Sucrose is molecular, and should therefore have a van’t Hoff factor of 1.

Coincidentally, the solubility of both sodium chloride and sucrose in water at room temperature is just over 6 mol/L. We’ll therefore prepare 6 molal samples of both of these compounds, and then by serial dilution prepare 3 molal and 1.5 molal samples of each.

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Use the 100 mL graduated cylinder to transfer 200 mL of water to the 250 mL beaker.
Weigh 70.13 g of sodium chloride, add it to the 250 mL beaker, and stir until the sodium chloride is completely dissolved. This solution contains 6 mol/kg of sodium chloride (6 molal).
Transfer 100 mL of the 6 molal sodium chloride solution to beaker A.
Add 100 mL of water to the 250 mL beaker to dilute the remaining 100 mL of 6 molal sodium chloride solution to 200 mL of 3 molal solution.
Transfer 100 mL of the 3 molal sodium chloride solution to beaker B.
Add 100 mL of water to the 250 mL beaker to dilute the remaining 100 mL of 3 molal sodium chloride solution to 200 mL of 1.5 molal solution.
Transfer 100 mL of the 1.5 molal sodium chloride solution to beaker C and discard the remainder. (Or transfer all 200 mL; the quantity of the solution has no effect on the boiling point elevation.)
Use the 100 mL graduated cylinder to transfer 200 mL of water to the 250 mL beaker.
Weigh 410.76 g of sucrose, add it to the 250 mL beaker, and stir until the sucrose is completely dissolved. This solution contains 6 mol/kg of sucrose (6 molal).
Transfer 100 mL of the 6 molal sucrose solution to beaker D.
Add 100 mL of water to the 250 mL beaker to dilute the remaining 100 mL of 6 molal sucrose solution to 200 mL of 3 molal solution.
Transfer 100 mL of the 3 molal sucrose solution to beaker E.
Add 100 mL of water to the 250 mL beaker to dilute the remaining 100 mL of 3 molal sucrose solution to 200 mL of 1.5 molal solution.
Transfer 100 mL of the 1.5 molal sucrose solution to beaker F and discard the remainder. (Or transfer all 200 mL; the quantity of the solution has no effect on the boiling point elevation.)

Part III: Determine the Boiling Points of Sodium Chloride and Sucrose Solutions

In this part of the lab, we determine the boiling points of the sodium chloride and sucrose solutions.

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Place beaker A on the heat source. Add one boiling chip, and apply heat with constant stirring until the contents of beaker A come to a full boil.
Immerse the thermometer in the beaker, making sure that it does not contact the beaker itself, allow the thermometer to stabilize, and record the temperature reading in Table 8-1. Place the beaker aside to cool.
Repeat steps 2 and 3 for beakers B, C, D, E, and F.

Figure 8-1. Determining the boiling point of a solution

Disposal

Retain all of the solutions for use in the following laboratory session. Allow all six beakers to cool to room temperature, and then place them in the refrigerator to chill them. A typical refrigerator maintains a temperature of about 5°C. Having the solutions at this temperature reduces the time required for the freezing point determinations in the following laboratory session. (I violate my general rule of avoiding mixing laboratory materials with kitchen materials because these beakers contain only sodium chloride and sucrose solutions, both of which are harmless.)

Table 8-1. Boiling point elevation—observed and calculated data

Beaker/solution	Boiling point	Calculated formula weight
Water	______.__°C
A. Sodium chloride, 6 mol/kg	______.__°C	_________.___ g/mol
B. Sodium chloride, 3 mol/kg	______.__°C	_________.___ g/mol
C. Sodium chloride, 1.5 mol/kg	______.__°C	_________.___ g/mol
D. Sucrose, 6 mol/kg	______.__°C	_________.___ g/mol
E. Sucrose, 3 mol/kg	______.__°C	_________.___ g/mol
F. Sucrose, 1.5 mol/kg	______.__°C	_________.___ g/mol

Review Questions

Q1: Name and define the unit of concentration used in calculating values for colligative properties of solutions. Why is this unit of concentration used rather than molarity?

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Q2: The K_bp for H₂O is 0.512°C/molal. Applying this value and the van’t Hoff factors for sodium chloride (2) and sucrose (1) to the boiling point elevation values you recorded in Table 8-1, calculate the formula weights of sodium chloride and sucrose. Record these calculated values in Table 8-1.

Q3: What effect on boiling point would you expect if you dissolved sufficient ethanol in water to produce a 1 molal solution?

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Laboratory 8.2: Determine Molar Mass by Freezing Point Depression

Just as a dissolved solute increases the boiling point of a solvent by an amount proportional to the quantity of the solute, it also decreases the freezing point of that solvent. The same formula is used to calculate either value, except that a negative sign is added for freezing point depression to account for the fact that the freezing point is lowered rather than raised.

Boiling point elevation = ΔT_b = iK_bm
Freezing point depression = ΔT_f = –(iK_fm)

In fact, some sources specify freezing point depression constants, K_f or K_fp, as negative signed values, eliminating the need to reverse the sign in the formula. In either case, the calculations are done the same way for boiling point elevation and freezing point depression. Only the values of the constants differ. For example, the K_b for water is 0.512°C/molal and the K_f for water is 1.858°C/molal. In other words, assuming a van’t Hoff factor of 1 for the solute, a 1 molal solution has a boiling point 0.512°C higher than pure water, and a freezing point 1.858°C lower. A 2 molal solution has a boiling point 1.024°C higher than pure water, and a freezing point 3.716°C lower. And so on.

In this laboratory, we’ll use the sodium chloride and sucrose solutions prepared in the preceding lab to test the effect of dissolved solutes on freezing points.

Procedure

This lab has two parts. In Part I, we’ll prepare an ice/salt bath and determine the freezing point of water. In Part II, we’ll determine the freezing points of the solutions we prepared in the preceding laboratory.

Part I: Prepare the Ice Bath and Determine the Freezing Point of Water

Unlike the boiling point of water, which varies significantly with ambient local pressure, the freezing point of water is almost unaffected by pressure. But thermometers vary in accuracy, so before we test freezing point depressions in the sodium chloride and sucrose solutions, we’ll calibrate our thermometer by using it to determine the freezing point of water. (For example, if our thermometer indicates a freezing point for pure water of +0.5°C and we subsequently measure a depressed freezing point at –2.5°C, the actual freezing point depression is 3.0°C rather than 2.5°C.)

Caution

Wear splash goggles, gloves, and protective clothing.

Because the presence of a solute reduces the freezing point of an aqueous solution below 0°C, the freezing point of pure water, we need a way to cool the solutions below 0°C. In an elegant application of using freezing point depression to test freezing point depression, we’ll use the phenomenon itself to provide the conditions necessary for the test. We’ll make an ice/salt bath by mixing crushed ice with sodium chloride. The phenomenon of freezing point depression means that the temperature of this ice/salt bath will be lower than the temperature of a pure ice bath.

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Half fill the large beaker with crushed ice, pour a 15 mm to 25 mm layer of sodium chloride on top of the ice, and stir until the ice and salt are thoroughly mixed.
Immerse the thermometer in the ice/salt bath, allow a minute or so for it to stabilize, and record the temperature of the ice/salt bath in Table 8-2.
Half fill a test tube with water and carefully press it down into the ice bath until the water is completely below the surface of the ice bath. See Figure 8-2.
Gently stir the water with the thermometer continuously until ice crystals begin to form on the side of the test tube. You will probably have to remove the test tube from the ice bath periodically to check the status. Ice crystals tend to form first near the bottom of the tube.
When ice crystals begin to form, record the temperature of the water in Table 8-2.

Disposal

Dispose of the solutions in the test tubes by flushing them down the drain. Allow the beakers with the remainder of the six solutions to come to room temperature. These solutions will be used in the following laboratory session.

Part II: Determine the Freezing Points of Sodium Chloride and Sucrose Solutions

In this part of the lab, we determine the freezing points of the sodium chloride and sucrose solutions.

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Label six test tubes A through F.
Transfer approximately 10 mL (about half a test tube) of solution from beaker A to test tube A.
Repeat step 3 for the solutions in beakers B through F.
Carefully press test tube A down into the ice bath until the solution is completely below the surface of the ice bath.
Gently and continuously stir the solution with the thermometer until ice crystals begin to form on the side of the test tube. You will probably have to remove the test tube from the ice bath periodically to check the status. Ice crystals tend to form first near the bottom of the tube.
When ice crystals begin to form, record the temperature of the solution in Table 8-2. If no ice crystals appear within a few minutes, note that fact in Table 8-2, record the lowest temperature reached for that solution, and go on to the next solution.
Repeat steps 5 through 7 for test tubes B through F.

Table 8-2. Freezing point depression—observed and calculated data

Test tube/solution	Freezing point	Calculated formula weight
Ice/salt bath	______.__°C
Water	______.__°C
A. Sodium chloride, 6 mol/kg	______.__°C	_________.___ g/mol
B. Sodium chloride, 3 mol/kg	______.__°C	_________.___ g/mol
C. Sodium chloride, 1.5 mol/kg	______.__°C	_________.___ g/mol
D. Sucrose, 6 mol/kg	______.__°C	_________.___ g/mol
E. Sucrose, 3 mol/kg	______.__°C	_________.___ g/mol
F. Sucrose, 1.5 mol/kg	______.__°C	_________.___ g/mol

Figure 8-2. Determining the freezing point of a solution

Review Questions

Q:	Q1: Did all of the solutions that you tested freeze? If not, propose and quantify an explanation for this observation. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________
Q:	Q2: In the early days of automobiles, methanol was used universally as antifreeze. In 1937, antifreeze solutions based on ethylene glycol were introduced, and quickly gained almost 100% market share, despite the fact that they cost much more than methanol. Propose an explanation for this rapid adoption of ethylene glycol antifreeze solutions. (Hint: look up the physical properties of methanol and ethylene glycol.) __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________
Q:	Q3: The K_fp for H₂O is 1.858°C/molal. Applying this value and the van’t Hoff factors for sodium chloride (2) and sucrose (1) to the freezing point depression values that you recorded in Table 8-2, calculate the formula weights of sodium chloride and sucrose. Record these calculated values in Table 8-2.
Q:	Q4: What effect on freezing point would you expect if you dissolved sufficient ethanol in water to produce a 1 molal solution? __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________

Laboratory 8.3: Observe the Effects of Osmotic Pressure

Osmotic pressure is the hydrostatic pressure that exists when two solutions of differing concentrations are in contact with the two sides of a semipermeable membrane. The semipermeable membrane allows small molecules (such as water) to pass through it freely, but physically prevents larger molecules (such as a solute) from passing through the membrane. In such a system, water passes from the side that contains the less-concentrated solution to the side with the more-concentrated solution until the concentrations of the solutions are equal, at which point the system is in equilibrium.

If a biological cell is placed in a hypotonic environment (one in which the solution inside the cell is more concentrated than the solution surrounding the cell), water flows into the cell from the outside solution through the semipermeable membrane that surrounds the cell, causing the cell to expand and gain mass. Conversely, in a hypertonic environment (one in which the outside solution is more concentrated), water flows outward from the cell, shrinking it and reducing its mass.

Children sometimes unwittingly illustrate the principle of osmotic pressure by pouring table salt on a garden slug. This creates an extreme hypertonic environment, sucks water out of the slug, and thereby kills it by quick and extreme dehydration. My editor told me to make sure that no slugs were harmed in running this experiment, so I decided to use celery instead. If that offends People for the Ethical Treatment of Vegetables (http://www.petv.org), I’m sorry, but I had to use something.

In this laboratory, we’ll use the sodium chloride and sucrose solutions left over from the first laboratory session in this chapter to observe the effects of osmotic pressure on the semipermeable membrane that surrounds celery cells.

Disposal

Dispose of the solutions by flushing them down the drain.

Caution

Wear splash goggles, gloves, and protective clothing. Be careful with the knife.

Procedure

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Cut the celery stalk into seven pieces of approximately equal mass.
Label the celery pieces A through G.
Weigh each piece of celery and record the mass of each in Table 8-3.
Drop each piece of celery, A through F, into the corresponding beaker, making sure that the celery is completely submerged.
Drop celery piece G into a container of pure water.
Allow the celery pieces to soak for at least one hour.
Remove each piece of celery from its beaker, rinse it briefly in running water, pat it dry with a paper towel, and reweigh it (Figure 8-3). Record the mass of each piece in Table 8-3.

Table 8-3. Osmotic pressure—observed and calculated data

Celery piece/solution % mass gain (loss)	Initial mass	Final mass	Mass gain (loss)
A. Sodium chloride, 6 mol/kg	________.____ g	________.____ g	_________.____ g
B. Sodium chloride, 3 mol/kg	________.____ g	________.____ g	_________.____ g
C. Sodium chloride, 1.5 mol/kg	________.____ g	________.____ g	_________.____ g
D. Sucrose, 6 mol/kg	________.____ g	________.____ g	_________.____ g
E. Sucrose, 3 mol/kg	________.____ g	________.____ g	_________.____ g
F. Sucrose, 1.5 mol/kg	________.____ g	________.____ g	_________.____ g
G. Water	________.____ g	________.____ g	_________.____ g

Figure 8-3. Determining the mass of a celery sample

Review Questions

Q1: Calculate the mass gain (or loss) for each of the samples, in grams and percentage, and enter the results in Table 8-3. Note which samples gained mass and which samples lost mass. Propose an explanation.

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Q2: If you submerged a celery sample in a 70% ethanol solution, would you expect that sample to gain or lose mass? Why?

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Q3: Using only the materials and equipment listed in this chapter, propose an experimental method to accurately determine the concentration of the solution in the celery cells.

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	goggles, gloves, and protective clothing
	balance and weighing paper
	beaker, 150 mL (6)
	beaker, 250 mL (1)
	graduated cylinder, 100 mL
	alcohol lamp or other heat source
	tripod stand and wire gauze
	beaker tongs
	stirring rod
	thermometer
	boiling chips
	sodium chloride (70.13 g)
	sucrose (410.76 g)
	tap water

	goggles, gloves, and protective clothing
	beaker, 600 mL (1)
	test tubes (6)
	thermometer
	stirring rod
	chilled solutions A through F from preceding lab
	chipped ice sufficient to half fill beaker
	sodium chloride sufficient to make a 15 mm to 25 mm layer on top of crushed ice

	goggles, gloves, and protective clothing
	balance
	beaker, 150 mL (1)
	marking pen
	sharp knife
	paper towels
	solutions A through F from Laboratory 8.1: Determine Molar Mass by Boiling Point Elevation
	celery stalk