Materials Included In Kit

Additional Materials Required

Safety Precautions

Aluminum chloride and calcium chloride are slightly toxic by ingestion. Sodium chloride and sucrose are not considered hazardous; however, the chemicals provided are for laboratory use only and are not intended for human consumption. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulation that may apply, before proceeding. Dispose of the ice-water mixtures by pouring them down the drain with plenty of water according to Flinn Suggested Disposal Method 26b.

Teacher Tips

• There are enough materials for a class of 30 students or 15 groups of students working in pairs.
• The lab works significantly better with crushed ice rather than with ice cubes. If crushed ice is not available, place some ice cubes in two layers of zipper-lock bags. Using caution, pound the ice with a hammer or other hard object to crush it into small pieces.
• Have each student group use the same thermometer throughout the lab to reduce errors resulting from inaccurate thermometers.
• Distilled or deionized water will produce the best results; however, tap water can also be used as the source of water and ice in this lab. Any ions in the water will be present in every mixture, including the control, and thus will not affect comparative results. The concentration of ions in the local tap water, however, may cause the results to vary slightly from the sample data provided.
• Remind the students to use the wooden stirrer to stir the solution and the thermometer strictly to measure the temperature. Also caution students not to leave thermometers standing in beakers unattended as this is a common cause of broken thermometers and spilled solutions.
• The temperature of the ice/water mixture in step 5 should be 0 °C. This step is a good calibration check to measure the accuracy of the thermometer. If the temperature does not read 0 °C, correct for this error in all further thermometer readings.
• Consider providing students with an unknown salt and having them repeat this freezing point depression lab. Have them determine n, the number of particles or ions, from each unit of unknown. Provide a list of possible unknowns.
• Demonstrate the change in freezing point depression with increasing amounts of the same solute—use 10 g, 20 g and 30 g of sodium chloride in three different ice-water mixtures. Compare ΔT_f values. The more additive added, the greater the ΔT_f values.

Sample Data

1. Determine the formula weight of each solute. Record this in the Analysis Table.
   Formula weights can be found on the chemical label on the bottle, or they can be calculated using the atomic weights from the periodic table. For example, the formula weight of NaCl = 23.00 g/mol + 35.45 g/mol = 58.45 g/mol.
2. Calculate the freezing point depression, ΔT_f, using the following equation. Record values for ΔT_f in the Analysis Table.
   

   T_f (pure solvent) – T_f (solution) = ΔT_f (in °C)

   T_f(H₂O) – T_f (NaCl solution) = 0.0 °C – (–16.7 °C) = 16.7 °C = ΔT_f
3. Calculate the number of moles used for each additive using the exact mass of additive from the Data Table and the formula weights. Record values in the Analysis Table.

4. Calculate the ΔT_f per mole of solute. Record values in the Analysis Table.
   {11912_Data_Equation_2}
5. Determine n for each solute. This is the number of particles or ions formed from the dissociation of each formula unit. Write a balanced equation for the dissociation of each solute. Record n for each solute in the Analysis Table.

   NaCl(s) → Na⁺(aq) + Cl^–(aq) n = 2
   C₁₂H₂₂O₁₁(s) → C₁₂H₂₂O₁₁(aq) n = 1
   CaCl₂(s) → Ca²⁺(aq) + 2Cl^–(aq) n = 3
   AlCl₃(s) → Al³⁺(aq) + 3Cl^–(aq) n = 4

6. Calculate the concentration of each solution in molality, m, which equals moles of solute per kilogram of solvent as shown in the equation below. Remember to use exact measured values for mass of solvent from the Data Table. Record values for m in the Analysis Table.

   Molality = m = moles solute/kg solvent

   {11912_Data_Equation_3}
7. Calculate the theoretical value for the freezing point depression for each solute. This is the freezing point depression that should have been observed using the starting amount of solute and solvent. Use Equation 1 from the Background section. Use K_f = 1.86 °C/m, and calculated values for m and n. Record values for ΔT_f (Theoretical) in the Analysis Table.

   ΔT_f (theoretical) = K_f x m x n
   ΔT_f = 1.86 °C/m x 5.17 m x 2 = 19.2 °C for NaCl
   ΔT_f = 1.86 °C/m x 0.872 m x 1 = 1.62 °C for C₁₂H₂₂O₁₁
   ΔT_f = 1.86 °C/m x 2.05 m x 3 = 11.4 °C for CaCl₂
   ΔT_f = 1.86 °C/m x 1.24 m x 4 = 9.23 °C AlCl₃

8. Determine the percent error for the ΔT_f for each solute. Use the following equation. Record values in the Analysis Table.
   {11912_Data_Equation_4}
9. Prepare a graph of n versus ΔT_f /mole. Describe the relationship.

   The graph is linear showing a directly proportional relationship between n and ΔT_f/mole.

   {11912_Data_Figure_1}

Answers to Questions

1. Make a general statement regarding the effect of additives or impurities on the melting (or freezing) point of a pure substance.

   An additive or impurity lowers the freezing point of a substance; in other words, a solution always has a lower freezing point than its pure solvent.

2. Why does an impurity (such as a salt) have the effect of lowering the freezing point of a solvent?

   Salt and other dissolved impurities interfere with the ability of the solvent to crystallize (solidify) and the solution remains liquid even at a lower temperature.

3. Using the same mass of additive (i.e., 30 g), which additive to the ice water lowered the freezing point the most (with the greatest ΔT_f)?

   The NaCl lowered the freezing point the most.

4. Using the same mass of additive (i.e. 30 g), which additive to the ice water lowered the freezing point the least (with the lowest ΔT_f)?

   The sucrose lowered the freezing point the least.

5. Which additive has the greatest freezing point depression per mole? Which has the least? Is this what would be expected? Explain.

   Aluminum chloride has the greatest freezing point depression per mole and sucrose has the lowest. This is expected because each mole of aluminum chloride dissociates into four particles (ions) in solution while sucrose remains as one single particle when in solution, and freezing point depression (a colligative property) depends on the number of particles or ions in solution.

6. What should be true about the freezing point depression per particle or ion? Does your data verify this?

   The values for freezing point depression per particle or ion should be the same for any solute in a given solvent. Yes, the data verifies this within experimental error.

7. Percent error was calculated by comparing the freezing point that should have been observed for the given amount of each solute with that actually observed. Discuss possible sources of error in the lab. How might the errors have been prevented?

   (1) The ions or impurities in tap water and in the crushed ice may have affected the freezing point. To prevent this, distilled or deionized water could have been used to make the ice. (2) Solutes dissolve slower at low temperatures; temperature readings may have been taken before all of the solid was dissolved. (3) Spillage of the solute or the solvent may have occurred. (4) The thermometer may be out of calibration. Correct for this by measuring the temperature of pure ice water and subtracting this out of each measurement.

8. What factors were held constant in this experiment?

   Same thermometer throughout; same amount of ice-water mixture; same mass of solute; solutions were stirred equally throughout.

9. Why was it important to keep the thermometer off of the bottom of the beaker?

   For a most accurate reading, the thermometer bulb must be surrounded on all sides by the solid-liquid mixture.

10. Why was it necessary to measure the temperature of the pure ice water mixture, instead of assuming it to be 0.0 °C?

    The most important reason for determining the freezing point of water is to calibrate the thermometer.

11. Which would be a better de-icer—sodium chloride or potassium chloride? Why?

    Sodium chloride. Both have same n value—two particles per mole. Sodium chloride has a lower formula weight, so less grams are needed to get the same number of particles.

12. Given the following sample cost data, which de-icing chemical would you recommend a road crew using as the most effective and most cost-effective agent at preventing road icing?
    {11912_Answers_Table_1}

    Although aluminum chloride forms the most particles per mole of compound, it is the least cost-effective. Sodium chloride is the most effective choice for the cost to use as a de-icing chemical because of costs/kg and moles/kg.

Teacher Handouts

11912_Teacher1.pdf

Introduction

Determine the effect of dissolved impurities on the freezing point of a substance and determine which additive has the greatest effect on the freezing point.

Concepts

• Freezing/Melting
• Freezing point depression

Background

Inhabitants of northern states are familiar with winter and the snowy, icy roads that go with the season. Road crews spread salt (sodium chloride, calcium chloride, or a variety of salt mixtures) on the road in order to lower the temperature at which freezing occurs. Thus, the surface of the treated road does not freeze until the temperature gets down to –10 or –20 °C. If the road already has ice on it, the salt helps to melt the ice, forming a solution with a lower freezing point than that of pure water.

The freezing point of a liquid is the temperature at which the forces of attraction among molecules are just great enough to cause a phase change from the liquid state to the solid state. Strictly speaking, the freezing (or melting) point of a substance is the temperature at which the liquid and solid phases are in equilibrium.

During the freezing process of water, for example, water molecules come together to form the more orderly, crystalline pattern of ice molecules. When any solute (such as salt) is added to a solvent (such as water), the pattern is interrupted by the presence of the salt “impurity.” Salt and other dissolved impurities interfere with the ability of the solvent to crystallize (solidify) and the solution remains liquid even at a lower temperature. Thus a solution always has a lower freezing point than its pure solvent. This phenomenon is termed freezing point depression.

Freezing point depression (ΔT_f) is defined as the difference in temperature between the freezing point of a solution and that of its pure solvent. The freezing point of the solution after the addition of a solute can be calculated using Equation 1.

Here, ΔT_f is the change in freezing point (the freezing point depression) in °C, K_f is the freezing point depression constant (1.86 °C/m for water solutions), m is the molality of the solution (the solution concentration in moles of solute per kilogram of water), and n is the number of ions (particles) formed from the dissociation of each formula unit of solute.

Freezing point depression is a colligative property of a solution. A colligative property is directly dependent on the concentration of the solution—that is, the number of solute particles that are formed when the material is put into water—where concentration is expressed in molality. It is not dependent on the size or the identity of the particles, as is a common misconception.

Certain substances (solutes) lower the freezing point more than other substances. When a molecular substance such as sucrose (C₁₂H₂₂O₁₁) or sand (SiO₂) is placed into water, the molecule does not dissociate and remains as just one particle. Ionic solutes, on the other hand, dissociate into ions when put into water. That is, one unit of an ionic salt, such as sodium chloride (NaCl), dissociates in water to produce two particles—one sodium ion (Na⁺) and one chloride ion (Cl^–). One unit of calcium chloride (CaCl₂) when placed in water dissociates into three particles—one calcium ion (Ca²⁺) and two chloride ions (Cl^–). Looking at Equation 1, it can be seen that the freezing point depression depends on the number of particles in solution—the more particles in solution, the greater the change in freezing point.

Experiment Overview

In this lab, the goals are to measure the freezing point of pure water and measure the freezing point depression for each solution to determine the effect of dissolved impurities on the freezing point of a substance.

Materials

Prelab Questions

• Temperature should be measured in °C and will be accurate only if both water and ice are present in the beaker, indicating that the mixture is at the melting/freezing point.
• Hold the thermometer slightly off the bottom of the beaker when reading temperature so that the thermometer bulb is surrounded on all sides by the mixture.
• Stir only with the wooden stirrers and not with the thermometers. Hold the thermometer in the beaker with one hand to read the temperature and stir with the wooden stirrer with the other hand.

Safety Precautions

Aluminum chloride and calcium chloride are slightly toxic by ingestion. Sodium chloride and sucrose are not considered hazardous; however, the chemicals provided are for laboratory use only and are not intended for human consumption. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Procedure

1. Label four 250-mL beakers with numbers 1–4.
2. Place Beaker 1 on the balance and tare the balance, if the balance is electronic.
3. Add 100 g of ice water to the beaker by (a) first adding approximately 70–80 g of crushed ice and (b) then adding enough water so the total mass of ice plus water is 100.0 g. Record the precise mass of the ice-water mixture to the nearest tenth of a gram in the data table. (Note: If the balance does not have a capacity large enough, weigh the ice in a small weighing dish and then place it in the beaker. Measure the remaining water using a balance or graduated cylinder, where d = 1.0 g/mL for water.)
4. Stir the ice-water mixture with a wooden stirrer.
5. Carefully insert a thermometer into the ice-water mixture. Wait for the temperature reading to stabilize. Record the temperature of the pure ice-water mixture in °C to the nearest tenth of a degree on the data sheet in the space provided above the table.

Beaker 1—Sodium Chloride

6. In a weighing dish, weigh out 30.0 g of sodium chloride. Record the precise mass of the sodium chloride in the data table.
7. Add the sodium chloride to the ice-water mixture in Beaker 1.
8. Stir the contents of the beaker with a wooden stirrer until the mixture has a slushy appearance.
9. Carefully insert a thermometer into the mixture and measure the temperature. Continue to stir the mixture with the wooden stirrer. Record in the data table the lowest temperature that the mixture reaches before rising again, recording temperature in °C to the nearest tenth of a degree. This may take some time, as the salt does not immediately dissolve in the ice water.

Beaker 2—Sucrose

10. Repeat steps 2–9 for Beaker 2, using 100.0 grams of ice water and 30.0 grams of sucrose. Remember to record precise masses and temperature readings to the nearest tenth of a degree.

Beaker 3—Calcium Chloride

11. Repeat steps 2–9 for Beaker 3, using 100.0 grams of ice water and 30.0 grams of calcium chloride. Remember to record precise masses and temperature readings to the nearest tenth of a degree.

Beaker 4—Aluminum Chloride

12. Repeat steps 2–9 for Beaker 4, using 100.0 g of ice water and 30.0 grams of aluminum chloride. Remember to record precise masses and temperature readings to the nearest tenth of a degree.
13. Dispose of the solutions by pouring the mixtures down the drain with plenty of water. Rinse the beakers with tap water.
14. Compile class results for each mixture on the board. Calculate the class average freezing temperature for each mixture. Record the class average freezing point values in °C in the data table.
15. Complete the post-lab calculations 1–9 and answer the Post-Lab Questions 1–11.

Calculations 

Complete the following calculations. Show all work on a separate sheet of paper. Fill in values on the attached Data Analysis Table. 

1. Determine the formula weight of each solute. Record this in the Analysis Table.
2. Calculate the freezing point depression, ΔT_f, using the equation below. Record values for ΔT_f in the Analysis Table.

T_f (pure solvent) – T_f (solution) = ΔT_f (in °C)

3. Calculate the number of moles used for each additive using the exact mass of additive from the Data Table and the formula weight of the additive. Record values in the Analysis Table.
4. Calculate the ΔT_f per mole of solute. Record values in the Analysis Table. Note: This is the key column to use when comparing effectiveness of additives; it corrects for the fact that unequal moles of additive (equal mass, but unequal moles) were used.
5. Determine n for each solute. This is the number of particles or ions formed from the dissociation of each formula unit. Write a balanced equation for the dissociation of each solute. Record n for each solute in the Analysis Table.
6. Calculate the concentration of each solution in molality, m, which equals moles of solute per kilogram of solvent as shown in the equation below. Remember to use exact measured values for mass of solvent from the Data Table. Record values for m in the Analysis Table.

Molality = m = moles solute/kg solvent

7. Calculate the theoretical value for the freezing point depression for each solute. This is the freezing point depression that should have been observed using the starting amount of solute and solvent. Use Equation 1 from the background section. Use K_f = 1.86 °C/m, and calculated values for m and n. Record values for ΔT_f (Theoretical) in the Analysis Table.
8. Determine the percent error for the ΔT_f for each solute. Use the following equation. Record values in the Analysis Table.
   {11912_Discussion_Equation_1}
9. Prepare a graph of n versus ΔT_f/mole. Describe the relationship shown by the graph.

Student Worksheet PDF

11912_Student1.pdf