Experimental Procedure

A. Determining the Freezing Point of Pure Solvent.

1. Set up the Logger Pro as indicated in the introduction.
2. Fill a 600 mL beaker ~ 1⁄2 with ice. Add water just below the top of the ice level.
3. Add ~ 30 g NaCl to the ice water. This

value can be approximate and does not have any influcence on experimental calculations.

4. Tarea50mLgraduatedcylinder.
5. Measureout~12mLofdeionizedwaterina graduated cylinder. Record the exactmass measured in your data section.
6. Transfer the deionized water to the testtube. Cover with either foil or a rubberstopper.
7. Insert the temperature probe so that probeis ~ 1⁄2 way into the solvent (it should not touch the bottom of the test tube) and in the middle of the test tube.
8. Place the stirrer into the solvent and ensure it will be able to agitate the solution without bumping the temperature probe.
9. Secure the test tube apparatus to a ring stand with a clamp.
10. Once the test tube is secured, lower it so that the level of the solvent in the test tube is

BELOW the level of ice water in the beaker. (Set up your apparatus as indicated in the Figure).
11. Begin recording temperature in 10 second intervals with constant agitation. (You may need to move BOTH the temperature probe AND the stirrer to prevent the stirrer from hitting the probe). You do not need to include the temperature data in your data section until the temperature reaches ~15 ◦C. At that point, begin recording the data in your data section.

12.The temperature will decrease until it reaches the freezing point. As long as both solid and liquid levels are present, the temperature will remain constant. Once the solution has completely frozen, the temperature will cool again. You should be able to see the freezing point as 3-4 temperature measurements that remain constant.

13.When you feel you have sufficient data. Stop the temperature collection.
14.Use the ring stand clamp to raise the test tube out of the ice bath so that the solvent can thaw for Part B.
15.While waiting for your solvent to thaw, graph your data (Temperature vs. Time) in the space provided.

B. Measuring the Freezing Point of the Solution (Covalent Compounds)

1. Measure out 0.05 -0.1 g of a nonionizing unknown solute on the analytical balance. (Youneed at least 2 significant figures).
2. Transfer the unknown to the test tube and stir to dissolve.
3. Leaving the ice, pour the water from the 600 mL beaker down the drain. Add additionalice, if necessary, to ensure the beaker is still ~ 1⁄2 full of ice.
4. Add an addition ~20 grams of NaCl to the ice bath. This value can be approximate andwill not influence your calculations.
5. Lower the solution into the ice bath.
6. Begin recording temperature in 10 second intervals with constant agitation. (You mayneed to move BOTH the temperature probe AND the stirrer to prevent the stirrer from hitting the probe). You do not need to include the temperature data in your data section until the temperature reaches ~15 ◦C. At that point, begin recording the data in your data section.
7. When you feel you have sufficient data. Stop the temperature collection.
8. Usetheringstandclamptoraisethetesttubeoutoftheicebathsothatthesolventcanthaw for Part B.
9. While waiting for your solvent to thaw, graph your data (Temperature vs. Time) in thespace provided.

10.Repeat this process for a second trial by adding an additional 0.05 to 0.1 g of the same unknown.
C. Measuring the Freezing Point of the Solution (Ionic Compounds)

1. Clean and dry your test tube. Measure out another ~12 mL of deionized water in a graduated cylinder. Record the exact volume measured in your data section.
2. Set up your apparatus as used above.
3. Measure out 0.05 -0.1 g of an ionizing solute on the analytical balance. (You need atleast 2 significant figures).
4. Transfer the solute to the test tube and stir to dissolve.
5. Leaving the ice, pour the water from the 600 mL beaker down the drain. Add additionalice, if necessary, to ensure the beaker is still ~ 1⁄2 full of ice.
6. Add an addition ~20 grams of NaCl to the ice bath. This value can be approximate andwill not influence your calculations.
7. Lower the solution into the ice bath.
8. Begin recording temperature in 10 second intervals with constant agitation. (You mayneed to move BOTH the temperature probe AND the stirrer to prevent the stirrer from hitting the probe). You do not need to include the temperature data in your data section until the temperature reaches ~15 ◦C. At that point, begin recording the data in your data section.
9. When you feel you have sufficient data. Stop the temperature collection.

10.Use the ring stand clamp to raise the test tube out of the ice bath so that the solvent can thaw for Part B.
11.While waiting for your solvent to thaw, graph your data (Temperature vs. Time) in the space provided.
12.Repeat this process for a second trial by adding an additional 0.05 to 0.1 g of the same solute.
13.All solutions must be added to the CHM 112 Waste container in the back hood. 14.Clean the glassware and return all equipment to its proper location.

Pre-lab Assignment/Questions

N o t e – this pre-lab must be finished before you come to lab. (Please see syllabus for how to submit this assignment.)

1. You add 0.991 g of a nonionizing solute to 10.972 grams of water. The freezing point of the solution is 3.48◦ lower. The Kf of water is 1.86 ◦C/m.

a. Determinethemolalityofthesolution.

b. Calculate the molar mass of the solute.

c. If the solute had been one that ionizes into 2 ions, what would the molar mass have been?

2. Which of the following solutes should have the greatest impact on freezing point depression? Explain why using an equation.

MgCl2
KNO3
Glucose (C6H12O6)

Experimental Data and Results

*Include all work, units and write answers in scientific notation (if applicable) using the correct number of significant figures for full credit. Obtain a signature prior to leaving lab.

A. Determining the Freezing Point of Pure Solvent.

B. Determining the Freezing Point of Solution (Covalent Compounds).

C. Measuring the Freezing Point of the Solution (Ionic Compounds) van’t Hoff Factor ________________

Experimental Data and Graphs

A. Determining the Freezing Point of Pure Solvent.
Time Temperature Time Temperature Time Temperature

Jessica Morales. Modified from http://www.bc.edu/schools/cas/chemistry/undergrad/genexp.html Boston College Chemistry Experiments by Lynne O’Connell is licensed under a Creative Commons Attribution-NonCommercial 4.0 International License.

*Kf and Kb data from Merck Index

Experimental Data and Graphs

B. Measuring the Freezing Point of the Solution (Covalent Compounds)
Time Temperature Time Temperature Time Temperature

*Kf and Kb data from Merck Index

Experimental Data and Graphs

C. Measuring the Freezing Point of the Solution (Ionic Compounds)
Time Temperature Time Temperature Time Temperature

Check this box if all materials in your tray have been cleaned, the trays and other materials have been returned to their proper position, and all items in your drawer are accounted for.

Instructor Signature_______________________________________________________________

Post Lab Questions

*See this syllabus for instructions on how to turn in this section of the lab handout.

1. A student failed to clean their test tube and had an impurity that is insoluble in water. What effect would this have on the molar mass determination (too high, too low or no change)? Explain your answer.
2. A spike in temperature is often seen when a solid is first dissolved into water. Provide an explanation for this initial increase. (See the chapter on solutions in your textbook).
3. Compare the molar mass you calculated for the unknown in Part B to the possible compounds used in Table 2. Which compound do you think your unknown was? Determine the percenterror of your calculation. 𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝑒𝑟𝑟𝑜𝑟 = 𝐴𝑐𝑡𝑢𝑎𝑙−𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑥 100 𝐴𝑐𝑡𝑢𝑎𝑙
4. List 2 sources of error that you could have had in Part B of this experiment. For each indicate whether the error would cause the molar mass of the solute to be calculated as too high or too low.