12_Lab_113

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sFreezing Point Determination for Various Aqueous Solutions Introduction The purpose of this experiment is to determine how the concentration of solute in an aqueous solution impacts the freezing point of the solution. In chemistry, colligative properties exist which depend on the number of dissolved particles in a solution instead of the individual identity of the particles (French, et al. 66). Within this experiment, as the particles in a solution increase, the freezing point decreases. According to French, et al. “this can be represented by the equation: ΔT = i K f m where ΔT is the change in temperature from the original freezing point, i is the van’t Hoff factor, K f is the freezing point depression constant for the solvent, and m is the molality of the solution expressed in moles of solute per kilogram of solvent.” This equation can be used with the molality calculated in the experiment, along with the observed change in freezing point temperature to determine the freezing point depression constant. This value will be an integral part of the experiment to determine how varying concentrations of solute impact freezing points of solutions. In the experiment, freezing point depression constants will be calculated through various experiments that measure the freezing point of sugar solutions. A test tube of water will be prepared as a constant and an additional test tube of NaCl which will explore the impact the van’t Hoff factor has on the freezing point (French, et al. 66). To do this, dry ice must be used to create a temperature bath with temperatures low enough to test the freezing point of various solutions. Obtaining the freezing point in °C, will be determined graphically through extrapolation of flat temperature regions and initial drops (French, et al. 67). Procedure

Procedure and list of materials was inspired from (French, et al. 66-68). List of Materials  Temperature Probe  Beakers of Assorted Sizes  25ml Graduated Cylinder  Weigh boats  Analytical Balance  Tongs  Test tubes  Wash Bottle  250ml Erlenmeyer flask  Salt (NaCl)  Sucrose (table sugar, C 12 H 22 O 11 ) Measurenet 1. Connect temperature probe to workstation and power on 2. Press main menu and function key to select the temperature probe and then again to select “temperature vs time” experiment. 3. Calibrate the probe to room temperature 4. Set up axis limits for display purposes 5. Press “START/STOP” to begin or stop data collection Actual Experiment

1. Zero out the analytical balance with the weigh boat on it 2. Fill each weigh boat with 2g, 4g, 6g, and 8g of sugar 3. Prepare 4 sugar solutions with each of the masses of sugar and 25 ml of water. Ensure that 5ml of water is poured into a test tube at once with a maximum of 2g for each 5ml poured into the test tube while mixing at each stage to ensure solute is fully dissolved 4. Pour 5ml of each solution into a separate test tube 5. Weigh out .342g of NaCl and dissolve in a test tube with 25ml of water. 6. Pour 5ml of the salt solution into a separate test tube 7. Prepare a test tube with 5ml of water as a control 8. Weigh out 50g of rock salt and mix in a 250ml beaker with 200ml of water. Stir until dissolved 9. Use tongs to transport 150ml of dry ice pellets in a 400 ml beaker 10. Pour the saturated NaCl solution over the dry ice pellets and insert the temperature probe 11. Stir the solution until the ice bath is constant (-18°C to -25°C) 12. Insert the temperature probe into the test tube containing solely water and begin data collection 13. After 5 seconds, place tube in the prepared ice bath and continue to collect data until the graph has flatlined and the temperature has stopped changing or reversed, save as “001” 14. Repeat steps 12-13 for the 4 sugar solutions (save as “002-005”) and the salt solution (save as “006”) while drying and rinsing probe between each trial Discussion

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The purpose of this experiment was to determine how freezing points of aqueous solutions are impacted by the addition of different solutes with varying concentrations. The summary of the results is presented in the following paragraph. The freezing point for the solutions are as follows: 0.00°C for water, 2.95°C for the .2343 molal sugar solution, 0.20760°C for the .4695 molal sugar solution, -1.6984°C for the .7045 molal sugar solution, -1.5999°C for the .9377 molal sugar solution, and -3.6611°C for the .2334 molal salt solution. The freezing point depression constant, K f is 1.86°C/m for water, -12.6°C/m for the .2343 molal solution, 2.41°C/m for the .7045 molal solution, 1.706°C/m for the .9377 molal solution, and -0.442°C/m for the .4695 molal solution, with an average value of -2.23°C/m for the sugar solutions. Overall, the results do not support the hypothesis because the data did not fall in line with the colligative property. Although it is possible to understand colligative properties through lectures, the data did not provide any insight into how freezing point depression works. If the percent error for the Kf was lower, it is absolute that the freezing point would be decreased as the solute concentration increased. The set of freezing points are counterintuitive to the colligative property freezing point depression. It is expected that as the concentration of a solute within an aqueous solution increases, the freezing point would decrease. However, there was no continuity within the data. At low concentrations, the freezing point is dependent on the solute concentration and independent of the identity of the solute. This is not displayed by the data as the freezing point for the sugar and salt solution are not like each other. The percent error for the freezing point for the .2343 molal sugar and salt solution is 195.7% and 18.8% respectively. The average value for the sugar solutions is also counterintuitive to the expected values. With an

expected freezing point constant of 1.86°C/m for H 2 O, the graphical value 6.6°C/m has a percent error of 254.8% and the calculated value -2.23°C/m has a percent error of 219.8%. One possible source of error within the experiment is the cooling of the ice bath. Waiting too long to start the series of experiments or not adding enough dry ice could result in this. This would have affected the results by potentially inhibiting the solution from reaching its freezing point or causing the solution to take substantially more time to freeze. To mitigate this, it is important to measure the temperature of the ice bath before beginning the experiments to ensure it is below -18°C as well as adding more ice if necessary to ensure this temperature stays somewhat constant through the trials. Another possible source of error within the experiment is supercooling. This is the lowering of a temperature of a liquid below the freezing point without freezing the liquid. This would result in the temperature dropping below the theoretical freezing point which would skew data significantly for the K f values. This could be mitigated by ensuring the temperature probe is fully in the test tube solution and by continuously stirring. Another source of error is the lack of time available to work on the lab. This limits the amount of time available to work which pressures students to get the data even if it means sacrificing its validity. This would impact the data severely because students would not have enough time to preform the experiments until completion so the temperatures, they are gathering may be higher than the actual freezing points. This could be prevented by spending less time discussing the procedure for the beginning of class and allowing students to preform the experiments from the start. Conclusion

In this experiment, I learned more about how freezing point depression operates mostly through the chem21 activities. In real life, it is interesting to think about this colligative property within bodies of water. In sub zero temperatures for prolonged periods of time, bodies of salt water may not freeze because of freezing point depression. In lakes, the temperature needs to reach 0°C to freeze, but in oceans it is much lower due to the presence of salt within it. Regarding questions, I am still confused the role the van’t Hoff factor played in this experiment. It may be because my experiment was not performed correctly to completion but observing a set of data that is more in line with the expected values would provide more insight. I believe I learned more about experimental processes and how to enhance that more than I did about the topic. This is mostly since I have already taken CHE 107 and understand the content, but the lab experiments are a completely new activity to me. Analyzing the sources of error really allowed me to reflect on my behavior in the experiment and how to recognize and change it moving forward to improve the quality of the experiment. Next time, I will spend more time maximizing the use of my lab materials to maximize the use of both my partner and my time. This would allow us to still have enough time to complete the experiments in full while collecting accurate data. I do believe that writing the procedure in the notebook has helped with the learning process. Works Cited French, April, et al (Department of Chemistry, University of Kentucky). “Freezing Point Depression.” Macmillan Learning Curriculum Solutions: Plymouth, MI, 2020; p. 65-68. Web. Accessed September 21, 2021.

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