Chemistry Lab Report #5 PDF

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Chemistry Lab Report #5 for General Chemistry 122L. Has all data and typed in correct format. ...

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Solubility and Thermodynamics of Borax Experiment: 5 Chem 121 Lab: Group #: Date: 04/17/2017 Report Writer: Group Leader: Spokesperson: Technical: I certify that I have read and have approved the submission of this group lab report:

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Abstract This lab explored solubility and thermodynamic properties by using a naturally occuring compound called Borax. The purpose of this lab was to determine the Ksp of borax and to

determine the value of Δ H ° , Δ S ° and Δ G ° for the dissolution of borax in water. This was done by preparing a saturated solution of borax and using the solution to prepare samples of borax at different temperatures. The different samples of borax were then each titrated with hydrochloric acid and the volume was recorded. For each titrated sample, the molarity of the tetraborate ion was determined using the number of moles of the tetraborate ion and the specific sample volume. Next the Ksp for borax was determined at each temperature. Using the Ksp and temperature data, Δ H ° , Δ S ° were determined to be 105.9 kJ and 329.3 J respectively. Δ G ° was determined to be 7.58kJ for the dissolution of borax at 25℃. Introduction Borax is a naturally occurring compound and is a very complicated ionic salt as both a solid and when dissolved in water. Borax is used commercially as a water softening agent or for cleaning purposes.[1] The products of dissociation of borax are two sodium ions and “tetraborate,” plus eight waters. The tetraborate ion is a weak base, which means it can be titrated with a strong acid, such as HCl. The equation for this titration is as follows: +¿(aq)+ H 2 O (l )→ 4 H 3 B O3 (aq) ¿ 2−¿(aq)+2 H 3 O ¿ OH ¿ 4 B 4 O5 ¿ To determine Ksp from this reaction the moles of the tetraborate ion is used to calculate the molar concentration of the ion which are used in the Ksp expression. The free energy change in a reaction, Δ G, is related to the equilibrium constant for the reaction by the equation[1]: K Δ G °=−RTln ¿ The free energy expresses the net effect of two thermodynamic quantities: the enthalpy, H, and the entropy, S. Under constant temperature conditions, this is expressed by the equation[1]: Δ G °=ΔH °−TΔS °

The two physical contributions to the free energy are the heat absorbed or released during a reaction, Δ H, and the molecular disorder created during a reaction, S.1 Because the above two equations are expressions of Δ G , the right hand side of each of the two equations above must be equivalent. If the equilibrium constant can be determined at several temperatures the combined equation can be used for determining the values of Δ H and Δ S by linear regression.

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Experimental Method The objective of this experiment is to determine the solubility (Ksp) and thermodynamic values for the dissolution ( ΔG ° , ΔH ° , ΔS ° ¿ of borax. To begin the experiment, prepare a saturated solution of borax, which is done by first adjusting a hot plate to a setting of 6 or 7. Weigh out 25 grams of borax and slowly add to 80 mL of distilled water in a 100-mL beaker. Once the borax is added, place the beaker onto the hot plate and carefully watch the temperature. Frequently stir the solution and record temperature readings, however make sure the thermometer is not resting on the bottom of the beaker. When the temperature of the borax solution reaches between 57 and 58 degrees Celsius, remove the beaker from the hot plate and place on counter. The next steps of this experiment include obtaining 5 sample collections of borax. Adjust the hot plate to a setting of 3 or 4 and place a 100mL beaker with approximately 50 mL of of water on it. Obtain a 10-mL graduated cylinder and place this into the water bath for later use in the experiment. Gather 5 clean Erlenmeyer flasks; two 250-mL and three 125-mL, and label them 1-5. For the first borax sample, stir the saturated borax solution and monitor the temperature until it is approximately 55 degrees Celsius. Once the temperature has been reached, refrain from stirring and allow for the solid to settle out of the solution. When the solid has settled, read and record the temperature to the nearest 0.1 degrees Celsius, then pour between 7 to 9 mL of the borax solution into the 10-mL graduated cylinder sitting in the warm water bath. Record the volume of the borax solution to the nearest 0.1 mL in data section. Transfer the borax solution in the graduated cylinder into the 250-mL Erlenmeyer flask labeled “#1.” Rinse the graduated cylinder several times with distilled water and pour these into the Erlenmeyer flask also. Once rinsings are finished, place the graduated cylinder back into the warm water bath. To collect the second borax sample, allow the saturated solution to cool down to a temperature of approximately 45 degrees Celsius. Be sure to frequently stir the solution. Once the desired temperature has been reached, let the solid settle out of the solution, while the thermometer is still in the beaker. Record the temperature to nearest 0.1 degrees Celsius when the solid has settled out, and again quickly pour between 7 and 9 mL of the borax solution into the warmed 10-mL graduated cylinder. Record the volume of the solution to the nearest 0.1 mL. Repeat the steps of transferring the solution to the second 250-mL Erlenmeyer flask labeled “#2”, rinsing the graduated cylinder several times with distilled water, adding those to the Erlenmeyer flask, and placing the graduated cylinder back into the warm water bath. For the third, fourth, and fifth borax samples, repeat the procedures explained previously with the temperatures of 35, 25, and 15 degrees Celsius and transferring the contents to the labeled “#3”, “#4”, and “#5” 125-mL Erlenmeyer flasks. A cool water and an ice bath may be needed to lower the temperatures to 25 and 15 degrees Celsius. It is important to try and allow for at least 10 minutes for each 10 degrees Celsius decrease in temperature. The final step in determining the solubility and thermodynamic data of borax is to titrate the collected samples. Obtain a 50-mL buret, buret clamp, ring stand, and funnel and assemble together. Collect 200 mL of 0.20 M hydrochloric acid (HCl) in a dry and clean 250-mL beaker. 2

Record the exact molarity of the HCl in data section. Rinse and fill the prepared buret with the HCl acid. Add enough water, but no more than 50 mL, to the labeled 1-5 Erlenmeyer flasks to dissolve the solid borax, if any. Next, add 5 or 6 drops of methyl red indicator to the each of the flasks. Stir the solutions to ensure that all of the solid borax is dissolved. If the borax in samples #1 and #2 are not completely dissolved, the Erlenmeyer flasks may need to be warmed. Titrate each of the samples with the 0.20 M HCl and once the methyl red indicator turns from a yellow to a salmon pink color, end the tiratrion. Be sure to use a full buret on the more concentrated samples, like #1 and #2. Record the volume in milliliters of the HCl used in the data section. Once all the samples have been titrated and volume of HCl has been recorded, clean and return all of the general equipment used back to its original place. Results Sample

Temperature Volume of Volume of Sample (mL) HCl (mL) (℃)

[Borax](M)

Ksp

1

43.1

7.1

43.9

0.62

0.95

2

43.1

7.6

35.8

0.47

0.42

3

32.1

7.5

23.2

0.31

0.11

4

23.5

7.4

13.3

0.18

0.023

5

13.5

8.1

9.8

0.12

0.007

Moles of tetraborate ion Sample 1: 0.2 mol HCl

0.0439L HCl

0.00878mol HCl

1 mol Borax

0.00439 mol tetraborate ion

2 mol H+

1L Sample 2: 0.2 mol HCl

0.0358L HCl

0.00716 mol HCl

1 mol Borax

0.00358mol tetraborate ion

2 mol H+

1L Sample 3: 0.2 mol HCl 1L

0.0232L HCl

0.00464 mol HCl

1 mol Borax

0.00232 mol tetraborate ion

2 mol H+

3

Sample 4: 0.2 mol HCl

0.0133L HCl

0.00266mol HCl

1 mol Borax

0.00133 mol tetraborate ion

2 mol H+

1L Sample 5: 0.2 mol HCl

1L

0.0098 L HCl

0.00196 mol HCl

1 mol Borax

0.00098 mol tetraborate ion

2 mol H+

Molarity of tetraborate ion Sample 1: ● 0.00439 mol tetraborate ion / 0.0071L = 0.62M Sample 2: ● 0.00358 mol tetraborate ion / 0.0076L = 0.47M Sample 3: ● 0.00232 mol tetraborate ion / 0.0075L = 0.31M Sample 4: ● 0.00133 mol tetraborate ion / 0.0074L = 0.18M Sample 5: ● 0.00098 mol tetraborate ion / 0.0081L = 0.12M Acid Molarity Sample 1: ● 2(0.62M) = 1.24M Sample 2: ● 2(0.47M) = 0.94M Sample 3: ● 2(0.31M) = 0.62M Sample 4: ● 2(0.18M) = 0.36M Sample 5: ● 2(0.12M) = 0.24M Ksp for Borax Sample 1: ● [(1.24)2(0.62)] = 0.95 Sample 2: ● [(0.94)2(0.47)] = 0.42 Sample 3: ● [(0.62)2(0.31)] = 0.11

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Sample 4: ● [(0.36)2(0.18)] = 0.023 Sample 5: ● [(0.24)2(0.12)] = 0.007 Sample

lnK

1/T (K)

1

- 0.05129

0.0031

2

-0.86750

0.0032

3

-2.20727

0.0033

4

-3.77226

0.0034

5

-4.96185

0.0035

[2]

[2]

Finding ΔH M = -ΔH / R -12720 = -ΔH / 0.00831446 ΔH = -(-12720 x 0.00831446kJ) ΔH = 105.76kJ Finding ΔS

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b=ΔS/R 39.606 = ΔS/8.31446J ΔS = 39.606 x 8.31446J ΔS = 329.30J or 0.32930kJ Finding ΔG ΔG = ΔH - TΔS ΔG = 105.76 - (298.15)(0.32930) ΔG = 7.58kJ Discussion The solubility of borax increases as the temperature increases because as the temperature was lowered to temperatures of 35, 25, and 15 degrees Celsius, the precipitate was much more visible than at the higher temperatures of 45 and 55 degrees Celsius. The solubility increases as the temperature increases due to the molecules within the solution moving around at faster rates, essentially increasing the amount of space available. The dissolution of borax is an endothermic process, due to the fact that as the temperature increases, the solubility of the borax also increases, causing the reaction to shift to the right to restore equilibrium. This is an example of Le Chatelier’s principle, which states that when a system experiences some sort of disturbance, such as a change in pressure, temperature, or concentration, it will respond by creating a new equilibrium state[1]. Using mathematics, the slope of the results was determined to be negative, M = -12720 and when set equal to the equation of -ΔH/R (R = 8.31446 Joules), ΔH is calculated to be positive, meaning the reaction is endothermic. Entropy is defined as “the amount of molecular randomness in a system[1].” ΔS of the reaction increased because the reaction went from a solid state to a liquid state, making the system became less ordered. When a system becomes less ordered, the entropy or “molecular randomness” is increased. The dissolution of borax is a temperature dependent reaction. The concluded calculations found for the dissolution of borax to a be a non-spontaneous process, meaning that some sort of energy input is required. Though borax is soluble in water, the amount that completely dissolves depends on the temperature. [3] When conducting the experiment, it was found that when the temperature was increased, the amount of borax that dissolved also increased. ΔH and ΔS of the concluded results were positive and when placed into the equation of ΔG = ΔH - TΔS, ΔG is found to also be positive, meaning that the reaction can only be spontaneous at high temperatures. Experimental errors that could cause invalid data could be not waiting for the borax solution to completely settle so the temperature was measured correctly. Another error that could be observed is using too much HCl to titrate which would influence the calculation to find the moles of the tetraborate ion which would then affect the molarity and the ksp value of borax. Conclusion This experiment was performed to determine the solubility (Ksp) and thermodynamic information of borax. Borax is used for multiple different purposes, such as cleaning or as a water-softening agent. Like almost every other element, the solubility increases as the 6

temperature increases, which is important to know when trying to make a neutral solution. The Ksp of borax was calculated to be from 0.007 to 0.95 with the varying samples and concentrations. The thermodynamic data determined that it is a non-spontaneous process as ΔG = 7.58 kJ, ΔH = 105.76 kJ, and ΔS = 0.32930kJ. This experiment could have been expanded by testing the solubility of borax in solvents other than water. While the water is always a good solvent to use, the experiment could offer more information regarding borax solubility when introduced to other solvents, as well as to see how it will react once the solution is placed on the hot plate. This experiment could have also used different amounts of borax dissolved at the same water temperature. Although these two ideas were not offered in the experiment, the original experimental method did get its point through. References [1]. McMurry J. E.; Fay R. C. General Chemistry: Atoms First, 2nd ed.; Pearson, 2014 [2]. Robinson, Bill. "Thermodynamic Information from the Temperature Dependence of Keq." Temperature Dependence of K. Accessed April 17, 2018. https://www.chem.purdue.edu/gchelp/howtosolveit/Thermodynamics/TemperatureDependanceOf K.html [3].Birney, D. Equilibria, ∆G, ∆H and ∆S https://www.depts.ttu.edu/chemistry/Faculty/birney/delta G.pdf (accessed Apr 17, 2017).

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