experiment 15 chem 113 PDF

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experiment 15 chem 113...

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Using Calorimetry to Determine Equilibrium Constant

Introduction: The purpose of this experiment is to determine the equilibrium constant, Kc, by preparing a standard solution and finding its concentrations using a calorimeter. The Kc value found will be for the following reaction between iron(III) and thiocyanate to produce thiocyaniron(III): SCN ¿ ¿ 2+ ¿(aq) −¿(aq)↔ Fe ¿ 3+¿ ( aq )+ SC N ¿ ¿ Fe

The three components of this reaction will reach equilibrium when combined. If the concentrations for the three ions are found for the reaction, the Kc value can then be calculated using the equation below: K c=

[ products] [ reactants]

In order to compare the results for Kc, four different concentrated solutions for this equilibrium reaction will need to be prepared; the different concentrated solutions will allow for the obtained value for Kc to be checked for its accuracy. Making an ICE table will be used for the calculations for each of the reactions in this experiment. The concentration of [Fe(SCN)2+] will be found by using a colorimeter. This ion produces a solution with a red color, therefore the blue LED (470 nm) setting is used on the colorimeter because “red solutions absorb blue light very well” (French, et al. 88). The amount of blue light absorbed, or the absorbance, is measured by the colorimeter. After each absorbance for the

equilibrium system, Aeq, is found, it can then be compared with the known “absorbance of a standard solution, Astd” that contains a known Fe(SCN)2+ concentration (French, et al. 88). The standard solution will be prepared with a large concentration of Fe3+, approximately 100 times larger than the amount in the other four solutions, and a very small concentration of SCN-. It is important to note that the Fe3+ concentration for the standard solution is large because it “forces the reaction far to the right” and uses up almost all SCN- ions; this concept is understood through Le Châtelier’s principle (French, et al. 89). The standard solution will demonstrate that the absorbances have a direct relationship because one mole of Fe(SCN)2+ is produced for every one mole of SCN- used up. This direct relationship between the absorbances is known as Beer’s law. This relationship can be illustrated through the equation below: 2+ ¿¿std

SCN ¿

=

Fe ¿ ¿ 2 +¿¿ SCN ¿ ¿ Fe ¿ ¿ ¿

A eq A std eq

(French, et al. 89). If the [Fe(SCN)2+]eq concentration is known, the other two ion concentrations can be found since it is observed that there is a 1:1 ratio of coefficients; the ratio demonstrates that “for each mole of Fe(SCN)2+ produced, one less mole of Fe3+ ions” are found in the solution (French, et al. 89). The equation below describes the equal relationship between consumption of Fe3+ and formation Fe(SCN)2+:

2+¿ ¿ eq ¿ SCN ¿ Fe ¿ 3+¿ ¿ i−¿ ¿ Fe 3+¿ ¿eq =¿ ¿ Fe ¿ (French, et al. 89). Consequently, the equation below displays the relationship between SCN- and Fe(SCN)2+; showing that “one mole of SCN- is used up for each mole Fe(SCN)2+ produced”: 2+ ¿ ¿eq SCN ¿¿ Fe ¿ −¿ ¿i −¿ SC N ¿ −¿ ¿eq =¿ ¿ SC N ¿ (French, et al. 89). After using the information and the equations above to find the concentrations of all three ions, the value of Kc can then be determined.

Methods: Materials: (French, et al. 89).         

MeasureNet Colorimeter Two cuvettes with caps Five 6” test tubes Thermometer Three small beakers 0.00200 M KSCN 0.00200 M Fe(NO3)3 0.200 M Fe(NO3)3

   

Kimwipes Funnel Pipets Pipet bulb

Procedure: 1. Gather all materials and label each beaker for each solution that will be used. 2. Collect around 20mL of 0.002M Fe(NO3)3, 0.002M KSCN, and 0.2M Fe(NO3)3 in three separate beakers. 3. Fill a beaker with around 20mL of deionized water. 4. Create sample 1 by adding 5mL 0.002M Fe(NO3)3, 2mL 0.002M KSCN, and 3mL DI water to a test tube. Record exact measurements. 5. Create sample 2 by adding 5mL 0.002M Fe(NO3)3, 3mL 0.002M KSCN, and 2mL DI water to a test tube. Record exact measurements. 6. Create sample 3 by adding 5mL 0.002M Fe(NO3)3, 4mL 0.002M KSCN, and 1mL DI water to a test tube. Record exact measurements. 7. Create sample 4 by adding 5mL 0.002M Fe(NO3)3 and 5mL 0.002M KSCN to a test tube. Record exact measurements. 8. Create the standard solution by adding 2mL 0.002M KSCN and 18mL 0.200M Fe(NO3)3 to a test tube. Record exact measurements. 9. Record observation of each solution after mixing them. 10. Stir each solution with stirring rod. Make sure to rinse stirring rod with deionized water and dry it before stirring each solution. 11. Set up MeasureNet station with colorimeter. Press “Main Menu”, “Colorimetry/Fluor/Turb”, “Colorimetry”, “Blue LED”, “Kinetics Experiment”. 12. Wash each cuvette with deionized water.

13. Fill each cuvette with each sample. Wipe fingerprints off with Kim wipe. 14. Make the reference cuvette by filling it with deionized water. Place in “R” slot of colorimeter. Make a blank cuvette by filling it with deionized water. Place in “S” slot. Close lid to colorimeter. Calibrate Colorimeter. 15. Remove blank cuvette and place sample 1 in “S” slot and close lid. Measure and record absorbance. 16. Repeat step 15 with samples 2, 3, 4, and standard solution. 17. Once data is collected, clean all materials used and put everything away.

Discussion: The purpose of this experiment is to find the equilibrium constant, Kc, of an equilibrium reaction by using a colorimeter to measure absorbance and comparing the results with a standard solution of known concentrations. The Kc value determined is for the equilibrium system found below: SCN ¿ ¿ 2+ ¿(aq) −¿(aq)↔ Fe ¿ 3+¿ ( aq )+ SC N ¿ ¿ Fe The average Kc value found is 143.515 with a standard deviation of 5.856. The results of this lab support the purpose because the equilibrium constant was determined for the given equilibrium system. In addition, the equilibrium constant calculated is shown to be accurate because of the small standard deviation for the average Kc of four trials.

Several sources of error could have still affected the results even though the standard deviation was very low. The first source of error could have been from not rinsing the stirring rod well enough with DI water when stirring each of the solutions with the same rod. This could have cross-contaminated solutions therefore changing the concentrations that were found and created for each solution. This mistake could have been prevented by either cleaning the stirring rod very thoroughly after each use or by using a different stirring rod for each solution. A second source of error could have occurred from cross-contaminating solutions when measuring all components with the same pipet. This would affect the volumes and the concentrations of each trial solution if there is extra liquid or chemical residue not rinsed off from the pipet between chemicals. This could be prevented by using a different pipet for each solution or thoroughly rinsing and drying the pipet between each use. A third source of error could have been from not wiping all fingerprint marks off of the cuvette with the Kimwipe before placing it in the colorimeter. This would affect the absorbance determined because the fingerprints would cause the light to scatter and result in less light detected passing through the sample, which then results in a greater absorbance value determined from the colorimeter. This could be prevented by thoroughly wiping the colorimeter with the Kimwipe and making sure to handle the cuvette in a way that leaves as little fingerprints as possible.

Conclusion: From this experiment I learned how to use absorbance to find the equilibrium constant and how to use a standard solution to compare different trials by means of diluted calculations to observe ion concentrations for the equilibrium equation.

Chemical equilibrium is a well-known and important concept in the real world. An important example of chemical equilibrium are the enzymes that are in our body; these enzymes work to speed up the reaction rate for chemical reactions in our body so that the reaction reaches equilibrium faster. Standard solutions are also frequently used in labs or places like pharmacies as a reference solution when compounding drugs like IVs. This experiment could be improved by allotting each lab group different pipets and stirrings rods for every solution in order to prevent cross-contamination and to save time from having to wash the materials off every time they are used.

Work Cited: French, April N., Allison Soult, Stephen Testa, Meral Savas, Francois Botha, Carolyn Brock, Charles Griffith, Darla Hood, Robert Kiser, Penny O’Conner, William Plucknett, Donald Sands, Diane Vance, William Wagner. “Chemical Equilibrium: Finding a Constant, Kc.” Chem 113 General Chemistry II Lab Manual. Plymouth, M: Hayden-Mcneil, 2018. 87-90. Web. 30 March 2020. https://www3.chem21labs.com/labfiles/36194_43_Exp%2015_FrenchA %202187-1%20W20..pdf?rf=9783....