Equilibrium - Lab Report PDF

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Equilibrium Name: Shaili Batsri Date: 09/28/2021 Lab Partners: Chris Zalayes

Introduction: In a chemical reaction, chemical equilibrium is the state in which both the reactants and products have the same rate. As a system approaches equilibrium, both the forward and reverse reactions are occurring. When the system is at equilibrium, the forward and reverse reactions are proceeding at the same rate and the amount (concentration) of each reactant and product remains constant. When there is a disturbance to the system at an equilibrium, the LeChâtelier’s Principle is applied. According to this principle, if a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance. We are using LeChâtelier’s Principle qualitatively to predict shifts in equilibrium based on changes in the system conditions. In this lab, we are going to determine the equilibrium constant Kc of the reaction between Fe3+ and SCN-, by calculating the ions concentrations.

Procedure: 1. 2. 3. 4.

Set up a spectrophotometer at 447 nm wavelength. Mix Fe(NO3)3 solution with distilled water and transfer the new solution to a beaker. Mix Fe(NO3)3, KSCN, HNO3 solutions, and distilled water to a cuvet. Measure the absorbance using the spectrophotometer at 447 nm wavelength.

*Repeat the steps but vary the volume of Fe(NO3)3, KSCN, and distilled water. 5. Warm the cuvet from step 3 using boiling water and measure the absorbance. 6. Prepare a new cuvet containing the same volumes used in step 3 and add AgNO3 until change in the solution is observed. 7. Dispose all the chemical solutions in the proper waste container.

Date Sheet: -

Data table:

Experiment Number

Volume Fe(NO3)3

1 2 3 4 5

1.00 mL 2.00 mL 0.50 mL 1.00 mL 1.00 mL

6 Molarity

1.00 mL 0.01 M

Experiment Number 1 2 3 4 5 6

-

Volume Volume Volume Absorbance Kc value KSCN HNO3 Water ROOM TEMPERATURE 1.00 mL 2.00 mL 2.00 mL 0.295 53 1.00 mL 2.00 mL 1.00 mL 0.590 59 1.00 mL 2.00 mL 2.50 mL 0.253 91 2.00 mL 2.00 mL 1.00 mL 0.554 50 0.50 mL 2.00 mL 2.50 mL 0.149 53 o 100 C 1.00 mL 2.00 mL 2.00 mL 0.196 34 0.002 M 1.00 M -

Equilibrium Concentration Initial Concentration Fe(SCN)2+ of Fe3+ ROOM TEMPERATURE -5 2.6818x10 M 1.67x10-3 M -5 5.3636x10 M 3.33x10-3 M 2.3x10-5 M 8.33x10-4 M -5 5.0364x10 M 1.67x10-3 M 1.3545x10-5 M 1.67x10-3 M o 100 C -5 1.7818x10 M 1.67x10-3 M

Initial Concentration of SCN3.33x10-4 M 3.33x10-4 M 3.33x10-4 M 6.667x10-4 M 1.667x10-4 M 3.33x10-4 M

Observations after adding AgNO3 to solution: The solution became less brown and had a precipitate in the bottom of the test tube.

Post-Lab Questions: 1. In all cases you used a total of 6.00 mL of solution. Calculate the initial concentr concentration ation of Fe3+ for all five of the room temperat temperature ure experiments, after dilution tto o 6.oo mL.

2. Calculate the initial concentration of SCN- for aallll five of the room temperature experiments, after dilution to 6.00 mL.

3. Determine the equilibrium conc concentration entration of Fe(SCN)2+ for all five room temperature experiments. Assume that the pathlength of the cuvet or ttest est tube is 1.00 cm.

4. Using ICE tables, determine the K c value for all five room temperature experiments. - Experiment Number 1:

I C E

-

Fe3+ 3.33x10-3 -x 3.27637x10-3

SCN3.33x10-4 -x 2.7937x10-4

Fe(SCN)2+ 0 +x 5.3636x10-5

Fe3+ 8.33x10-4 -x 8.10x10-4

SCN3.33x10-4 -x 3.10x10-4

Fe(SCN)2+ 0 +x 2.3x10-5

Experiment Number 4:

I C E

-

Fe(SCN)2+ 0 +x 2.6818x10-5

Experiment Number 3:

I C E

-

SCN3.33x10-4 -x 3.06182x10-4

Experiment Number 2:

I C E

-

Fe3+ 1.67x10-3 -x 1.643182x10-3

Fe3+ 1.67x10-3 -x

SCN6.667x10-4 -x

Fe(SCN)2+ 0 +x

1.619636x10-3

6.16306x10-4

5.0364x10-5

Experiment Number 5:

I C

Fe3+ 1.67x10-3 -x

SCN1.667x10-4 -x

Fe(SCN)2+ 0 +x

E

1.656455x10-3

1.53125x10-4

1.3545x10-5

5. Determine the average value of K c at room temperature.

6. Did the absorbance of the solution go up or down when it has heated? Is t his reaction endothermic or exothermic? Explai Explain n how you know. The absorbance of the solution went down when the solution was heated. According to Beer’s Law, when the absorbance value is lower, the concentration is lower. Since the equilibrium constant is the concentration of the products divided by the concentration of the reactant, getting a lower product concentration will result in a lower kc value. At a lower kc value, the reaction shifts to the left toward the reactants. Therefore, the reaction is exothermic exothermic. Another explanation is when the equilibrium mixture is heated, by Le Châtelier’s Principle, the equilibrium position will shift to the left to consume some of the additional heat. Therefore, the reaction is exothermic. The K c of the reaction in 100oC is lower than the kc values at room temperature, showing the reaction is exothermic:

I C

Fe3+ 1.67x10-3 -x

SCN3.33x10-4 -x

Fe(SCN)2+ 0 +x

E

1.652182x10-3

3.15182x10-4

1.7818x10-5

Increase in temperature, decrease Kc value, the reaction is exothermic.

7. How does increasing the concentration of FFe e3+ affect the value of kc? Explain. Changes in concentrations does not affect the kc value. If the concentration of a reactant, such as Fe3+, increases, then one can expect that the value of Kc to decrease (because the denominator would increase). However, when the concentration of a reactant, Fe3+, increases, the forward reaction will be favored. So, the concentration of the product Fe(SCN)2+ will increase, the concentration of Fe3+ will decrease, and K c will be kept constant. Therefore, increasing the concentration of Fe3+ will not affect the value of Kc. Another explanation, when we add a reactant to the reaction, the reaction quotient Q gets a lower value than kc. Since the reaction want to restore the equilibrium, Q=kc , the concentrations of Fe3+ and SCN- decreases to form more product, Fe(SCN)2+, therefore the reaction shifts to the right until Q returns to be equal to kc which did not change when the concentration of Fe3+ increased.

8. Does changing the temperature alter the value of kc? Explain. Yes, temperature affects the value of Kc because changes in temperature shift the reaction toward a certain side, shifts from endothermic / exothermic. - Increase in temperature, decrease Kc value, the reaction is exothermic. A ⇌ C + HEAT exo exothermic (system release heat) -

Decrease in temperature, increase Kc value, the reaction is endothermic. A + HEAT ⇌ C endothermic (system absorb heat)

9. Silver ion reacts with thiocyanate in the following reaction: Ag+ + SCN- ⇌ AgSC AgSCN N Châtelier’s r’s principle to explain your AgSCN, unlike Fe(SCN)2+, is colorless. Use Le Châtelie observations when you added A AgNO gNO3 to your solution. After adding AgNO3 the solution turn out to be less brown and had a precipitate in the bottom of the test tube. When we added the AgNO3 we created a precipitation reaction when the Ag+ reacted with SCN- to produce insoluble AgSCN, which settles to the bottom of the test tube, and since AgSCN is colorless we got a less brown solution. The Le Châtelier’s principle states that when a chemical system at equilibrium is disturbed, it returns to equilibrium by undo the disturbance. Therefore, by adding AgNO3 to the solution we disturb the equilibrium system and caused it to shift to the left.

By adding AgNO3 we disturbed the system by removing SCN- ions from the reaction. Therefore, the reaction will shift to the left to create more SCN- ions to restore to the equilibrium state.

Conclusion: In this lab, we examined the equilibrium constant, kc. the objectives of the experiment were to determine the equilibrium constant K c of the reaction between Fe 3+ and SCN-, by calculating the ions concentrations. The results showed that the average value of the equilibrium constant Kc was 61.2. In addition, the results also showed that the reaction between Fe3+ and SCN- is an exothermic reaction. In this lab I learned how to find the value of Kc for a reaction and how to calculate concentrations using Beer’s Law . I learned about the spectrophotometer and its use to find equilibrium concentrations. In addition, I learned that the value of the equilibrium constant, kc, is not dependent on the concentrations of the reactants and the products. However, it is affected with change in temperature....