Calculate the concentration of the dilute thiocyanate

Question

Assignment: Determining the Equilibrium Constant of a Chemical Reaction

Purpose

To determine the equilibrium constant, Kc (also called Kf) for the formation of a complex ion by measuring equilibrium concentrations of the reacting species involved.

Calculations

Part A

1. Calculate the concentration of the dilute thiocyanate solutions ([SCN-]_initial) using the concentration of the stock solution and the dilution equation (M₁V₁ = M₂V₂).

2. Make a graph of the corrected absorbance (y-axis) vs. *[FeSCN²⁺]_eq (x-axis). This is a Beer's law Plot. Add a linear best-fit trend line to the data. Include the equation for the best-fit line and the R² value.

* Remember In part A: [FeSCN²⁺]_eq = [SCN^-]_initial

Part B

3. Calculate the concentration of the dilute thiocyanate solutions ([SCN^-]initial) using the concentration of the stock solution and the dilution equation (M₁V₁ = M₂V₂).

4. Calculate the concentration of the dilute iron solutions ([Fe³⁺]initial) using the concentration of the stock solution and the dilution equation (M₁V₁ = M₂V₂).

(Remember in Part B [Fe³⁺]=0.004M)

5. Using the corrected absorbance data for the six solutions in Part B and the equation for the linear fit to the Beer's Law plot in Part A, determine the concentration of iron thiocyanate complex at equilibrium. ([FeSCN²⁺]_eq) for each solution.

Remember In Part B: [FeSCN²⁺]_eq ≠ [SCN-]_initial.

5. Set up an ICE-box table similar to the one shown on the first page of this procedure. Using the initial concentrations of Fe³⁺ and SCN^- and the equilibrium value for the iron complex, determine equilibrium values for Fe³⁺ and SCN^-.

6. Calculate the value of K_c using the equilibrium values determined above. Repeat this for each of the six samples in part B.

7. Calculate an average value and the standard deviation for K_c.

8. Calculate the 95% confidence interval for K_c. Look up the true value in your textbook. (These are referred to as Kf in your text, page A-12, Complex ion Formation constants in Water). Does your calculated value agree with the true value at the 95% confidence interval?

Questions

1. In Part A the initial Fe³⁺ concentration was made to be very large compared to the initial thiocyanate concentration. Explain why?

2. In part A a blank solution was made containing Fe³⁺, nitric acid and water but no thiocyanate.

a) What is a blank solution? Why was it necessary to use a blank in this analysis?

b) True or False - Once corrected with the blank solution the straight line on the Beer's law plot should pass through the origin. (In other words, zero concentration will give zero absorbance.) Explain your answer.

3. In part B the concentrations of Fe³⁺ and SCN^- for each of the six sample were all very different. As such should the calculated value of Kc be different as well or should it be the same (within error) for all samples. Explain.

4. The value of Kc from the data table in your textbook is 890. At the 95% confidence interval does your result agree or disagree with the published value. In other words does the accepted (true) value from your book fall in the range of your calculated value?

5. Your value is likely much lower. This has to something to do with the HNO₃ in solution. Use your knowledge of Le Chatlier's principle to explain why Kc is lower than the table value.

Attachment:- Experiment.pdf