Determination of an Equilibrium Constant

Introduction

A state of chemical equilibrium exists when the rate of the forward reaction is equal to the rate of the reverse reaction. Once equilibrium has established itself, the amounts of products and reactants are constant. Furthermore, if one of the product or reactant concentrations can be measured, it can be used to determine the remaining concentrations and the equilibrium constant.

In this experiment, we will begin by combining aqueous solutions of iron (III) ion (Fe3+) and thiocyanate ion (SCN-). The reaction that occurs produces the thiocyanoiron (III) complex ion responsible for the equilibrium mixture's deep red color.

The final state of the equilibrium (i.e., the final concentrations of products and reactants) depends on the relative amounts of reactants before the reaction occurs. However, regardless of the initial concentrations, the final equilibrium concentrations must satisfy the following relationship:

where the bracketed terms are molar equilibrium concentrations of the different species, and Kc is the temperature dependent equilibrium constant.

In today's experiment, we will study the chemical equilibrium described above by varying the initial concentrations of each reactant. After equilibrium is achieved, the colorimeter will be used to determine the concentration of thiocyanoiron (III) complex ion. This is then used to determine the concentrations of iron (III) ion and thiocyanate ion . These concentrations are then substituted into the expression for Kc to check and determine whether the equilibrium constant is really constant from one case to the next.

The molar absorptivity, e, will be used to convert measured absorbances (via the colorimeter) to FeSCN2+ concentrations. The molar absorptivity of FeSCN2+ will be obtained by using a standard solution that contains an initial 100x excess of Fe3+ ion in comparison to SCN-. According to LeChatelier's principle, this high concentration shifts the reaction far to the right in favor of the product using up approximately 100% of the SCN- ions. Consequently, the [FeSCN2+] in the equilibrium mixture is approximately equal to the original [SCN-].  The [ FeSCN2+] for each of the equilibrium mixtures (vials) may then be determined from their absorbances and the molar absorptivity.

Knowing the equilibrium concentration of the product FeSCN2+ species, it is possible to determine the equilibrium concentrations of reactants. First calculate the number of moles of original reactant before equilibrium was achieved using initial volume and concentration information. Next calculate the number of moles of FeSCN2+ formed. Since products and reactants are in a 1:1 mole ratio and any product formed reduces the amount of reactant remaining in the reaction mixture, the number of moles of remaining reactant is determined by:

Convert the number of moles of reactants by dividing each by the mixture's total volume. Use these concentrations to calculate the equilibrium constant for each of the mixtures.

The colorimeter used in today's experiment will be set to a wavelength of 470 nm. This wavelength is used since it readily absorbed by the FeSCN2+ complex ion. As the concentration of the sample increases, the amount of light that passes through the sample decreases, and the absorbance increases.

Procedure

  1. Obtain a pair of goggles. You will be working in pairs or individually today.
  2. Obtain six vials with screw-on lids and prepare the following solutions using two burets (Use the 0.00200 M Fe(NO3)3 and 0.00200 M KSCN solutions). Be sure to thoroughly mix the contents of each of the six vials before proceeding.

Vial Number

1

2

3 4 5 6

Fe(NO3)3 (mL)

1.0

2.0

 4.0 5.0 6.0 8.0
KSCN (mL)

9.0

8.0

6.0

5.0 4.0 2.0
  1. Preparation of the standard: Using your 10 mL graduated cylinder to measure 9.0 mL of 0.200 M Fe(NO3)3 (read carefully) into a seventh vial. Next, add 1.0 mL of 0.0020 M KSCN from the buret. Mix well.
  2. In the "Experiments" folder of Logger Pro, you will find another folder called "Probes & Sensors". Open the "colorimeter" file from the list of probes and sensors.
  3. Calibrate the colorimeter using the following procedure:
  4. Remember to handle the cuvette by the ribbed sides. If necessary, wipe the smooth sides of the cuvette with a Kimwipe®. Rinse the cuvette twice with ~1 mL portions of the vial #1 contents. Fill the cuvette ~3/4 full. Now, insert the cuvette into the colorimeter, close the lid and record the absorbance reading. Repeat this step for each of the six remaining numbered vials, including the standard.
  5. Place your used solutions in the designated waste container(s). 
  6. Remove your cuvette from the colorimeter since it may severely damage the sensor if left in the sample compartment! Clean all vials and cuvettes before returning them to their appropriate location.

Data Analysis and Questions

  1. Determine the [FeSCN2+] for the standard solution, as described in the Introduction. Then calculate the molar absorptivity, e.
  2. Use the molar absorptivity, e, to determine the [FeSCN2+] in each of the 6 unknown equilibrium mixtures  (writing formulas in Excel will simplify these calculations...don't be afraid to ask for help). 
  3. Calculate the number of moles initially present for Fe3+ and SCN- in each of the 6 vials. Calculate the number of moles of FeSCN2+ present at equilibrium in each of the six trials. Determine the number of moles of each reactant left when equilibrium is reached (see the Introduction) and convert these to concentrations. (Once again, you are encouraged to use formulas in Excel to simplify these calculations).
  4. Use the equilibrium concentrations to calculate equilibrium constants for each of the six trials. 
  5. Calculate the average Kc.
  6. How constant were your Kc values at room temperature? Explain any variation.
  7. See if you can find literature and/or Internet references for the equilibrium constant for this equilibrium. Please cite the reference(s). How does your result compare with the one(s) you cited? Perhaps you will even find a similar experiment, with sample results, posted out on the Web.

Data Table
Use Excel to create a data table that includes six columns—one for each of the six vials. The following information should be included as rows for each trial :

Also include the absorbance for the standard solution and the average values of Kc at the bottom of the table.

Lab Report
You may choose to submit an individual or group report for this lab. Follow the guidelines for Laboratory Reports located at http://webs.anokaramsey.edu/chemistry/Chem1062. For this lab report, you will need to include a title, abstract, procedure, results, discussion, and references. You may use the above questions to guide your discussion, but the discussion should be more than just answering the questions and should flow logically as you discuss the lab and the results.

Follow your instructor’s directions for submitting this lab report. Remember to name the file as specified (Lastname Lab 8 or Lastname1LastName2 Lab8). If you are emailing your report, use the subject line “Chem 1062: Equilibrium”. Also, embed the Excel table in the Word document so that the professor may view the formulas in each cell.  If you submit a paper copy of your report, include one set of sample calculations for one of the trials, either handwritten or typed.  If you worked in pairs and are submitting this assignment on an individual basis, please underline your own name and include your lab partner’s name on the assignment.

Portions of this lab courtesy of Kirk Boraas, Minneapolis Community and Technical College, and Vernier Software, "Chemistry With Computers". Edited by Lance S. Lund, Anoka-Ramsey Community College. Updated July, 2008.