Equilibrium - Thiocyanate

Measuring an Equilibrium Constant

Learning Goals

     Through the use of spectrometry and an ICE table, the concentrations of all species of a reaction at equilibrium will be     calculated, and the equilibrium constant for the reaction can be determined.

Overview

You will use spectrometry to measure the concentration of iron (III) thiocyanate ions, Fe(SCN)2+, in various aqueous mixtures of iron (III) nitrate and potassium thiocyanate. You will use the results of these measurements to determine the equilibrium constant for the formation of Fe(SCN)2+, described by this equation:

                                      Fe3+ (aq) + SCN- (aq) <==> Fe(SCN)2+ (aq) 

Lecture Connections 

Experiment Background

a.) Determining Solution Concentrations by Spectrometry

Measuring the concentration of a solute by its absorbance of light is an example of a spectrophotometic or colorimetric method. The amount of light absorbed (called absorbance, A) by solutes depends on three factors:

1)  Nature of the solute

Some solutes, such as sugar, absorb no light, while others, like the pigments in grape juice, absorb light strongly. Solutes that absorb visible light produce colored solutions.

2)  Distance light travels through the solute

Obviously, light travelling one meter through a solution encounters more solute particles than light travelling only a centimeter through the same solution. Grape juice in a slender glass looks lighter in color (absorbs less light) than does grape juice in a glass of larger diameter. The absorbance of a solution is directly proportional to the path length (usually called b), the distance light travels through the solution.

3)  Concentration of the solute

Light traveling through a concentrated solutions encounters more particles of solute than does light traveling the same distance through a dilute solution. Grape juice concentrate is more darkly colored (absorbs more light) than is dilute grape juice. The amount of light absorbed by a solution is often, but not always, directly proportional to the concentration of the absorbing solute. If by experiment, we find that absorbance is proportional to concentration, we say that such solutes obey Beer’s law:

A = ebC

This simple law relates absorbance A, the light absorbed on passing through b cm of a solution, to the concentration C of the absorbing solute. The constant of proportionality e is called the molar absorptivity (units: M-1 • cm-1).

As its name suggests, you can also think of e as the amount of light absorbed when light passes through 1.00  cm of a 1.00 molar solution of the solute. The molar absorptivity is a characteristic of a particular solute (factor #1 above) at a particular wavelength of light. If a solute absorbs visible light (wavelength 400-700 nm), a large value of e means that that even a dilute solution of the solute exhibits noticeable color. In older chemical literature, e is called the extinction coefficient, and A is called the optical density.

When light of intensity I0 passes through an absorbing medium, such as a solution containing an absorbing solute, the intensity of the light is reduced to intensity I.

The fraction or ratio I/I0 is called the transmittance, T, of the solution. Many instruments for measuring absorption (called spectrophotometers) allow determination of this ratio, as well as conversion of this ratio to absorbance A by the relationship A = - log T. Many instruments display both absorbance and transmittance, often as a percentage (T x 100%).

If a solute obeys Beer’s law, then a graph of A vs C for a series of solutions is a straight line of slope equal to be. This graph is called a Beer’s-law plot, and it has two functions: a) to determine whether the solute obeys Beer’s law, and if so, b) to determine e.

b.) Equilibrium Formation of Iron (III) Cyanate Ion (FeSCN2+)

 Iron (III) ions react with thiocyanate ions (SCN-) to form iron (III) thiocyanate, FeSCN2+, which is a stable deep red complex ion:

Fe3+ (aq) + SCN- (aq) <==> FeSCN2+ (aq, deep red)

The equilibrium constant expression for this reaction is

   Keq = [FeSCN2+]/([Fe3+]•[SCN-]

Solutions of iron (III) ions are weakly colored and thiocyanate ion is colorless, so the primary absorber in a mixture of these components is FeSCN2+. This ion obeys Beer’s law at 460 nm over a fairly wide range of concentrations, allowing measurement of its concentration in equilibrium mixtures. From its measured equilibrium concentration and the initial concentrations of Fe3+ and SCN-, it is possible to determine the concentrations of all three components, and from them, the equilibrium constant for the reaction.

Iron(III) ion introduces a complication because of its reaction with water to form iron hydroxide, which is insoluble in water:

Fe3+ (aq) + 3 H2O (l) <==> Fe(OH)3 (s) + 3 H+ (aq)

To avoid precipitation of iron (III) hydroxide, you will include excess nitric acid (HNO3) in all solutions, to shift this equilibrium far to the left. Because neither hydrogen ions nor nitrate ions are components of the iron (III) thiocyanate equilibrium, nitric acid does not affect the equilibrium position of the reaction that produces FeSCN2+.

In Part 1 of this experiment, you will prepare a series of calibration solutions having known concentrations of the iron (III) thiocyanate ion. To do so, you will use a small, known amount of thiocyanate, and a large excess of Fe3+ to drive the reaction to the right, incorporating all of the thiocyanate into FeSCN2+. In each of these solutions, therefore, the equilibrium concentration of FeSCN2+ equals the initial concentration of SCN-. You will determine the absorbance of each solution, prepare a Beer’s law plot, and obtain the equation of the line relating A and C.

In Part 2, you will prepare five mixtures with known initial concentrations of iron (III) and thiocyanate ion. You will determine the absorbance of each mixture after it reaches equilibrium, and then use the equation of your Beer’s law plot (Part 1) to determine the concentration of FeSCN2+ in each equilibrium mixture. Finally, from the measured [FeSCN2+], and the known initial [Fe3+] and [SCN-], you will determine the equilibrium constant Keq for this reaction.

In the five different equilibrium solutions, the equilibrium concentrations, and hence the equilibrium positions, are different. But your measured values of the equilibrium constant should be the same, within expected experimental variation.

The following problems require skills and calculations similar to those called for in the report form on this experiment. Learn how to work these problems, showing all calculations with units. For problems involving calculation, answers are provided.

 Practice Problems:

1)  The absorbance of a solution of solute X in a 1.5-cm sample tube is 0.755. Use the Beer’s law plot above to determine the concentration of X in this solution. (Answer: 0.132 M)

2)  A calibration solution contains 1.50 mL of 0.00200 M NaSCN, 5.00 mL of FeCl3, and enough

0.1 M HNO3 to make a total volume of 25 mL. Assuming that [Fe3+] is sufficient to drive the reaction to completion, what is the concentration of FeSCN2+ in this solution? (Answer: 1.20 x 10-4 M)

3)  A student prepares an equilibrium solution by mixing 5.00 mL 0.002 M Fe3+, 2.00 mL 0.00200 M SCN-, and 3.00 mL 0.1 M HNO3. What are the initial concentrations of Fe3+ and SCN- in this solution? (Answers: [Fe3+] = 1.00 x 10-3; [SCN-] = 4.00 x 10-4)

4)  By measuring the absorbance of the solution in Question 3, the student finds that [FeSCN2+] =

8.70 x 10-5. Set up an ICE table, using the information in Question 3 and this question. Use the ICE table to calculate the equilibrium concentrations of Fe3+ and SCN- in this solution. (Answers: [Fe3+] = 9.13 x 10-4; [SCN-] = 3.13 x10-4)

5)  Use the equilibrium concentrations from question 4 to calculate Keq for the formation of FeSCN2+. (Answer: 3.04 x102)

Experimental Procedures

Work in pairs to collect data. Carry out all other work and reporting individually.

Part 1. Obtaining Calibration Data for FeSCN2+ Solutions

 a)   Calibration Solutions

CAUTION: The Fe(NO3)3 reagents for Parts 1 and 2 are different: 0.200 M in Part 1, and 0.00200 M in Part 2. Read the instructions and reagent labels carefully.

In your lab notebook, set up a data table in the format shown in Part 1 of the Report Form (Table 1). Enter data and results in your lab notebook as you proceed.

Prepare each calibration solution by pipetting the amounts of 0.200 M Fe(NO3)3 and 2.00 x 10-3 M KSCN shown in Table 1 into a clean 25-mL volumetric flask.   Use a repipetter for the 0.200 M Fe(NO3)3 and a graduated pipet for the 2.00 x 10-3 M KSCN. Increase the solution volume to precisely 25.00 mL by adding 0.1 M HNO3 until the flask is filled to the mark. Add the last mL with a dropper in order to avoid filling past the mark. Stopper the flask securely and mix completely by inverting and rotating the flask several times. Transfer each solution to a clean, dry flask labeled with the calibration solution number. Rinse the volumetric flask thoroughly with distilled water before preparing the next solution. Make all five solutions before proceeding to Part 1b of the procedure.

b)    Calibration Data

You will use a Vernier Spectrophotometer for the measurements. Set the wavelength of the spectrophotometer to 460 nm. You can view a YouTube video of the operation of the instrument at URL: https://www.youtube.com/watch?v=diqdn-xalqM.

Obtain a sample cuvette. Make sure it is clean inside and out. Avoid touching the bottom third of the tube, to keep its light path clear. Rinse the cuvette in, with small amounts (~1 mL) of Solution C1 (the blank). Then fill the cuvette about two-thirds full. Wipe off the outside of the cuvette, and insert it into the sample compartment with its orientation mark aligned with the mark in the spectrophotometer. Calibrate the spectrophotometer. Rinse in the cuvette with solution C2, wipe it off, insert it into the instrument, and read the absorbance. Repeat for the remaining calibration solutions, recording absorbances to three decimal places in a data table in your lab notebook.

 LANGUAGE NOTE: Because you have set the instrument to A = 0.000 with solution C1, we say that you have “determined the absorbance of solutions C2 through C5 against C1 as a blank.”

Finally, check to be sure that the instrument is still properly blanked and zeroed, as follows: rinse in and fill the cuvette again with Solution C1. Read its absorbance. It should still be very close to 0.000. It should still be close to 0.00. If either reading has changed by more than 5 in the last decimal place, you should repeat Part 1b.

c)   Beer’s-Law Plot

After lab, use Excel to plot your data as absorbance (vertical axis) vs. concentration (horizontal axis). Fit a straight line to your data, and obtain the equation of the line and the regression coefficient (R2). Label the axes of the graph and give it an informative title from which the reader can readily understand the relationship represented on the graph. Make two copies of the graph. Hand in one with your notebook pages and report form, and tape the other one on the corresponding white pages of your lab notebook. Working in your lab notebook, determine e for the FeSCN2+ ion.

Part 2. Obtaining Absorbance Data for Equilibrium Mixtures

a)   Equilibrium Absorbance Data

In your lab notebook set up a data and results table in the format shown in Part 2 of the Report Form

(Table 2). Record data in the table as you proceed.

Using a repipetter for the 0.00200 M Fe(NO3)3, and graduated pipets for the KSCN and HNO3, prepare the following solutions in clean, dry, labeled test tubes that will easily hold 10 mL of solution and allow mixing.

Mix each solution thoroughly for 1 to 2 minutes. Make up all five solutions before measuring absorbances. As in Part 1, zero the spectrophotometer, and then determine and record the absorbances of solutions E2 through E4 against E1 as a blank.

b)    Calculating the Concentration of FeSCN2+ in Equilibrium Solutions

After lab, working in your lab notebook, set up a table for your calculated results, using the table format shown in the Report Form (Table 2). Use the calibration equation obtained in Part 1c to calculate [FeSCN2+] in each of the equilibrium solutions E2 through E4.

c)   Calculating Keq

Working in your lab notebook, calculate the initial concentrations of Fe3+ and SCN- in solutions E2 through E4. From the measured concentrations of FeSCN2+ in each equilibrium solution (calculated in Part 1b), calculate the final concentrations of Fe3+ and SCN- (use an ICE table). Enter all results in the table.

For each equilibrium solution, compute Keq. Enter results in the table of your lab notebook. Determine the mean and standard deviation for your four determinations of Keq. In your notebook, comment on the level of agreement among your values.

 Pre-lab Assignment

Notebook Preparation

    In your lab notebook, prepare the following information:

·      A brief (2-3 sentence) objective of the lab.

·      A table of glassware, equipment and chemicals to be used. Include relevant properties and safety information for each chemical. Use this helpful link for online SDS https://chemicalsafety.com/sds-search/ to find out the hazards associated with substances you use and make in this experiment.

·      Several “bullet points” summarizing the tasks involved in the procedure.

.    Solutions to the Practice Problems - be sure to show your work!

Practice Problems - repeated from previous

1)  The absorbance of a solution of solute X in a 1.50-cm sample tube is 0.755. Use the Beer’s law plot above to determine the concentration of X in this solution. (Answer: 0.135 M)

2)  A calibration solution contains 1.50 mL of 0.00200 M NaSCN, 5.00 mL of 0.200 M FeCl3, and enough 0.100 M HNO3 to make a total volume of 25.0 mL. Assuming that [Fe3+] is sufficient to drive the reaction to completion, what is the concentration of FeSCN2+ in this solution? (Answer: 1.20 x 10-4 M)

3)  A student prepares an equilibrium solution by mixing 5.00 mL 0.00200 M Fe3+, 2.00 mL 0.00200 M SCN-, and 3.00 mL 0.100 M HNO3. What are the initial concentrations of Fe3+ and SCN- in this solution? (Answers: [Fe3+] = 1.00 x 10-3; [SCN-] = 4.00 x 10-4)

4)  By measuring the absorbance of the solution in Question 3, the student finds that [FeSCN2+] = 8.70 x 10-5. Set up an ICE table, using the information in Question 3 and this question. Use the ICE table to calculate the equilibrium concentrations of Fe3+ and SCN- in this solution. (Answers: [Fe3+] = 9.13 x 10-4; [SCN-] = 3.13 x 10-4)

5)  Use the equilibrium concentrations from question 4 to calculate Keq for the formation of FeSCN2+. (Answer: 3.04 x 102)

Be sure to read and study the Background material, along with the procedures, before coming to lab. Keep full, legible records of your work, data, and observations in your Laboratory Notebook.

 Experimental Videos

 (1) Preparation of calibration solutions 1 (C1-C5), 5.1 min

  (2) Preparation of calibration solutions 2 Use of volumetric pipettes, 7.7 min

   (3) Absorbance acquisition of calibration solutions, 9.9 min

   (4) Preparation and absorbance acquisition of equilibrium solutions (E1-E5), 12.5 min

Experimental Data

 (pdf)

References

Grading Rubric 

 (pdf)

Post Experimental Analysis complete the following REPORT FORM (pdf version) on your own.

Report Form

Name                                                                                                  CHY 116: Measuring an Equilibrium Constant

Keep complete records of your lab work, including observations and data, in your Laboratory Notebook. A student with your background should be able to repeat your work and compare their results against yours, using only your lab notebook.

Calculations and Results

Working in your lab notebook, complete all tasks requested in the Procedure.

Part 1. Calibration Data and Results

 1.) The following data table should be in your laboratory notebook and completed as you collect results (before lab, set up this table in your lab notebook):

Table 1. Calibration Data and Results (found on page          in your lab notebook)

 2.) Using Excel, prepare a Beer’s-law plot. Graph the data as absorbance (vertical axis) vs. concentration(horizontal axis).  Fit a straight line to the data, and obtain the equation of the line and  the regression coefficient (R2). Label the axes of the graph and give it an informative title from which the reader can readily understand the relationship represented on the graph. Make two copies of the graph. Hand in one with your notebook pages and report form, and tape the other one on the corresponding pages of your lab notebook.

3.) Enter the value you obtained for e, with proper units and significant digits:

e =                              

Part 2. Equilibrium Data and Results

1.) The following data table should be in your laboratory notebook and filled in as you collect results and complete calculations. Calculations can be done in Excel. If you choose to do this, be sure to include the spreadsheet showing the data.

Table 2. Equilibrium Data and Results (from page              in your lab notebook)

 2.) Find the mean (average) value of Keq from your four equilibrium solutions. Answer: _______________

3.) Find the standard deviation for your four values of Keq. 

Answer: _________________

4.) Are there any data points that are outliers that you could omit? If so, recalculate the mean Keq and standard deviation omitting this outlier.

5.) The literature Keq value is 280. How does your value compare to this? What could account for any discrepancies?

Part 3. Questions

1.) What are the units of A?  Show that the product ebC has the same units as does A.

2.) When you plot a graph of the absorbance (A) versus the concentration of a substance, how do you know if the data obey Beer's law?

Turn in as one pdf document:

1.) all pages from your laboratory notebook for this experiment, 

2.) your Excel spreadsheet including graphs, 

3.) all pages of this Report Form, and 

4.) a typed Summary.

Staple them into a single package. The following should be included in the summary: topic sentence describing the goal of the experiment; in the body of the summary, focus on the results and outcomes (not procedural details, not intermediate calculations, but final results); a final conclusion sentence. Your summary should consist of three paragraphs with each paragraph addressing the following information:

Paragraph 1: What was the purpose of the lab that you performed and what did you do?

Paragraph 2: What were your results and what was the error associated with the results. Include numerical results and how do your results compare to literature values (if appropriate include percent error from literature values)?

Paragraph 3: Provide an error analysis. Address the precision of the instrument(s) and method(s) that you used. Analyze the %RSD of the method(s) used if appropriate.