Chemical Equilibrium:
Finding a Constant, Kc

The purpose of this lab is to experimentally determine the equilibrium constant, Kc, for the following chemical reaction:

Fe³+(aq) + SCN^–(aq) FeSCN²⁺(aq)

iron(III) thiocyanate thiocyanoiron(III)

When Fe³⁺ and SCN^– are combined, equilibrium is established between these two ions and the FeSCN²⁺ ion. In order to calculate Kc for the reaction, it is necessary to know the concentrations of all ions at equilibrium: [FeSCN²⁺]eq, [SCN^–]eq, and [Fe³⁺]eq. You will prepare four equilibrium systems containing different concentrations of these three ions. The equilibrium concentrations of the three ions will then be experimentally determined. These values will be substituted into the equilibrium constant expression to see if Kc is indeed constant.

You will use a Colorimeter or a Spectrometer to determine [FeSCN²⁺]eq,. The FeSCN²⁺ ion produces solutions with a red color. Because the red solutions absorb blue light very well, so Colorimeter users will be instructed to use the 470 nm (blue) LED. Spectrometer users will determine an appropriate wavelength based on the absorbance spectrum of the solution. The light striking the detector is reported as absorbance or percent transmittance. By comparing the absorbance of each equilibrium system, Aeq, to the absorbance of a standard solution, Astd, you can determine [FeSCN²⁺]eq. The standard solution has a known FeSCN²⁺ concentration.

To prepare the standard solution, a very large concentration of Fe³⁺ will be added to a small initial concentration of SCN^– (hereafter referred to as [SCN^–]i. The [Fe³⁺] in the standard solution is 100 times larger than [Fe³⁺] in the equilibrium mixtures. According to LeChatelier's principle, this high concentration forces the reaction far to the right, using up nearly 100% of the SCN^– ions. According to the balanced equation, for every one mole of SCN^– reacted, one mole of FeSCN²⁺ is produced. Thus [FeSCN²⁺]std is assumed to be equal to [SCN^–]i.

Assuming [FeSCN²⁺] and absorbance are related directly (Beer's Law), the concentration of FeSCN²⁺ for any of the equilibrium systems can be found by:

[FeSCN²⁺]eq = (Aeq/ Astd)x [FeSCN²⁺]std

Knowing the [FeSCN²⁺]eq allows you to determine the concentrations of the other two ions at equilibrium. For each mole of FeSCN²⁺ ions produced, one less mole of Fe³⁺ ions will be found in the solution (see the 1:1 ratio of coefficients in the equation on the previous page). The [Fe³⁺] can be determined by:

[Fe³⁺]eq = [Fe³⁺]i – [FeSCN²⁺]eq

Because one mole of SCN^– is used up for each mole of FeSCN²⁺ ions produced, [SCN^–]eq can be determined by:

[SCN^–]eq = [SCN^–]i – [FeSCN²⁺]eq

Knowing the values of [Fe³⁺]eq, [SCN^–]eq, and [FeSCN²⁺]eq, you can now calculate the value of Kc, the equilibrium constant.

OBJECTIVE

In this experiment, you will determine the equilibrium constant, Kc, for the following chemical reaction:

Fe³⁺(aq) + SCN^–(aq) FeSCN²⁺(aq)

iron(III) thiocyanate thiocyanoiron(III)

MATERIALS

PROCEDURE

Both Colorimeter and Spectrometer Users

1. Obtain and wear goggles.

2. Label four 20 x 150 mm test tubes 1– 4. Pour about 30 mL of 0.0020 M Fe(NO₃)₃ into a clean, dry 100 mL beaker. Pipet 5.0 mL of this solution into each of the four labeled test tubes. Use a pipet pump or bulb to pipet all solutions. CAUTION: Fe(NO₃)₃ solutions in this experiment are prepared in 1.0 M HNO₃ and should be handled with care. Pour about 25 mL of the 0.0020 M KSCN into another clean, dry 100 mL beaker. Pipet 2, 3, 4 and 5 mL of this solution into Test Tubes 1–4, respectively. Obtain about 25 mL of distilled water in a 100 mL beaker. Then pipet 3, 2, 1 and 0 mL of distilled water into Test Tubes 1– 4, respectively, to bring the total volume of each test tube to 10 mL. Mix each solution thoroughly with a stirring rod. Be sure to clean and dry the stirring rod after each mixing. Measure and record the temperature of one of the above solutions to use as the temperature for the equilibrium constant, Kc. Volumes added to each test tube are summarized below:

3. Prepare a standard solution of FeSCN²⁺ by pipetting 18 mL of 0.200 M Fe(NO3)3 into a 20 Í 150 mm test tube labeled “5”. Pipet 2 mL of 0.0020 M KSCN into the same test tube. Stir thoroughly.

4. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use cuvettes, remember:

• Wipe the outside of each cuvette with a lint-free tissue.

• Handle cuvettes only by the top edge of the ribbed sides.

• Dislodge any bubbles by gently tapping the cuvette on a hard surface.

• Always position the cuvette so the light passes through the clear sides.

Colorimeter Users Only (Spectrometer users proceed to the Spectrometer section)

5. Connect the Colorimeter to LabQuest and choose New from the File menu.

6. Calibrate the Colorimeter.

a. Place the blank in the cuvette slot of the Colorimeter and close the lid.

b. Press the < or > button on the Colorimeter to select a wavelength of 470 nm. Press the CAL button on the Colorimeter. When the LED stops flashing, the calibration is complete.

7. Set up the data collection mode.

a. On the Meter screen, tap Mode. Change the mode to Events with Entry.

b. Enter the Name (Concentration) and Units (mol/L). Select OK.

c. Proceed directly to Step 8.

Spectrometer Users Only

5. Connect the Spectrometer to LabQuest and choose New from the File menu.

6. Calibrate the Spectrometer.

a. Place the blank cuvette in the Spectrometer.

b. Choose Calibrate from the Sensors menu. The following message is displayed: “Waiting … seconds for lamp to warm up.” After the allotted time, the message will change to “Warmup complete.”

c. Select Finish Calibration. When the message “Calibration completed” appears, select OK.

7. Determine the optimal wavelength for creating the standard curve and set up the data-collection mode.

a. Empty the water from the blank cuvette. Using the solution in Test Tube 1, rinse the cuvette twice with ~1 mL amounts and then fill it 3/4 full. Wipe the outside with a tissue, place it in the Spectrometer.

b. Start data collection. A full spectrum graph of the solution will be displayed. Stop data collection.

c. The wavelength of maximum absorbance (λ max) is automatically identified. To choose a different wavelength, tap the graph (or use the ◄ or ► keys on LabQuest).

d. Tap the Meter tab. On the Meter screen, tap Mode. Change the data-collection mode to Events with Entry.

e. Enter the Name (Concentration) and Units (mol/L). Select OK.

Both Colorimeter and Spectrometer Users

8. You are now ready to collect absorbance data for the four equilibrium systems and the standard solution.

a. Leave the cuvette containing the Test Tube 1 mixture in the device (Spectrometer or Colorimeter). Start data collection.

b. When the value displayed on the screen has stabilized, tap Keep and enter 1 (the trial number). Select OK. The absorbance and concentration values have now been saved for the first solution.

c. Discard the cuvette contents as directed. Rinse the cuvette twice with the Test Tube 2 solution, fill the cuvette 3/4 full, and place it in the device. After the reading stabilizes, tap Keep and enter 2. Select OK to continue.

d. Repeat the procedure to find the absorbance of the solutions in Test Tubes 3, 4, and 5 (the standard solution).

e. Stop data collection. To examine the data pairs on the displayed graph, tap any data point. As you tap a data point, the absorbance and concentration values are displayed to the right of the graph. Record the absorbance values, for each Trial, in your data table.

f. Dispose of all solutions as directed by your instructor.

PROCESSING THE DATA

1. Write the Kc expression for the reaction in the Data and Calculation table.

2. Calculate the initial concentration of Fe³⁺, based on the dilution that results from adding KSCN solution and water to the original 0.0020 M Fe(NO₃)₃ solution. See Step 2 of the procedure for the volume of each substance used in Trials 1–4. Calculate [Fe3+]i using the equation:

[Fe³⁺]i = (Fe(NO₃)₃ mL/total mL)x (0.0020 M)

This should be the same for all four test tubes.

3. Calculate the initial concentration of SCN^–, based on its dilution by Fe(NO₃)₃ and water:

[SCN^–]i = (KSCN mL/total mL) x (0.0020 M)

In Test Tube 1, [SCN^–]i = (2 mL / 10 mL)(0.0020 M) = 0.00040 M. Calculate this for the other three test tubes.

4. [FeSCN²⁺]eq is calculated using the formula:

[FeSCN²⁺]eq = (Aeq /Astd) x [FeSCN²⁺]std

where Aeq and Astd are the absorbance values for the equilibrium and standard test tubes, respectively, and [FeSCN²⁺]std = (1/10)(0.0020) = 0.00020 M. Calculate [FeSCN²⁺]eq for each of the four trials.

5. [Fe³⁺]eq: Calculate the concentration of Fe³⁺ at equilibrium for Trials 1–4 using the equation:

[Fe³⁺]eq = [Fe³⁺]i – [FeSCN²⁺]eq

6. [SCN^–]eq: Calculate the concentration of SCN- at equilibrium for Trials 1–4 using the equation:

[SCN^–]eq = [SCN^–]i – [FeSCN²⁺]eq

7. Calculate Kc for Trials 1–4. Be sure to show the Kc expression and the values substituted in for each of these calculations.

8. Using your four calculated Kc values, determine an average value for Kc. How constant were your Kc values?

DATA AND CALCULATIONS