Name______Date______

Determination of an Equilibrium Constant

When light is passed through a sample, some of the light may be absorbed. The absorbance, A, of a substance is directly proportional to the concentration of that substance according to Beer’s Law, A =  b c. Here “” is the molar absorptivity of the substance at any given wavelength, “b” is the pathlength of the sample in centimeters, and “c” is the concentration in moles/L.

The reaction between iron (Fe3+) and thiocyanate (SCN–) forms a deep red complex which absorbs the most light (has an absorbance maximum) at 447 nm.

Fe3+ + SCN–  FeSCN2+

In order to calculate the equilibrium constant for this reaction, in Part A a series of solutions of known concentration of FeSCN2+ are made and the absorbance of each solution is measured. A high iron concentration ensures that the equilibrium is driven all the way to the right and that all the thiocyanate is converted to the product. The initial concentration of thiocyanate is then equal to the final concentration of iron thiocyanate complex.

Once the concentration of SCN– is calculated for each solution, a calibration curve is made to determine the relationship between absorbance and iron thiocyanate concentration. This curve is then used to determine the concentrations of iron thiocyanate complex in Part B.

A. Establishing a Calibration Curve

1. Prepare a set of standard solutions. Pour 15 mL of 0.00200 M KSCN into a 50 mL beaker, 60 mL of 0.200 M Fe3+ into a 100 mL beaker, and 200 mL of 0.1 M HNO3 into a 400 mL beaker. Using a volumetric pipet for the KSCN and a graduated cylinder for the Fe3+, add the appropriate amounts to the 50 mL volumetric flask to make the solutions listed in the chart below. Once the flask is filled to the 50 mL mark with a solution, cap it and invert to mix thoroughly. Once each solution is prepared, measure its absorbance (or pour it into a labelled beaker if you are not ready to measure absorbance), then rinse the volumetric flask with deionized water and use it to make the next solution.

Solution / 0.00200 M KSCN / 0.200 M Fe3+ / 0.1 M HNO3 / Absorbance
blank / 0 mL / 10 mL / dil. to 50 mL / 0.00
#2 / 1 mL / 10 mL / dil. to 50 mL
#3 / 2 mL / 10 mL / dil. to 50 mL
#4 / 3 mL / 10 mL / dil. to 50 mL
#5 / 4 mL / 10 mL / dil. to 50 mL

2. Calibrate the spectrophotometer. Make sure the wavelength is set to 447 nm. Place a sample of your “blank” – the solution with 0 mL of thiocyanate solution – in a cuvette. Wipe the cuvette and put it in the sample chamber. Now zero the instrument: adjust the transmittance to 100 with the right front knob, remove the cuvette, and use the left front knob to adjust the transmittance to zero. Use the mode button to change from transmittance to absorbance.

3. Measure the absorbance of the standard solutions. Rinse the cuvette several times with a small portion of the solution being measured. Fill the cuvette approximately ¾ full, wipe the outside, put the cuvette into the instrument, and measure the absorbance. Repeat with each solution. Use the same cuvette for each measurement.

4. Graph the data. Plot absorbance vs. [FeSCN2+]. Create a best-fit line (trendline) that goes through zero.

B. Absorbance for a Set of Test Solutions

1. Prepare the solutions. Into three small beakers measure approximately 40 mL of 0.00200 M Fe3+, 20 mL of 0.00200 M KSCN, and 20 mL 0.1 M HNO3. In a 10-mL volumetric flask prepare the solutions below. Use volumetric pipets for the measurements. After you prepare a solution, measure its absorbance, then reuse the volumetric flask for the next solution.

Solution / 0.00200 M KSCN / 0.00200 M Fe3+ / 0.1 M HNO3 / Absorbance
blank / 0 mL / 5 mL / dil. to 10 mL / 0.00
#1 / 1 mL / 5 mL / dil. to 10 mL
#2 / 2 mL / 5 mL / dil. to 10 mL
#3 / 3 mL / 5 mL / dil. to 10 mL
#4 / 4 mL / 5 mL / dil. to 10 mL

2. Recalibrate the spectrophotometer. Use the blank solution.

3. Measure the absorbance of the test solutions. Rinse the cuvette with solution #1. Fill the cuvette ¾ full and wipe the outside. Measure the absorbance. Repeat with solutions 2 – 4.

4. Determine the concentration of the FeSCN2+ complex. Use the equation for the trendline from the calibration curve from part A and the absorbances measured here to calculate the concentration of FeSCN2+ at equilibrium. Once you have determined the equilibrium concentrations of reactants and products, calculate Kc for each trial and an average Kc.

(Hint: use an ICE chart to calculate the equilibrium concentrations of the reactants!)

Calculations

Test Solution / #1 / #2 / #3 / #4
[FeSCN2+] from calib. curve / ______/ ______/ ______/ ______
initial [Fe3+] / ______/ ______/ ______/ ______
equilibrium [Fe3+] / ______/ ______/ ______/ ______
initial [SCN–] / ______/ ______/ ______/ ______
equilibrium
[SCN–] / ______/ ______/ ______/ ______
[FeSCN2+] [Fe3+][SCN–] / ______/ ______/ ______/ ______
average Kc / ______

Lab Report

Include your calibration curve in your data section. It is not necessary to include an experimental procedure.

Questions

1. How does a dirty cuvette (caused by fingerprints, lint, etc.) affect the absorbance of a solution?

2. If all the measurements were done with a dirty cuvette, how would that affect K?

3. If the test solution measurements (but not the calibration curve measurements) were done using a dirty cuvette, how would that affect K?