Determination of Equilibrium Constant for FeSCN2+

Introduction:

Chemical reactions can proceed in both the forward and reverse directions until a state of dynamic equilibrium is reached. The equilibrium constant (K) characterizes this balance between reactants and products at a specific temperature. In this laboratory experiment, we aimed to determine the equilibrium constant for the formation of the complex ion FeSCN²⁺ using a spectrophotometer.

The formation of FeSCN²⁺ can be represented by the following equilibrium reaction:

Fe³⁺(aq) + SCN^-(aq) ⇌ FeSCN²⁺(aq)

This reaction involves the combination of iron(III) ions (Fe³⁺) with thiocyanate ions (SCN^-) to form the red-colored complex ion FeSCN²⁺. The equilibrium constant (K) is defined as the ratio of the product concentration to the reactant concentrations, each raised to the power of their respective stoichiometric coefficients.

The equilibrium constant (K) is a crucial parameter in chemical equilibrium calculations, and it depends on temperature. In this experiment, we will use a spectrophotometer to measure the absorbance of solutions at a specific wavelength, which is directly proportional to the concentration of FeSCN²⁺. By analyzing the absorbance data of standard solutions with known concentrations and applying the Beer-Lambert Law, we can construct a calibration curve. This curve will allow us to determine the concentration of FeSCN²⁺ in test solutions of unknown concentration.

This lab report will outline the experimental procedure, data, calculations, and a discussion of the results. By obtaining the equilibrium constant (K) for this reaction, we aim to gain a deeper understanding of chemical equilibrium and the principles governing it.

Objective:

The purpose of this lab is to determine the equilibrium constant (K) for the formation of FeSCN²⁺ based on the absorbance recorded with a spectrophotometer.

Equipment:

• Spectrophotometer (wavelength set at 450 nm)
• Volumetric flasks
• Test tubes
• Stirring rod
• Iron(III) nitrate 0.20 M
• Iron(III) nitrate 2.0x10^-3 M
• Potassium thiocyanate, 2.0x10^-3 M
• Nitric acid 0.50 M
• Baking soda

Procedure:

1. Use pipets to measure the KSCN solutions listed in the available table into 100-ml volumetric flasks and dilute to 100 ml with 0.20 M Fe(NO₃)₃ in 0.50 M HNO₃. The more concentrated iron nitrate solution should be used.
2. Assuming all of the SCN has reacted, calculate the concentration of FeSCN in each flask.
3. Measure the absorbance of each of the standard solutions at 445 nm with the spectrophotometer.
4. Plot molar concentration of FeSCN²⁺ vs. absorbance and draw the line of best fit through the data points (including the origin).
5. Prepare the test solutions based on the quantities listed in the second table and then measure the absorbance of each at 445 nm.

Data and Calculations:

Calibration Curve for Absorbance of FeSCN²⁺

Test Solutions Data:

Discussion Questions:

1. The equilibrium constant (K) characterizes an equilibrium in a reaction based on the final concentrations (M) of the compounds that are involved in the reaction. The value of K was nearly constant for all of the experiments within a few digits. It should be consistent because the constant is defined as the relation between all of the concentrations of products and reactants at equilibrium.
2. The calculated value of the equilibrium constant indicates that there are mostly products since 54.5 is larger than the Haber process's constant of approximately 30, which contains mostly reactants.
3. A spectrophotometer is a device that measures the amount of light that can pass through a substance. In this experiment, we used it to determine the transmittance which we used to thus determine the percent absorbance for each solution. The standard solutions were obtained by mixing SCN- with an excess of Fe3+. This would allow us to assume that the reaction went to completion and so we would have a set amount of FeSCN2+. We then found the absorbances of the solutions to be able to create a graph that relates concentration to absorbance which was then used to find the unknown concentrations with the measured absorbances of the test solutions.
4. The spectrophotometer should not be set to the same color of the solution because the visible color is that of which is transmitted so none of the light would be absorbed. The complement of the solution's color should be used instead. In this experiment, 450 nm light was used, which is blue. Since the FeSCN2+ complex ion is red, this blue light can be absorbed.
5. The spectrophotometer that was used went to one significant figure. The major source of error in this experiment is the spectrophotometer itself due to it only allowing one significant figure with a small area of precision.
6. A spectrophotometer could be useful in measuring the concentration of a solution containing copper due to its blue-colored nature. The amount of red wavelength that has been absorbed would then be recorded.

Conclusion

This lab clearly demonstrated the process of determining the equilibrium constant. The average equilibrium constant was determined to be 54.5. Although the readings caused us to have an accurate final answer, the individual values differed slightly. The values for the equilibrium constants of each reaction ranged from a low of 47.50 to a high of 58.64, which is not a tremendous difference but it would still affect the average constant to a minor degree. The cause for these incorrect readings was most likely due to incorrect measurement in solutions. The accuracy of the spectrophotometer, the purity of the solution, and the accuracy of the line of best fit drawn on the graph could also have affected the data. If water was present in the cuvette, this could have made the solution appear lighter, skewing the data as well. The intense color of the FeSCN2+ complex ion makes the determination of its equilibrium concentration quite simple. The results of the absorbance and concentration of the standard solution were used to create a graph with which the line of best fit was found, enabling us to locate values for the absorbencies of the test solutions.