Laboratory #8:

Spectrophotometric Determination of an Equilibrium Constant I: Iron (III) Thiocyanate*
*Adapted from "Chemistry in the Laboratory", 4th ed., J. L. Roberts, Jr., J. L. Hollenberg and J. M. Postma, Freeman, 1997.

August 14, 2001
 

Introduction: The compounds of transition metals are often highly colored. The intense pigment cobalt blue is an example. Small amounts of chromium give rise to the red color of rubies and the green color of emeralds; the blue color of sapphires is produced by trace amounts of iron. In solution, transition metals also can have deep colors. For example, when Fe³⁺ is mixed with the thiocyanate ion, SCN^-, a deep red color results. The reaction is:

The purpose of this experiment is to measure the equilibrium constant of the above reaction. This will be done by mixing solutions of known concentrations of Fe³⁺ and SCN^- and measuring the amount of product that forms. Since the product is the only colored species in the mixture, its concentration can be determined spectrophotometrically using Beer's Law as described below.

Procedure. Using deionized water as a reference, measure the absorbance of the Fe(NO₃)₃ solution and the KSCN solution at 450 nm. Record these values in your notebook.

Next, label a series of nine test tubes and prepare nine solutions of by varying the amounts of iron and thiocyanate in a mixture, as listed in the table below. (Be careful, each solution is also 0.5 M nitric acid) Use a buret to deliver the solutions and record the exact amounts of solutions added; it is not important to deliver each volume exactly as written below, but it is important to know the volumes delivered to the nearest tenth of a milliliter. Thoroughly mix the solutions to ensure uniformity.
 
 

Measure the absorbance spectrum of each solution individually. Save the solutions until the analysis is complete - this is so you can reanalyze the solutions if you suspect the validity of some of your measurements. Once Part 1 is complete, pour all of the solutions in a beaker, neutralize with sodium bicarbonate and dispose down the drain.
 

Analysis.

Part 1: Determination of Mole Ratio. Plot the absorbance at 450 nm as a function of mole fraction of Fe³⁺. The mole fraction of Fe³⁺ is simply the number of moles of the iron (III) ion divided by the total number of moles in the solution, i.e., moles Fe³⁺ plus moles SCN^-.  Thus, the range on the x-axis will be from 0 (the pure SCN^- solution) to 1 (the pure Fe³⁺ solution). Include the plot in your final report.

Part 2: Calculation of the Equilibrium Constant. Based on the reaction stoichiometry, above, write the equilibrium expression for the reaction. Evaluate the value of the equilibrium constant, K, by using the known initial concentrations of Fe³⁺ and SCN^- and the measured concentration of iron thiocyanate complex. This is done as follows:
 

[Fe(SCN)²⁺]. Using Beer's Law, where e = 4700 M^-1cm^-1 and b, the pathlength, is 1 cm, you can determine the concentration of the iron (III) thiocyanate. Note that this only works because the Fe³⁺ and the SCN^- are both colorless. If they were to absorb light, this analysis would be incorrect because the total absorbance would be due to multiple compounds.

[Fe³⁺]. The equilibrium iron concentration is equal to the initial concentration minus the concentration of iron thiocyanate that forms during the reaction.

Note that the initial concentration of iron will not be 0.00200 M; you need to calculate the iron concentration for each sample you prepare using v₁c₁=v₂c₂.

[SCN-]. This can be calculated in the same manner as the iron. Thus,

Use the spectral data from solutions #1 thru #9 to determine the value of K for each.

Report. Turn in the plot of absorbance vs. mole fraction of Fe³⁺as described above. Prepare a table that provides the volumes used to prepare each solution, the resulting initial concentrations of iron and thiocyanate as well as the equilibrium concentrations and calculated value of K. Also report an average value of K along with a standard deviation.