Lab 4 - Spectrophotometric Determination
of an Equilibrium Constant I: Iron (III) Thiocyanate*
April 25, 2000
*Adapted from "Chemistry in the Laboratory", 4th ed., J. L. Roberts,
Jr., J. L. Hollenberg and J. M. Postma, Freeman, 1997.
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 determine the ratio of iron to
thiocyanate that exists in the product, i.e., to determine the value of
x
in the above reaction. This will be done by mixing the concentrations of
these ions systematically. Provided that the total number of moles is constant,
the mix ratio that results in the highest concentration of the complex
corresponds to the ratio that exists in the compound. Once this is ascertained,
the value of the equilibrium constant can then be found.
Procedure. Prepare the solutions listed in the table below in
a series of labeled test tubes. Note that because the concentration of
the iron (III) ion and the thiocyanate are both 0.00200 M, the total number
of moles is the same in each case. Be careful when delivering these volumes.
Use a buret and record the amount of solutions added. If you add too much,
adjust the volume of the second reagent so that the volumes delivered will
be equivalent. Thoroughly mix the solutions to ensure uniformity.
| Sample Number | Volume 0.00200 M Fe3+
(mL) |
Volume 0.00200 M SCN-
(mL) |
| 1 | 10 | 0 |
| 2 | 8 | 2 |
| 3 | 6 | 4 |
| 4 | 5 | 5 |
| 5 | 4 | 6 |
| 6 | 2 | 8 |
| 7 | 0 | 10 |
Using the spectrophotometer, measure the absorbance spectrum of each
solution individually. Save each spectrum in a file. Save the solutions
until Part 1 of 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 lmax 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-. The plot should consist of two segments that intersect at the mole fraction of iron in the product. From this value, the coefficient x in the above equation can be obtained. Include the plot in your final report.
Part 2: Calculation of the Equilibrium Constant. Once the stoichiometry
of the reaction is determined, you can write the equilibrium expression
and define the equilibrium constant, K. To evaluate K, you
need to know the concentrations of all three species in solution at equilibrium.
This is done as follows:
[Fe(SCN)x(3-x)+]. 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 iron concentration can be calculated using the mass balance. This states that the total iron in solution at equilibrium must equal the number of moles added initially. That is,
total moles Fe3+ = moles Fe3+eq
+ moles Fe(SCN)x(3-x)+eq
Since you know the number of moles of iron that you added and you know the concentration of the iron(III) thiocyanate, the number of moles of iron can be calculated. Note that you will need to then calculate the concentration from the resulting value.
[SCN-]. This can be calculated in the same manner as the iron, except that you need to take into account that there might not be a 1:1 relationship between the moles of thiocyanate in the complex and the moles of complex, i.e., x might not be 1. Thus, the mass balance is:
total moles SCN- = moles SCN-eq
+ x (moles Fe(SCN)x(3-x)+eq)
Use the spectral data from solutions #2 thru #6 and determine the value of K for each. In your report, provide a table with calculated concentrations for each of the three species and the resulting K. Also report an average value along with a standard deviation. Comment on any descrepancies in the data and why you did not get the same value for each solution.