Experimental Determination of the Equilibrium
Constant for the Reaction of Iron(III) Ions and Thiocyanate Ions
Objective
Any chemical reaction will eventually proceed to a state of dynamic equilibrium, in which the reactants and products are present indefinite concentrations. In some cases, the equilibrium will be attained only after nearly all of the reagents have been consumed and converted to reaction products. Such a reaction is said to "go to completion."
Oftentimes however, a substantial fraction of the original reagent concentration remains after equilibrium has been achieved. In this experiment, you will determine the equilibrium constant, Keq, for the reaction of iron (III) ions and thiocyanate ions.
Introduction
In this experiment you will measure concentrations of the thiocyanate-iron(III) complex ion, Fe(SCN)2+, which forms when you mix solutions containing ferric ion, Fe3+ and thiocyanate ion, SCN-, as indicated by the following equation:
Fe3+(aq) + SCN-(aq) ↔ Fe(SCN)2+(aq)
Colorless Colorless Red-brown Complex
The equilibrium constant expression for this system is:
Keq =
The equilibrium system is prepared from known concentrations of Fe3+ and SCN-, carefully measured with pipets and mixed in volumetric flasks. Since the Fe(SCN)2+ formed is an intense red-brown colored complex ion with an absorption maximum at about 447 nm, its concentration can be determined with a spectrophotometer. The higher the concentration of FeSCN2+, the more intense is the red color of the solution.
By knowing the initial concentrations of Fe3+ and SCN-, and by estimating the Fe(SCN)2+ equilibrium concentration, the equilibrium concentrations of Fe3+ and SCN- can be calculated. Using these equilibrium concentrations, the Keq can be determined for this reaction.
You will first prepare a standard curve graph. This graph will allow you to extrapolate the equilibrium concentration of FeSCN2+ of a solution from its absorbance. The standard curve was prepared by using very high concentrations of Fe3+ to drive the reaction to completion by Le Chatelier’s Principle.
Next, you will prepare a series of solutions and measure the %T of each solution. After the percent transmission have been converted to absorbance data, you can use your graph to extrapolate the equilibrium concentration, [FeSCN2+]eq, for each solution. This gives us the numerator of the Keq equation above.
We need to calculate the equilibrium concentrations of Fe3+ and SCN- in order to calculate Keq. The equilibrium concentration of these reactants is just the amount of each reactant that is left over at the end of the experiment. Each reactant lost an amount equal to the amount of product that was formed since the stoichiometry is 1:1 in this reaction.
We can determine [Fe3+]o and [SCN-]o, the initial concentration of the reactants, from the number of milliliters of the stock solutions of each that was used in the reaction. We use the familiar M1V1 = M2V2 formula to calculate the initial concentrations.
If we use a table to organize this information, it will be clear how to determine the equilibrium concentration of Fe3+ and SCN-.
Solution B / Fe3+ / SCN- / FeSCN2+Initial: / Dilute conc. / Dilute conc. / 0.0 M
Change: / - [FeSCN-]eq / - [FeSCN-]eq / + [FeSCN-]eq
Equilibrium:
The dilute concentrations are the concentrations of each reactant after they are mixed together.
The equilibrium concentration of FeSCN2+ is extrapolated from the graph. You subtract this equilibrium concentration from the reactants and add this equilibrium concentration to the products. Initially there is no FeSCN2+ present.
Using the values you calculate on the Equilibrium row, you can calculate Keq.
The Standard Curve
The data you will need to prepare a standard curve is given in the table below. This data was determined experimentally by your laboratory instructors. Use Excel to graph the data and print a copy of the graph to use in class. One partner can do this while the other partner is preparing the test tubes for the next part of the experiment.
[FeSCN-] /Absorbance
5.00E-04 / 0.2077.00E-04 / 0.277
9.00E-04 / 0.354
1.10E-03 / 0.441
1.30E-03 / 0.519
1.50E-03 / 0.602
The Experiment
Turn on the spectrophotometer and set the wavelength to 447 nm. Allow the machine to warm up for 5-10 minutes.
Adjust the controls so that %T read 0% when no tube is in the spectrophotometer. Insert a cuvette containing solution A and adjust the control to read 100 %T. Record the %T of each of the remaining samples (B-D).
Solution / 0.0010 M Fe(NO3)3 / 0.0015 M NaSCN / 6.0 M HNO3A / 8.0 / 2.0 / 0.0
B / 6.0 / 2.0 / 2.0
C / 4.0 / 2.0 / 4.0
D / 2.0 / 2.0 / 6.0
The spectrophotometer expresses the amount of light absorbed in one of two ways. The first way of measuring the amount of light absorbed is called the percent transmittance, %T, which is defined as follows:
%T = I x 100
I o
The second way of measuring the amount of light absorbed is called the absorbance, A. It is defined as follows:
A = 2 - log %T
When a sample completely absorbs no light at all (all light is transmitted), %T = 100 and A = 0.000. As a sample absorbs a larger fraction of light (less light is transmitted), %T will decrease and A (the amount of light absorbed) will increase.
Calculations
1. Convert the %T data to absorbance data.
2. Calculate the dilute concentrations of Fe3+and SCN-.
MconcentratedVconcentrated = MdiluteVdilute
3. Using your graph, estimate the equilibrium concentration of FeSCN2+ for solutions B-D.
4. Calculate the equilibrium concentrations of Fe3+ and SCN- by subtracting the equilibrium concentration of FeSCN2+ from each dilute concentration. It is very helpful to fill out a table like this for each solution.
Solution B / [Fe3+] / [SCN-] / [FeSCN2+]Initial: / Dilute conc. / Dilute conc. / 0.0 M
Change: / - [FeSCN-]eq / - [FeSCN-]eq / + [FeSCN-]eq
Equilibrium:
5. Calculate the Keq for solutions B-D.
6. Calculate the average Keq for your reactions.