Kylie Case
Period 6
Emma McKee Rebecca Smith
Spectrophotometric Determination of an Equilibrium Constant
Purpose:
The purpose of this lab was to determine an equilibrium constant by observing and deriving
calculations from the reaction between aqueous iron III nitrate and potassium thiocynate.
Theory:
A big part of this lab was knowing (and in the end proving) that absorbance is directly
proportional to the concentration of a solution (Beer’s Law), and if Absorbance v. Concentration
is graphed, a straight line will result. The regression line’s equation is used to determine the
concentration of an unknown solution once the percent of T has been measured. The equation of
Beer’s Law (A=abc) is an efficient way of putting the relationship between absorbance, the path
length of the solution, and the concentration of the solution, into a numeric form, since it
displays the ideal condition in which these components are proportional to the absorbance.
Another important component in this lab was chemical equilibrium, which is the state in which
the concentration of both the reactants and products yield no net change with time (This
condition usually occurs when the forward reaction elapses at the same rate as the reverse
reaction). A big part of chemical equilibrium is reaction constant K=[C]c [D]d / [A]a [B]b
This equation can be calculated if the equilibrium concentration of each reactant and product in a
reaction at the equilibrium in known.
Procedure:
In this lab, a laptop and colorimeter were set up as indicated by the instructor, followed by the
group members obtaining the necessary testable chemicals. Next, five solution were prepared,
which had varied amounts of KSCN and distilled water, with a constant of five mLs of Fe
(NO3)3. Then the first solution was placed in a cuvette and the absorbance was tested using the
colorimeter, and this data was plotted on a graph. This was then done for the remaining
solutions, and using a compilation of the data, a best fit line was added to the graph, which was
then printed.
In part two, 5 mL of .2M Fe (NO3)3 were added to a 50 mL flask. The volumetric flask was filled
with distilled water and then mixed thoroughly. The value of the absorbance was found, and the
acquired data created a graph, was then interpolated to confirm the approximate value.
In part three, five test tubes were used to create new solutions yielding various amounts of SCNand distilled water, and the volume of the .002M Fe(NO3)3 was kept constant at 3 mL. The
colorimeter was used in order to find the absorbance of these solutions and collect data, and after
all data was collected, the solutions were disposed of.
Data:
Table 1. Parts I and II: Beer’s Law Data
Solution
1
.001
Absorbance
2
.255
3
.411
4
.567
5
.695
Unknown, Part II
.513
Best-fit line equation for Part I standard solutions: _y=3521x-0.008505____________________
Temperature: _____23.1°C___________________
Concentration of Unknown: _____1.48 * 10-4__________________________
Table 2. Part III: Absorbance Data
Test tube number
1
Absorbance
Net Absorbance
0.000
2
0.125
0.125
3
0.198
0.198
4
0.246
0.246
5
0.307
0.307
Table 3. Initial Concentrations of SCN- and Fe3+
Test tube #
[SCN-]i
2
4.00 * 10-4M
3
6.00*10-4M
4
8.00*10-4M
5
1.00*10-3M
[Fe3+]i
6.00 *10-4M
6.00*10-4M
6.00*10-4 M
6.00*10-4M
Table 4. Equilibrium concentration of FeSCN 2+
Test tube #
[FeSCN2+]eq
2
3
4
5
3.79*10–
Calculations:
1. Determine the initial concentration of SCN- and Fe3+ in the test tubes.
Trial 2
Fe3+
SCN-
Meq Veq= MiVi
Meq Veq= MiVi
Meq= MiVi/ Veq
Meq= MiVi/ Veq
Meq= (.0020)(3.00)/(.010)
Meq= (.0020)(.0020)/(.010)
Meq= 6.0 *10-4
Meq= 4.0*10-4
2. Use the net absorbance values in Part II and the best fit line equation in Part I to
determine the [FeSCN2+] at equilibrium.
Trial 2
Best-Fit Line Equation
Net Absorbance
Y=
Graph: See attached graph
Error Analysis:
In this lab, as every other lab, there were multiple potential sources of error. Had everything been
perfect, the Key values should have all been the same, which was not the case, though they were
in close relation to eachother. A large source of error in this experiment could have been not
making sure the test tubes were dry, as would have diluted the absorbance value which would in
turn lower the value of Keq. Another source of error could have been temperature, since it was
assumed that all solutions would remain at the same temperature, and body heat, or perhaps not
being the temperature of the room at the time of reading, could have altered results and
calculation.
Conclusion:
Throughout this lab, students determined an equilibrium constant by studying the reaction
between aqueous iron III nitrate and potassium thiocynate, which was the purpose of the lab.
Even though the results were not exactly as expected, the data helped the students derive values
that were respectably close to the desired readings. The results tended toward the uppers 100’s
lower 200’s, which were near what was expected. Overall, the students successfully carried out
their task and continued to familiarize themselves with the lab equipment, as well as hopefully
learn how to deal with the contingency of the temperature value fluctuating due to environmental
influences.