Determining an Equilibrium Constant Using Spectrophotometry
Adam Capriola
CHM 1122 Section 155
February 21, 2007
Introduction
A chemical reaction usually starts with reactants which react to yield products. Many
times the reactants are completely used up to make products. However, the reactants sometimes
do not completely turn into products. There is an equilibrium between the concentration of
reactants and products. At equilibrium, the reactants turn into product and the products
decompose into reactants at the same rate. This ratio of the products to reactants at equilibrium
is represented by the equilibrium constant, or K. K is found by taking the concentration and
order of the products and dividing by the concentration and order of the reactants. In this
experiment, iron(III) ion reacts with thiocyanate ion to produce thiocyanatoiron (III). The
reaction is represented by the following equation: Fe3+ + SCN- <----> FeSCN2+. Using a
spectrophotometer, the absorbance of FeSCN2+ is measured at different concentrations. The
absorbance in then put into Beer-Lambert’s law, A = εbc, to find concentration and ultimately
the equilibrium constant.
Experimental
First, a clean cuvette was obtained, rinsed, and filled three-fourths full with 0.5 M HNO3
solution. This was used as a the blank solution for the spectrophotometer, which was set at 447
nm. Next, a 100 mL volumetric flask was obtained and rinsed with distilled water. About 20
mL of 2.00 x 10-3 M KSCN solution was then dispensed into a clean, dry 50 mL beaker. 10.00
mL of this solution was pipeted into the volumetric flask. The volumetric flask was then filled
with distilled water to the line on the neck. The cap was put onto the volumetric flask and it was
agitated to ensure consistency of the solution. This solution was then transferred into a clean,
dry 250 mL beaker.
Next, about 20 mL of 1.00 x 10-1 M Fe(NO3)3 solution was dispensed into a clean, dry 50
mL beaker. 1.00 mL of this solution was pipeted into the 250 mL beaker with the KSCN
solution. The solution was then mixed with a glass stirring rod. A second clean, dry cuvette was
filled three-fourths full with this solution using a disposable Pasteur pipet. The
spectrophotometer was blanked with the cuvette filled with 0.5 M HNO3 solution and the
absorbance of the second cuvette was then measured and recorded.
The KSCN and Fe(NO3)3 solution was then poured back into the 250 mL beaker.
Another 1.00 mL of Fe(NO3)3 solution was the pipeted into the 250 mL beaker. The cuvette
used for measuring absorbance was filled with this solution using the disposable Pasteur pipet
and was rinsed twice. It was finally filled three-fourth full with the solution in the 250 mL
beaker. The spectrophotometer was blanked again with the 0.5 M HNO3 solution and the
absorbance of the solution in the other cuvette was recorded again. This process was repeated
until 10.00 mL of Fe(NO3)3 solution had been added to the 250 mL beaker.
Results
Molarity of stock KSCN solution, M
Volume of KSCN solution used, L
Molarity of diluted KSCN solution, M
Molarity of Fe(NO3)3 solution, M
2.00 x 10-3
0.01000
2.00 x 10-5
1.00 x 10-1
Mixture
Absorbance
[Fe*]
[SCN*]
[Fe(SCN)2+]
0.092
0.221
0.246
0.287
9.90 x 10-4 M
1.96 x 10-3 M
2.91 x 10-3 M
3.85 x 10-3 M
1.98 x 10-4 M
1.96 x 10-4 M
1.94 x 10-4 M
1.92 x 10-4 M
1.96 x 10-5 M
4.70 x 10-5 M
5.23 x 10-5 M
6.11 x 10-5 M
1
2
3
4
Volume of
Fe(NO3)3
added, mL
1.00
2.00
3.00
4.00
5
6
7
8
9
10
5.00
6.00
7.00
8.00
9.00
10.00
Mixture
1
2
3
4
5
6
7
8
9
10
0.437
0.401
0.505
0.508
0.518
0.527
[Fe3+] = [Fe*] [Fe(SCN)2+]
9.70 x 10-4 M
1.91 x 10-3 M
2.86 x 10-3 M
3.79 x 10-3 M
3.83 x 10-3 M
5.57 x 10-3 M
6.43 x 10-3 M
7.30 x 10-3 M
8.15 x 10-3 M
8.98 x 10-3 M
3.93 x 10-3 M
5.66 x 10-3 M
6.54 x 10-3 M
7.41 x 10-3 M
8.26 x 10-3 M
9.09 x 10-3 M
1.90 x 10-4 M
1.89 x 10-4 M
1.87 x 10-4 M
1.85 x 10-4 M
1.83 x 10-4 M
1.82 x 10-4 M
[SCN-] = [SCN*] [Fe(SCN)2+]
1.78 x 10-4 M
1.49 x 10-4 M
1.42 x 10-4 M
1.31 x 10-4 M
9.70 x 10-5 M
1.04 x 10-4 M
8.00 x 10-5 M
7.70 x 10-5 M
7.30 x 10-5 M
7.00 x 10-5 M
9.30 x 10-5 M
8.53 x 10-5 M
1.07 x 10-4 M
1.08 x 10-4 M
1.10 x 10-4 M
1.12 x 10-4 M
K
113.52
165.15
128.78
123.06
250.33
147.25
208.01
192.14
184.90
178.17
Average K: 169.13
Calculations
To find [Fe*], I took the volume of Fe(NO3)3 used in liters, multiplied by the molarity of
the Fe(NO3)3 concentration, 1.00 x 10-1 M, and divided by the total volume of the solution. For
example using the first mixture, (1.00 x 10-3 L) (1.00 x 10-1 M) / (0.101 L) = 9.90 x 10-4 M. I did
the same thing to find [SCN*]. For example using the mixture 1, (10.00 x 10-3 L) (2.00 x 10-3 M)
/ (0.101 L) = 1.98 x 10-4 M. To find [Fe(SCN)2+], I rearranged Beer’s law to c = A / (εb), with b
equaling 1 cm and ε equaling 4700 L mol-1 cm-1. Using the first mixture as an example,
c = 0.092 / (4700 L mol-1 cm-1 x 1 cm), c = 1.96 x 10-5 M. To find [Fe3+], I used the equation
[Fe3+] = [Fe*] - [Fe(SCN)2+]. For example, from mixture 1, [Fe3+] = 9.90 x 10-4 M - 1.96 x 10-5
M = 9.70 x 10-4 M. To find [SCN-], I used the equation [SCN-] = [SCN*] - [Fe(SCN)2+]. Using
the first mixture, [SCN-] = 1.98 x 10-4 M - 1.96 x 10-5 M = 1.78 x 10-4 M. To find K, I took the
concentration and order of the products and divided by the concentration and order of the
reactants. Using the concentrations from mixture 1, K = [Fe(SCN)2+] / ([Fe3+] x [SCN-]), K =
1.96 x 10-5 M / (9.70 x 10-4 M x 1.78 x 10-4 M), K = 113.52. Finally, to find the average K, I
added the 10 K values found and divided by 10.
Discussion/Conclusions
My results do not seem very accurate. Even during the experiment, my partner and I
agreed that our absorbance readings did not seem right. The readings jumped a significant
amount between the first and second readings and the fourth and fifth readings. After the
seventh reading, the absorbance readings seemed to be evening out, as the difference between
readings was getting smaller. When reading some of the absorbances, we took a second reading
because a few times the absorbance reading was smaller than the previous. This did not seem
right; the absorbance readings should have been going up each reading. Error may be attributed
to the fact that a few drops of solution escaped the 250 mL beaker during transfers between the
cuvette. Error could have also resulted from the beakers not being totally clean, or the pipetment
of solutions could have been inaccurate. This would have caused the volumes and
concentrations to be different than they actually were.