JWB-1 Spectrophotometric Determination of Kf
Formal Report
Heather Ryan
Abstract
Five solutions were created using different amounts of KSCN solution, 0.20 M Fe3+ ion
solution, 0.020 M Fe3+ ion solution, and 1.0 M HNO3 solution. The solutions were tested in a
spectrophotometer for absorbance. In testing the absorbance of various solutions, it was possible
to monitor concentration changes of iron thiocyanate complex and determine the equilibrium
constants of the solutions. The end result was an average equilibrium constant, for four out of
the five solutions, of 144.
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Introduction
The study of the interactions between photons and matter is known as spectroscopy. The
branch of spectroscopy used in this study was absorption spectroscopy, which is used to study
the position of equilibrium for complex-forming reactions (LabMaunual). When light is
absorbed by matter it causes an electron to move to a higher energy level. This is used to
measure the wavelength of the moving electron and since only particular wavelengths can be
absorbed by each solution used in the experiment, the wavelengths can be used to determine the
electronic structure of that solution (Lab Manual). The instrument used to measure these things
is a spectrophotometer. In this study a simple, manual spectrophotometer was used, which uses
light in the visible region and has a digital readout that shows both percent transmittance and
absorbance (Lab Manual). The measured result used in this experiment was the absorbance (A).
Beer’s Law states that at a given wavelength the absorption is directly proportional to the
concentration of the absorption species in the solution, which in this experiment was iron
thiocyanate complex. The absorption is also proportional to the thickness of the solution that the
light is passing through. The equation for Beer’s Law is:
in which ε is equal to molar absorptivity, b is equal to the optical path length of the curvette
which has the solution to be absorbed measured in centimeters, and c is equal to the molarity of
the solution that is going to absorb the light (Lab Manual). This is used for the absorbance, in
order to detect changes in concentration so that chemical equilibrium can be determined. As
previously stated, spectrophotometry was used to determine the formation equilibrium constant
of iron thiocyanate complex from the equation:
Fe+3(aq) + SCN-(aq)
FeSCN2+(aq)
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Using this equation formation constant can be determined with the formula:
The concentration of FeSCN+2 can be determined using the spectrophotometer, set at a
wavelength of 450 nm, because Fe3+ and SCN- do not absorb appreciable amounts of light but
FeSCN+2 does (Lab Manual). These measurements were used to determine the concentration of
FeSCN+2 in each solution. Use of spectrophotometry tested the absorbance in order to monitor
concentrations changes between solutions which made it possible to determine the equilibrium
constants (Lab Manual).
Experimental Method
Measured 10.0 mL KSCN solution and 10.0 mL 0.20 M Fe3+ ion solution separately
using a 10-mL volumetric pipet and a pipet bulb. Mixed the solutions in a small beaker.
Repeated using 10.0 mL KSCN solution and 10.0 mL 0.020 M Fe3+ ion solution; 10.0 mL KSCN
solution, 5.00 mL 0.020 M Fe3+ ion solution, and 5.0 mL 1.0 M HNO3; 10.0 mL KSCN solution,
2.50 mL 0.020 M Fe3+ ion solution, and 7.50 mL 1.0 M HNO3; and finally 10.0 mL KSCN
solution, 1.25 mL 0.020 M Fe3+ ion solution, and 8.75 mL 1.0 M HNO3. The 10.0 ml of KSCN
solution was added directly before placing the solution in the spectrophotometer to prevent
decomposition from occurring. Each mixture goes in a separate small beaker. Turned on the
spectrophotometer before use to warm it up, set the wave length to 450 nm. Filled a curvette
halfway with 1 M HNO3, to use as a blank, wiped the outside of the curvette with a Kimwipe and
placed the curvette into the slot in the spectrophotometer and pressed the “zero” button. Doing
so adjusts the instrument to 0.000 absorbance units. Filled another curvette with the first
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solution, placed in the spectrophotometer and recorded the value in absorbance units. Repeated
these steps for the remaining four solutions (Lab Manual).
Results and Discussion
Table 1 shows that as the absorbance of sample #2 through sample #5 decreases, the
concentration of FeSCN+2 also decreases. Also shown in Table 1, as the concentration of
FeSCN+2 decreases from sample #2 to sample #5, the constant (Kf) also decreases. This shows a
direct relationship between the absorbance, the concentration, and the equilibrium constant.
Table 1. Absorbance and Calculated Data for Iron(III) thiocyanate complex
Absorbance of Sample #1
First estimate of εb=
Sample # A
2
3
4
5
0.369
0.249
0.151
0.078
0.593
4235.7143
[FeSCN+2] CFE
8.71E-05
5.88E-05
3.56E-05
1.84E-05
0.01000
0.00500
0.00250
0.00125
[Fe +3]
-
[SCN ]
CSCNKf
0.00991 1.40E-04
5.29E-05 1.66E+02
0.00494 1.40E-04
8.12E-05 1.46E+02
0.00246 1.40E-04
1.04E-04 1.39E+02
0.00123 1.40E-04
1.22E-04 1.23E+02
Conclusion
Due to the calculations of the concentrations and equilibrium constant, the molar
absorptivity from the first solution was used, there is a possibility that the molar absorptivity
might not have been all the same for all the solutions (Lab Manual). This is due to the
assumption that SCN- had formed the complex in all the solutions, which might not have been
this case (Lab Manual). In order to account for that, the average equilibrium constant value was
calculated, 144, and would be used as the constant for all the solutions since concentration does
not affect equilibrium constant (Lab Manual).
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