A=εbc 3
Where,
A = absorbance (from the spectrophotometer)
ε = molar absorptivity coefficient M-1cm-1
b = path length in cm
c = analyte molar concentration
Note: Deviations from Beer's Law can be caused by:[2]
(a) Stray light: there should be no stray light that is outside the wavelength used.
(b) Unequal light path lengths across the light beam.
(c) Unequal absorber concentration across the light beam.
(d) Changes in refractive index of the solution at high analyte concentration.
(e) Light-scattering by the sample solution, especially in turbid samples: results in a significant absorption signal even when the absorber's concentration is zero
(f) Shifts in chemical equilibrium involving the absorber as a function of concentration.
This experiment was subdivided into two parts. In the first part of the experiment, finding the calibration curve, a graph was plotted using the data obtained. The concentration of the FeSCN2+ will be used as the x value and the absorbance as the y value. Then, the equation of the best fit line has to be found using linear regression. While on the second part, the absorbance from the five unknown solutions was taken. The equilibrium constant is now then determined with the help of the data obtained. The equation of the best fit line was associated with the Beer-Lambert’s Law (Equation 3) to obtain the concentration of FeSCN2+ at equilibrium. The answer obtained was then subtracted to the initial concentrations of Fe3+and SCN- to get their respective equilibrium concentrations. In determining the equilibrium constant, the equilibrium concentrations of FeSCN2+, Fe3+and SCN- are plugged in to Equation 1. The aim of this experiment is to show that the value of the equilibrium constant of a reaction is always the same even if different conditions are applied.
Experimental Detail
This experiment aims to calculate the equilibrium constant for the reaction of Fe3+ and SCN-. The reaction is prepared in under different conditions to ensure that the value of the equilibrium constant is always the same. There are two parts in this experiment:
In the first part, a set of standard solutions and unknown solutions were prepared. The standard solutions are prepared by adding much greater concentration of Fe3+ to a lesser concentration of SCN-. This will ensure a very forward reaction so that the reverse reaction will not occur. This will make sure that all SCN- will be converted to FeSCN2+, because of this, the equilibrium concentration of FeSCN2+ is basically equal to the initial concentration of the SCN-. For the unknown solutions, equal concentrations of Fe3+and SCN- are made to react.
In the second part, the absorbance of the standard and the unknown are measured with the use of a spectrophotometer. The spectrophotometer was set to a wavelength of 447nm, a wavelength that is part of the visible light spectrum. The blank standard solution was first to be tested. This blank solution does not contain any SCN-. This is needed in order to account only the absorbance of FeSCN2+. In getting the absorbance, the cuvette where the solutions are placed is washed first with distilled water then the solutions to be tested. The cuvette is placed inside the spectrophotometer then the absorbance was recorded.
Results and Discussion
The combination of Fe3+, a yellow solution, and SCN-, a colorless solution, forms a blood-red complex, FeSCN2+. This complex appears blood red because of the formation of [Fe(NCS)(H2O)5]2+ [3]. After using the spectrophotometer for the standard solutions, the data obtained is:
Table I: The Calibration Curve
As seen in the table, the equilibrium concentration of FeSCN2+ is equal to the initial concentration of SCN-. The reason for this is the SCN- acts as the limiting reactant of the reaction.
Figure I: The Graph of the Calibration Curve
Regression Line: 5150x-0.294
R2value: 0.999
Using the standard solutions 3, 4, and 5, the obtained regression line is 5150x-0.294. The R2 value which is approximately equal to 1 affirmed that this equation best fits the regression line. It can be observed in the graph that as the concentration of FeSCN2+ increases, the absorbance of the solution increases. This meant that as the concentration of FeSCN2+ increases, it absorbs more light making it more blood-red. This calibration curve is of great importance because it can be related to the Beer-Lambert’s law (Equation 3) to determine the equilibrium concentration of FeSCN2+ in the unknown solutions. The slope being the εb, the x value being the [FeSCN2+]eq and the y value being the absorbance.
Table II: Absorbance of Unknown Solutions
Table II shows that increasing the initial amount of SCN- the increasing Absorbance same as the data from Table 1. The initial concentrations of the reactants, Fe3+and SCN-, was obtained through the use of the dilution formula,
M1V1=M2V2 [4]
Table III: Equilibrium Constant Determination
Average Keq= 253.77
Literature Value= 890[1]
Comparing Table II and Table III, it can be observed that the difference between the initial concentrations of the reactants is equal to the respective equilibrium concentration of FeSCN2+. Also, the calculated average is far from the literature value with a percent difference of 71.46.This difference brings out questions regarding what went wrong in the experiment.
One possible error that might have occurred is that the stock solutions used were too diluted or does not have the right concentrations needed. The solutions were possibly contaminated. Also, the solutions were probably not all be transferred to the beaker due to the transfer from one container to another.
Human errors can also be held accountable to this large percentage error. When using the spectrophotometer, fingerprints might have stained the cuvette causing it to be hazy. Lastly, bubbles might have formed inside the cuvette when the solutions were transferred which will affect the absorbance that is being recorded.
Even with this large error, we can see that spectrophotometry can be used to determine the equilibrium constant of a reaction.
Conclusion
Answers to Questions:
- Discuss the significance of the HCl in the solution preparation.
Answer: The HCl causes the yellow color of the Fe3+ to minimize, making the complex, FeSCN2+, be the only substance that will absorb light.
-
The concentration of FeSCN2+ in the standard solution is equal to the concentration of SCN- , the limiting reagent. Is this condition always true? If not, what is (are) the condition(s) for this to be true?
Answer: This condition always holds true. Given that SCN- is the limiting reagent in this reaction, addition of other reactants will only shift the equilibrium forward. If all of the SCN- is completely reacted and converted to FeSCN2+, there will be no further reaction.
-
Solutions containing Fe3+ are colored, thus absorb at the visible region. Explain why the absorbance readings in the experiment correspond only to the absorption of the complex, FeSCN2+.
Answer: It is because HCl is added to the solution. It reduces the intensity of the color of Fe3+ until it is colorless. Thus, only the color of the FeSCN2+ is considered. Also, a blank solution without any SCN- was first placed in the spectrophotometer, AUTOZERO was performed. The reason why only the complex, FeSCN2+, correspond to the absorbance readings.
-
Can distilled water which has zero absorbance be used as a blank solution instead of the Fe3+ solution?
Answer: No, because the Fe3+solution was used as blank to negate its color thus the absorbance readings are only due to the FeSCN2+ complex.
- Account for the difference between the literature value and the experimentally determined value of the equilibrium constant.
Answer: The percent error obtained was 71.46. This may be due to measurement discrepancies and improper solution preparation. Measurements were made in graduated cylinders instead of a more accurate pipette. Some apparatus were not completely dried and may have diluted the solutions.
Application
References
[1] Bodner, G.M. Common Ions and Complex Ions. December 2011. Purdue University. 15 January 2012.
[2] O’Haver, T. Instrumental Deviation from Beer’s Law. 2010. University of Maryland. 15 January 2012.
[3]