To begin with, in every reaction involving chemical substances, the reaction does not always go under a completion reaction. Instead, such a system reacts to some intermediate level where the reverse reaction gets equal to the forward reaction. At this intermediate state, the concentration of products and reactants becomes constant with changes in time. A system of that type is referred to as a chemical equilibrium. It is worth noting that at an equilibrium state, at specific temperature, the reaction mixture is subject to the Chemical Equilibrium Law, which provides a condition to guide the concentrations of both products and reactants. The condition is expressed in the equilibrium constant Kf for the reaction. This experiment involves the study of equilibrium properties for a reaction between thiocyanate ion and iron(III) ion as shown below:
Fe3+(aq) + SCN-
(aq) ↔ FeSCN2+ (aq)
It is worth noting that when a solution containing thiocyanate and Fe3+
ions are mixed, a reaction takes place between the two to an extend where FeSCN2+, a complex ion, that is commonly known for its deep red color is formed. Due to the reaction, the amounts of equilibrium for SCN- and Fe3+
is likely to be much less compared to a state of no reaction between the two; this implies that for every reacted mole of SCN-, one mole of Fe3+ is reacted to form one mole of FeSCN2+. An expression for an equilibrium constant Kf for the reaction is determined as follows:
Kf= [FeSCN2+]/[ Fe3+(aq)][ SCN- (aq]
It is worth noting that value of Kf is always constant at a particular value of temperature. This implies that all mixtures having SCN- ions and Fe3+ ions will always attain a state of equilibrium at the same value of Kf regardless of the initial amounts of SCN- ions and Fe3+ used in the experiment. Furthermore, the formation constant is useful in this experiment because it explains the complex formation, and we can use the equation to determine the Kf value for the resultant complex. The first solution has 1800 number Fe3+ larger than the initial SCN-, thus causing the reaction reach early completion, and the whole thiocyanate amount is converted or reacted to form the complex. Moreover, in the third solution, the iron (III) thiocyanate concentration is calculated as shown below: FeSCN2+=AAstd FeSCN2+
Purpose:
In this experiment, solutions comprising of SCN and Fe3+ are used to verify the formula for FeSCN2+ complex and an iron (III) thiocyanate constant of formation (Kf) is determined through spectrometry. It is worth noting that the overall electrolyte concentration is greatly affected by the solutions’ equilibrium constant.
Procedure:
Part A:
The procedure for this part involves using 40 mL beakers of each 3.0 x 10-3 M aqueous KSCN and 3.0 x 10-3
M aqueous Fe(NO3). The quantities required to prepare the solutions are measured using a graduated pipette. Moreover, the quantities are measured into 13 x 100 mm test tubes with stoppers. The absorbance is measured for the wavelength that absorbs the highest intensity of light, and the complex function or absorption spectrum is used to determine the wavelength absorption.
Part B:
In this part of the experiment, three dry, clean, labeled, separate beakers of about 10 mL 1.0 x 10 -3 M KSCN, 50 mL 0.5 M nitric acid and roughly 20 mL of 0.20 M Fe (NO3)3 are used in a way similar to that employed in part A to prepare solutions and obtain the respective absorbance at maximum wavelength. Then, the [FeSCN2+ ] are calculated for the measured absorbances.
Results:
Part A: Determination of the formula of the Fe(III)-SCN- complex
Table 1. Absorbance vs. reading for solution A3
(nm)
A3
400
0.976
425
0.837
450
1.054
475
0.820
500
0.250
525
1.133
550
0.592
575
0.820
600
0.407
625
0.226
650
0.197
675
0.076
700
0.029
Table 1 shows the absorbance versus the lambda readings for solution A3, which consisted of 5.0 of Fe(III) and 5.0 SCN-. This table shows that the lambda max is at 525 nm.
Table 2. Absorbance readings for all solutions at max
Solutions
Absorbance (nm)
A1
0.014
A2
0.004
A3
0.005
A4
0.005
A5
0.007
Table 2 shows the absorbance readings for solutions A1-A5 at lambda max (525 nm).
Part B: Formation constant for the Fe(III)-SCN- complex
Table 3. Absorbance for solutions B1, B2, B3, and B4 at max
Solutions
Absorbance (nm)
B1
0.006
B2
0.033
B3
0.015
B4
0.002
Table 3 shows the absorbance for solutions B1-B4 at lambda max (525 nm).
Results Analysis
Part A
Figure 1: Absorbance vs. λ
Figure 1 above indicates a curve of spectrum absorption for Fe(III)SCN- complex. It is important to notes that the maximum lambda is at 525 nm, which indicates absorption of most light.
Figure 2. Absorbance vs. Fe(III)
Figure 2 above indicates the plot of absorbance against the Fe(III). Therefore, according to the above curve, we aee able to determine the number of moles of this reactant at the maximum, which is 1.0 mL.
Figure 3. Absorbance vs. SCN-
Figure 3 above indicates the plot of absorbance against the SCN. It is important to note that the above curve shows the number of moles of this reactant was determined to be 2.0 mL.
It is worth noting that the stoichiometry can be obtained from each curve by dividing the maximum moles of reactant from the above figures (Figure 1 and 2). Therefore, upon dividing the two, we 0.5.
That is,
1 mL Fe (III)/ 2 mL SCN-= 0.5
Part B
Kf
B1:B2
25.27
B1:B3
136.15
B1:B4
66.4
Average
75.94
The value for B1:B3 is obtained as follows: Kf=3.7×10−29.27×10−16.4×10−5=272.3
Average Kf Value= (25.27+136.15+66.4)/ 3= 75.94
Uncertainty
Volume for B1: x 100% =-0.2162
[Fe3+]initial: ( x 100%) + B1 = -0.0.0371
[Fe3+]initial: ( x 100%) + B1 = -0.2634
[Fe3+]initial: ( x 100%) + B1 =-0.1081
[SCN-]initial: ( x 100%) + B1=- 0.1081
Therefore, uncertainty is (−0.1081)2+(−0.0101)2+(0.2634)2+(−0.0371)2+(−0.1081)25=0.0185
Standard deviation= 0.09147
Discussion:
In the above analysis, we obtained the formula for iron (III) thiocyanate using a spectrometer and measured the absorbance so as to calculate the constant of formation. According to figure 1 above, 525 nm was the peak value. We then used this lambda as the maximum for the other measurements. We used figure 1 and 2 to obtain the reactants’ ratio, which helped us understand the stoichiometry reaction. We used one of the solutions to obtain the other measurements since at one point most of the light was already absorbed at high sensitivity, and that is why we picked on a wavelength of 525 nm.
Conclusion
This laboratory experiment clearly shows the procedure of obtaining the equilibrium constant. According to the result analysis, the average equilibrium constant was found to be 75.94. It is worth noting that even though we were able to achieve a final accurate answer, the various individual readings slightly differed. The equilibrium constant values were within a range between 25.27 on the lower side and 136.15 on the higher side. This implies that the difference was not tremendous but it is still likely to have some impact on the average constant though to a small extend. Sources of errors during the experiment probably originated from incorrect solution measurement during the study. Moreover, other sources of errors may be originated from the spectrophotometers’ inaccuracy, presence of impurities in the solution, and inaccurate line of best fit on the constructed graph. In a situation where some drops of water were left in the cuvette, there is a high likelihood that the solution will assume a lighter appearance, thus causing the data skewing. Furthermore, it was quite easy to determine the equilibrium concentration of FeSCN2+ because of the its high intensity of color. In addition, the results of concentration and the absorbance of the standard solution were utilized in creating a graph with which the line of best fit was developed, which enabled us to obtain values for the absorbencies of the test solutions.
Works Cited
Bauer, Richard C, James P Birk and Douglas J Sawyer. Laboratory inquiry in chemistry. Belmont, CA: Brooks/Cole, CENGAGE Learning, 2009.
Hobbs, Peter Victor. Basic physical chemistry for the atmospheric sciences. Cambridge: Cambridge Univ. Press, 2000.
I︠A︡t︠s︡imirskiĭ, K B and V P Vasil'ev. Instability Constants of Complex Compounds. Boston, MA: Springer US, 1995.
