We used Fe3+(aq) + HSCN(aq) FeSCN2+(aq) + H+(aq) to investigate equilibrium. By measuring the equilibrium concentration of FeSCN2+ via absorbence of 445 nm light, we were able to find that the K for this reaction is 78 with a standard deviation of 31.
Equilibrium constants quantify the ratio between the reactants and products when the forward and reverse reaction rates are equal. If a chemical reaction is of the form: aA + bB cC + dD, the expression for K is .
The lab uses the spectrophotometer and the linear relationship between concentration and absorbence of the sample. The ability to determine concentration using this technique is critical to establishing the concentrations at equilibrium.
The use of a spectrophotometer requires that we have known concentrations of the FeSCN2+ ion. However, because of the equilibrium constant that we are trying to find, the might not fully react and form. Thus, Le Châtelier's Principle comes into use. The principle says that when there is an excess of chemical on one side of the reaction, the reaction proceeds toward the opposite side. We used this principle to force the reaction to almost go completely to FeSCN2+, thus we could now find the absorbence curve for FeSCN2+.
Our experiment used the reaction: Fe3+(aq) + HSCN(aq) FeSCN2+(aq) + H+(aq). The equilibrium constant K = . We will measure the concentration of FeSCN2+ because it has a reddish color that can be used in the spectrophotometer at 445 nm. The whole experiment would be conducted in a 0.5M HNO3 solutions with the exception of the distilled water cuvette used to zero the spectrophotometer.
The process to obtain the known concentrations of FeSCN2+ is complicated. Using stock KSCN, Fe3(NO3)3, and HNO3, we were able to create a known quantity of FeSCN2+. By limiting 2.010-3M KSCN to 0.2, 0.3, 0.4, 0.5, and 0.6 mL, and providing excess ferric nitrate, the reaction would proceed after creating the weak, intermediate acid, HSCN. Le Châtelier's Principle argued that in the 10mL of reacting solution, much of the SCN- would turn into HSCN because the whole solution is 0.5M HNO3. Furthermore, the principle will also push much of the HSCN into FeSCN2+ because of the excess Iron (III). The class as a whole created the solutions and measured the absorbences. The solutions were allowed to react at room temperature for several minutes. Stirring was not performed. Adequate mixing occurs when the pipette pours in the 0.2M stock ferric nitrate solution.
The experiment then involved various mixtures of 210-3M Fe(NO3)3, 210-3M KSCN, and 0.50M HNO3 (See Table below for mixtures). The tubes were allowed to stand for several minutes and then their absorbences were taken.
Tube 1 Tube 2 Tube 3 Tube 4 Tube 5 Units Fe(NO3)3 5.0 5.0 5.0 5.0 5.0 mL KSCN 1.0 2.0 3.0 4.0 5.0 mL HNO3 4.0 3.0 2.0 1.0 0.0 mL
Table 1: Amounts for the Experimental Tubes
The lab involved use of acids and nitrates. Both cause irritation and the nitric acid could, at higher concentrations burn skin. Aprons and goggles were worn at all times. The eye wash and acid neutralizing kit were available in the case of any spills.
Tube 1 Tube 2 Tube 3 Tube 4 Tube 5 Units KSCN 200 300 400 500 600 microliters [KSCN] 0.000040 0.000060 0.000080 0.00010 0.000120 M Data 0.18 0.27 0.39 0.44 0.53 absorbence Collected 0.19 0.28 0.36 0.49 0.56 absorbence 0.19 0.29 0.49 0.55 absorbence 0.55 absorbence Average 0.19 0.28 0.38 0.47 0.55 absorbence
Table 2: Data for Graph of Absorbence vs. [FeSCN2+]
The best fit line is: Absorbence = 4600[FeSCN2+]
Tube 1 Tube 2 Tube 3 Tube 4 Tube 5 Units [Fe(NO3)3] 0.0010 0.0010 0.001 0.0010 0.0010 M [KSCN] 0.00020 0.00040 0.0006 0.00080 0.0010 M [HNO3] 0.50 0.50 0.5 0.50 0.50 M
Table 3: Initial Concentrations
Tube 1 Tube 2 Tube 3 Tube 4 Tube 5 Units Absorbence 0.19 0.21 0.31 0.39 0.45 absorbence [FeSCN ion] 0.000041 0.000045 0.000067 0.000084 0.000097 M [H ion] 0.50 0.50 0.50 0.50 0.50 M [Fe ion] 0.00096 0.00096 0.00093 0.00092 0.00090 M [HSCN] 0.00016 0.00036 0.00053 0.00072 0.00090 M K 130 67 67 64 59
Table 4: Final Concentrations
To find [FeSCN2+] = .
The average K of 78 with a standard deviation of 31. The large standard deviation is caused by the Tube 1's K which tends to pull up the average from 64. All values reflect significant figures. Percent error cannot be calculated because we were not given an established K for the lab's ambient temperature and pressure.
Our experiment found that K was 78. The expected result was in the arena of 50. Considering the large standard deviation, 31, our result comes within a standard deviation of the expected result. We had expected the reaction to move a bit forward (thus a K greater than 1) but not extremely forward (thus a K smaller than 106).
The greatest errors in this experiment are the procedural errors where we put too much or too little of the reactants. The handling, including the lack of stirring, could play some minor roles. The spectrophotometer drifted only 0.02 Absorbence from start to finish of our readings. The drift is not much of a concern since our Tube 1 reading is already far from the other tubes.
Some approximations were made during the analysis. Molarity was manipulated during the calculations of the equilibrium concentrations instead of conversion into moles then back to molarity. Furthermore, the origin was assumed to be a point during the plotting of the absorbence line for FeSCN2+ (the calculations without the origin only show 0.0065 y-intercept value which is minuscule compared to our absorbence readings on the order of 0.30). During the hand calculations, the concentration of H+ ions were approximated to remain 0.50 M. The approximation neglects the change on the order of 10-5 which is well within the tolerances of two significant figures.
The K value should be the same at equilibrium. Some random fuzz of the K values can be expected but the standard deviation of 31 is large. It suggests that the Tube 1 sample is either flawed or had the wrong amount of reactants. Our value of K, 78, tells us that there are more products at equilibrium.