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Evaluation of the Equilibrium
|In HNO3:||Fe3+ (aq) + SCN- (aq) FeSCN2+ (aq)||(1)|
for which the equilibrium constant expression is
To obtain the equilibrium constant, it is necessary to know the concentrations of all three ions present in an equilibrium mixture. The following information should allow you to develop a research plan for evaluating the equilibrium constant of the iron thiocyanate (lI) ion in aqueous solution.
Iron (III) has a coordination number of 6, meaning that there is a strong tendency for the Fe3+ ion to be surrounded by six molecules or ions. Aqueous solutions of iron (III) salts are generally yellow in color due to the presence of hydroxo-complexes of the iron (III) cation. These solutions become colorless in the presence of excess acid. The yellow hydroxo-complexes are converted to the colorless aqua-complex:
|[Fe(H20)5OH]2+ (aq) + H+ (aq) [Fe(H20)6]3+ (aq)||(3)|
The reaction to be studied is that of the colorless aqua complex with SCN-. Equation (4) is a complete representation of the reaction of equation (1)
|[Fe(H20)6]3+ (aq) + SCN- (aq) [Fe(H20)5SCN]2+ (aq) + H20||(4)|
It is customary to omit the molecules of water for the sake of simplicity.
The color of the complex ion, FeSCN2+, is sufficiently different from [Fe(H20)6]3+ and the SCN- ions so that a spectrophotometric method can be used to determine its equilibrium concentrations. The equilibrium concentrations of Fe3+ and SCN- can be found from the stoichiometry of the reaction and from a knowledge of the initial amounts of the reactants used to prepare the solutions.
Click here for prelab worksheet, Part 1. Please print using double-sided printer.
Click here for "prelab" worksheet,
(This worksheet will be done in lab, except the textbook problem which should be done before coming to lab.) Please print using double-sided printer.
1. Define the following terms in your own words:
DEVELOPMENT OF A RESEARCH PLAN:
Part 1—Establishment of the Beer’s Law Relationship for FeSCN2+ (The Standard Curve!)
Goal: Use the following questions/data table as a guide to plan Part 1 of your experiment where you establish the Beer's Law relationship for FeSCN2+ in which SCN- is the limiting reagent.
2. In order to determine the Beer's law relation for FeSCN2+, what type of data will you need to collect?
3. Given the following Materials and Equipment & Criteria below, fill in Table 1 on the next page to express how you would make 5 different solutions (where 1 of the 5 solutions is your blank) to establish your Beer's Law relationship (i.e. your standard curve). Hint: Read Part 1 of the experiment for help in filling out Table 1.
Your 5 solutions must fulfill the following criteria:
Table 1: Solutions needed for Beer's Law's Relationship for FeSCN2+Click to open a printable version of this table
|Fe3+ Vol. (ml)||Fe3+Initial Diluted Conc.** (M)||SCN- Vol. (ml)||SCN- Initial Diluted Conc.** (M)||HNO3Vol. (ml)||HNO3 Final Conc.* (M)||TOTAL Vol. (ml)||Ratio of Fe3+ to SCN- (should be greater than 650:1)|
|1.||1.2 x 10-4||0.5M||50.00|
*The final concentration of 0.5M HNO3 does not change since all the other solutions also contain 0.5M HNO3 as the diluent.
**Initial Diluted Conc. means the concentration immediately after reactants are mixed but before any reaction occurs to bring the system toward equilibrium.
4. When making the solutions in the table above, what glassware will you use for measuring each component? (Do not use volumetric glassware unless you need to.) If you use a volumetric flask, you will measure one of the components by filling the flask to the line.
5. Using your knowledge of color, predict an approximate value for the l max of an orange FeSCN2+ solution. [You will determine the exact l max in lab using one of your reactions and the Diode Array Spectrophotometer.
1. With your partner, devise a plan to determine the numerical value of the equilibrium constant for formation of the complex ion, FeSCN2+. To do this, you and your partner will need to create 3 different solutions/reactions that do NOT have an extreme excess Fe3+ to SCN- so as to determine the equilibrium constant in each solution and get an average equilibrium constant for the formation of FeSCN2+.Here are some hints to help you accomplish your goal!
Initial diluted concentration of reactants
Change(X) due to reaction going to equilibrium
Final equilibrium concentrations used to determine K
2. Complete the following textbook problem:
Chapter 6, Exercise 31, p. 220 Show all work, including an ICE table. (For assistance, see Example 6.3, p. 203).
Part 1: Suggested Strategy for Determining Beer’s Law Relationship for FeSCN2+
In developing your experimental plan, you must first establish a consistent relationship between the concentration of the complex ion and its’ absorption. Find the wavelength of maximum absorbance in the visible range. A suggestion for obtaining a known amount of the complex is to arrange experimental conditions so that the equilibrium is driven far to the right (an enormous amount of one reagent). Under these conditions, you can assume that all of the reagent in the lesser amount is converted to the complex. Test this idea experimentally by preparing three or four mixtures from the Fe3+ and SCN- solutions provided. You will also need to prepare a blank. The mixtures you prepare should have a large excess of Fe3+ (at least 650:1) with respect to SCN-. Show, experimentally, that all the SCN- has been converted to the complex. Dilute all solutions with nitric acid to keep the nitric acid concentration constant in all of them; the only thing that should be varied is the concentration of SCN-. Note that the HNO3 maintains the ionic strength (total concentration of all ions) of the solution. It also ensures that the Fe3+ stays in the colorless form, Fe(H2O)63+ .
Plot the absorbance at the appropriate wavelength versus concentration of FeSCN2+.
Has all the SCN- been converted to the complex? Justify your answer with lab data.
Part 2: Suggested Strategy for Determining Equilibrium Constant, K of FeSCN2+
Review the “hints” given in the Preparation sheet for Part 2. Having developed a method for determining the concentration of the complex ion in solution, prepare a solution for study that has roughly the same concentrations of Fe3+ and SCN- (plus the required nitric acid). Use the absorbance of this solution to determine the concentration of the complex ion (FeSCN2+). Knowing the initial amounts of Fe3+ and SCN- used to prepare the solution and the concentration of the complex from the absorbance data, calculate the concentrations of Fe3+ and SCN- present at equilibrium. Calculate the equilibrium constant. Repeat at least twice with different initial concentrations of Fe3+, SCN-, or both, to obtain at least three estimates of K. How can you use the plot from Part 1 to help you determine a final equilibrium concentration of [FeSCN2+] that you would like to make in Part 2?
Click here for results worksheet. Please print on double-sided printer.
Place all solutions in the waste beakers or bottles in the hood.