10. (Part 1) Equilibrium Constant of Iron Thiocyanate

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 (II) thiocyanate ion in aqueous solution. You will also need to generate a standard curve for iron (II) thiocyanate and it might be helpful to review the last time we used a standard curve (Lab 4) and why.

Background

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:

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)

It is customary to omit the molecules of water for the sake of simplicity (see equation 1).

The color of the complex ion, FeSCN²⁺, is sufficiently different from Fe³⁺ 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 then be found from the stoichiometry of the reaction and from a knowledge of the initial amounts of the reactants used to prepare the solutions.

Experiment – Part 1

Equipment and Materials

• • 0.2 M iron (III) nitrate[Fe(NO₃)₃]^.9H₂O in 0.5 M HNO₃ (known to more than 1 sig. fig. - be sure to record exact molarity in lab and use it in your calculations)

  • 2 × 10^-3 M potassium thiocyanate [KSCN] in 0.5 M HNO₃ (known to 4 sig. figs. - be sure to record exact molarity in lab and use it in your calculations)

  • 6M sodium hydroxide (NaOH)

  • 0.5M nitric acid (HNO₃)

  • iron (III) nitrate, Fe(NO₃)₃^.9H₂O_(s) (Molar mass = 404.00 g/mole)

  • 13 x 100mm glass test tubes

  • disposable plastic graduated pipets

  • Vernier Spectrophotometers with plastic cuvettes

  • 1.00-9.00 mL volumetric pipets

  • 50.00 & 100.00 mL volumetric flasks

  • 10, 25, and 100 mL graduated cylinders

Part A: Understanding Le Chatelier's Principle

Work in pairs unless otherwise instructed.

Le Chatelier's Principle states that, "When a system at equilibrium is subjected to a stress, the system will react so as to relieve the stress." Some examples of stresses that can be applied to a system are changes in concentration (both increasing and decreasing), pressure (for systems involving gases), and temperature. The system we will study is:

Fe³⁺ (aq) + SCN¯(aq) ⇄ FeSCN²⁺ (aq)

With this system, it is easy to observe the shifts in equilibrium because the reactants are colorless while the product has a distinct orange color.

Experiment A (Record all observations in Table 1 below.)

1. Prepare a solution, that we will label as "Control", using approximately 1.0 mL of 0.2M FeNO3, 5.0 mL of 0.002M KSCN and approximately 44 mL of deionized water.

2. Mix the solution well.

3. Fill four test tubes half-way (approx. 5 mL) with the Control Solution.

   • Do not add anything to one test tube. This is your Control.

   • Place one test tube in the heat block at approx. 95°C for 20-45 minutes to investigate the effect of heat on this reaction.

   • To a second test tube, add 5-10 drops of 6M NaOH with a disposable pipet. NaOH will react with Fe³⁺ and should cause an orange iron hydroxide precipitate to form. Wait for the orange precipitate to settle to the bottom of the test tube.

   • To the third test tube, add approximately 0.25 grams of solid iron (III) nitrate. Mix well and note if any solid remains undissolved.

4. Compare the color of the solution in the first three test tubes to the control and record the results in Table 1. Use a white background for observation. You may want to bring your Control test tube to the heat block when removing your test tube from the heat block in order to make an immediate comparison between the hot and room temperature solutions.

5. Save your solutions until the end of lab and make additional observations before disposing the solutions and rinsing the test tubes with water into the appropriate waste container.

6. Dispose of the rinsed-out test tubes into the "broken glass" box.

Questions for Thought:

• 1. The orange color of the product (FeSCN²⁺) allows us to monitor its concentration using Beer's Law. Using your knowledge of color, state a wavelength value (in nm) that would allow us to get the maximum absorbance value for our orange FeSCN²⁺ solution. (This wavelength is referred to as the λ_max.)

  2. If one reactant is added in extreme excess, the reaction will approach complete reaction. For example, if Fe³⁺ were in excess, the added concentration of SCN¯ would equal the concentration of the product, FeSCN²⁺. Does one of your solutions fit this description? If so, which solution(s)? Are these solution(s) at equilibrium?

  3. Devise a quick spectrophotometric experiment that can test your answer(s) to (2).

Part B: 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 at the wavelength of maximum absorbance (λ_max). A known amount of the FeSCN²⁺ complex can be obtained if the equilibrium is driven far to the right.

Which one of the conditions in Part 1 will cause the equilibrium to be driven far to the right (towards products)?

Under these conditions, all of the limiting reagent (SCN¯) should be converted to the iron thiocyanate complex.

You will now devise an experimental plan with your partner to measure the molar extinction coefficient of the iron thiocyanate complex. Recall Beer's law:

A = εlc

A = Absorbance (unitless)

ε = Molar extinction (or absorptivity) coefficient (cm^-1 M^-1)

l = Path length (cm)

c = Concentration (M)

A graph of absorbance vs. concentration will be linear, and the slope of this plot will give you the extinction coefficient. Thus, to measure the extinction coefficient, we need to make a standard curve, just like we did earlier this semester! Meaning we need to measure the absorbance of several solutions with known [FeSCN²⁺].

In the provided planning table below, please plan with your lab partner four different mixtures of the Fe³⁺ and SCN¯ solutions that satisfy the following criteria:

Here are the factors that will impact the design of your experiment:

• • Your solutions must be measurable with our spectrophotometers: 0.1 to 1.0 absorbance units.

  • You only have access to 1.00-9.00mL volumetric pipets and 50.00 mL volumetric flasks.

  • The mixtures you prepare should have a large excess of Fe³⁺ (at least 800:1) with respect to SCN¯.

We will be making some solutions that have known concentration. The first bullet point above dictates what concentrations you need to aim for. Perform some calculations, outlined below, to figure out the dilutions you will need to perform:

• • Assume that the extinction coefficient is approximately 4500 L/mol-cm. Use Beer's Law to figure out what concentration will cause A = 0.1, and what concentration will cause A = 1.0. All of your concentrations must be inside this range!

  • Next, look at the glassware available. How many mL of stock SCN- solution will you need to dilute to 50 mL in order to get a concentration somewhere within the range you calculated above? Note that since SCN- is the limiting reagent, and the stoichiometry of the reaction is 1:1, [SCN-] = [FeSCN²⁺]. In total, you will need to make four solutions, with four different concentrations inside the range calculated above.

  • Next, figure out how much Fe³⁺ to add. In all of your solutions, [Fe³⁺] should be constant! So you will need to determine how many mL of the Fe³⁺ stock to add so that the ratio of Fe³⁺:SCN- is at least 800:1 (it can be more!)

  • Finally, prepare your solutions!

    • Pipette your calculated volume of SCN- to each 50.00 mL volumetric flask using a volumetric pipette.

    • Add the same volume of Fe³⁺ to all flasks

    • Dilute to the line with nitric acid (HNO₃). This keeps the nitric acid concentration constant in all solutions.

    • Finally, you will also need to prepare a blank, which contains all solutions except SCN¯.

Use the table below to devise your test mixtures from the following stock solutions that will be available in our lab:

--> Stock [Fe³⁺] = 0.2 M in 0.5 M HNO₃

--> Stock [SCN¯] = 2×10^-3 M in 0.5 M HNO₃
(Be sure to record the exact concentration of the SCN¯ solution written to 4 sig figs once you get to lab.)

--> Stock [HNO₃] = 0.5 M

* This column in the table below is the concentration immediately after the reactants are mixed but before any reaction occurs to bring the system to equilibrium.

** The final concentration of 0.5 M HNO₃ does not change since all the other solutions also contain 0.5 M HNO₃ as the diluent.

Part B: Using Beer's Law to find a Molar Extinction Coefficient

Once you have planned your experiment with your partner and/or lab bench, please get together with everyone at your lab bench and select one set of solutions to make for your entire lab bench.

Submit your planning table (above) with your planned mixtures to your instructor for approval prior to proceeding with your experiment.

As a lab bench of students, please divide up the solutions to be made (including the blank) and create the 5 solutions from your table above and use them for the entire lab bench. Split back into groups of two and measure your solutions with a spectrophotometer to determine the absorbance of each solution. Each pair of students at a lab bench should create their own standard curve, even though you are sharing the solutions at the lab bench. Note: the absorption λ_max for FeSCN²⁺ is 454nm, yet what's most important is that you and your partner(s) use the exact same wavelength when measuring the absorbance of your solutions for this lab and next week's lab.

Question for Thought:

• 1. You want to know the extinction coefficient of FeSCN²⁺, so that you can determine [FeSCN²⁺] for your samples next week. So you need to plot A vs. [FeSCN²⁺]. How do you know what [FeSCN²⁺] is for your samples?

  2. Feel free to compare standard curves made using the same solutions. How similar are they? How similar should they be?

Part C: Develop a Plan for Determining Equilibrium Constant, K, of FeSCN²⁺

If time allows and your instructor asks you to, complete the Prelab assignment for next week.