Just for fun!
Where else do you see iron thiocyanate and pseudo-first order reactions?  
Here are some titles of articles that use one of those concepts in their scientific work.  Notice the cool titles of the journals…if you can read them?! 

*Noroozifar, M.; Khorasani-Motlagh, M.; Rafahmand, A.-R.   Application of  iron (III)- thiocyanate  complex in the spectrophotometric investigation of ascorbic acid.    Asian Journal of Spectroscopy (2003),  7(2),  81-86.

*Litvinenko, G. P.; Natansov, M. E.; Pyasik, I. B.   Increasing the wear resistance of the effective area of teeth of gear drives using  iron   thiocyanate.    Problemy Treniya i Iznashivaniya  (1974),  6  123-6.

*Pandey A; Iyengar L   Modification of arginine residues at the substrate binding site of yeast glutathione reductase. Indian journal of biochemistry & biophysics  (1998),  35(3),  157-60.

Equilibrium is EVERYWHERE!  
Here are just a few from David Brooks' teaching website:
*The “bends” for scuba divers.
*The binding of oxygen to hemoglobin.
*The “Predator-prey” system in nature.

Jokes
Q: What do you call a tooth in a glass of water?
A: One molar solution.

Q: Why do chemists like nitrates so much?
A: They're cheaper than day rates.

Laughing face image taken from Microsoft Office Clip Art.

Evaluation of the Equilibrium
Constant of Iron Thiocyanate

Goals

• Develop a better understanding of equilibrium, including the effects of concentration and temperature on equilibrium position.


• Determine the numerical value of the equilibrium constant for formation of the iron thiocyanate(II) ion,



In HNO3:	Fe3+ (aq) + SCN¯(aq) FeSCN2+ (aq)		(1)
		orange

for which the equilibrium constant expression is

(2)

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.

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:

	[Fe(H20)5OH]2+ (aq) + H+ (aq) [Fe(H20)6]3+ (aq)		(3)
	yellow	colorless

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)
		orange

 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.

Preparation – Week 1

Questions:

Fill out the prelab worksheet that can be found after the Experiment section for Week 1. One prelab is due for each week of the two-week experiment.

Print:

Print and bring to lab the Prelab for next week in pdf or Word format. Please print double-sided. Alternatively, bring your computer and plan to do the prelab electronically.

Experiment – Week 1

To print instructions, select the portion that you with to print, choose File/Print, and choose "selection" to prevent printing the entire document.

Safety

• Wear safety glasses
• Pour all waste solutions into the appropriate waste container.

Equipment and Materials

• 0.2 M iron (III) nitrate[Fe(NO₃)₃] 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 x 10^-3 M 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(s) 
• Spectronic 20 with test tube cuvettes
• 1-5 mL volumetric pipets
• 10, 15, & 25 mL volumetric pipets (week 2 only)
• 50 & 100 volumetric flasks
• 10, 25, and 100 mL graduated cylinders

Experimental

Part 1: Understanding Le Chatelier's Principle

Work in pairs unless otherwise instructed.

AIM: To investigate the effects of temperature and concentration upon the position of equilibrium.

Many chemical reactions do not go to completion. In these reactions, there will be measurable amounts of both reactants and products, even after a considerable period of time. For reversible reactions, an equilibrium is established and the ratio of products to reactants will remain constant if temperature is kept constant. It is possible to shift the equilibrium in a desired direction by applying a stress to the system. This process is explained by Le Chatelier's Principle,  which 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. In this experiment the effects of changes in temperature and concentration will be observed. The system studied is:

Fe³⁺ (aq) + SCN¯(aq) FeSCN²⁺ (aq)

Experiment 1

1. Prepare solution 1, described in the table below:

   	0.2M Fe3+ in 0.5M HNO3		0.002M SCN¯ in 0.5M HNO3
   	Fe3+  volume (mL)	Fe3+  conc* (M)	SCN¯  volume (mL)	SCN¯  conc* (M)	water volume (mL)	TOTAL volume (mL)	Molar ratio of Fe3+ to  SCN¯
   1.	1.0		5.0		44	50.0

   * This is the concentration immediately after the reactants are mixed but before any reaction occurs to bring the system to equilibrium.
   
   
2. Mix the solution well.
3. Fill three test tubes half-way with solution 1.
1. Place one test tube in the95°C heat block for 5 minutes.
2. Add 5-10 drops of  6M NaOH to the second test tube with a disposable pipet (iron hydroxide precipitate should form). Wait for the precipitate to settle before recording observations. This may take a while.
3. Use the third test tube as a comparison.
4. For each of your test tubes, record your observations.
5. Solution 2 starts with solution 1, and perturbs the equilibrium by adding Fe³⁺. Prepare solution 2 by putting 5.0 mL of solution 1 in a test tube and adding the appropriate amount of iron (III) nitrate[Fe(NO₃)₃] so that the initial Fe³⁺ concentration is 0.18M (concentration before reaction occurs to bring the system to equilibrium; you calculated this amount in your prelab). 
6. Mix well. Note any remaining solid.
7. Record your observations for solution 2.
8. Save your solutions until the end of lab.

Data table of observations:

Questions for Thought:

1. Explain the chemical basis for your observations when you compare the two room temperature solutions. In your comparison, explain how the equilibrium responded and why.
2. Based on your observations when heat was introduced to the system, is the reaction endothermic or exothermic? Explain how the equilibrium responded and why.
3. Fe³⁺ in the presence of hydroxide ions forms an insoluble precipitate. Did this happen when you added the NaOH solution to the equilibrium system? Explain how the equilibrium responded and why. 
4. The orange color of the product allows us to monitor its concentration using Beer's Law. Using your knowledge of color, predict a range of wavelength values that should describe the λ_max for the orange FeSCN²⁺ solution.
5. 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?
6. Devise a quick spectrophotometric experiment that can test your answer(s) to (5). 

Part 2: 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. A known amount of the complex can be obtained if the equilibrium is driven far to the right (an enormous amount of one reagent). Based on the knowledge gained in Part 1, you should have confidence that under these conditions, all of the limiting reagent (SCN¯) is converted to the complex. In the provided planning table below, plan four different mixtures of the Fe³⁺ and SCN¯ solutions that satisfy the following criteria:

1. The mixtures you prepare should have a large excess of Fe³⁺ (at least 800:1) with respect to SCN¯. 
2. Only the concentration of SCN¯ is allowed to vary. This means that you must:
1. Keep the concentration of Fe³⁺ constant.
2. Dilute all solutions with nitric acid (HNO₃) to keep the nitric acid concentration costant in all of them; 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⁺.
3. The added SCN¯ concentrations must yield a colored product within an absorbance range of 0.15 to 1.0. Assume that the extinction coefficient is approximately 4500 L/mol/cm, and use Beer's Law to calculate the concentration range needed for SCN¯. Perform this calcualtion first. Ultimately, volumes of SCN¯ stock will be delivered with the volumetric pipets available. Therefore, once the stock concentration of 2x10^-3 M and total volume of 50.00 mL are additionally defined, as well as the available volumetric pipettes, only a finite number of concentrations of SCN¯ are possible.
4. Lastly, you will also need to prepare a blank (all solutions except SCN¯).
5. Submit your planning table with your planned mixtures to your instructor for approval prior to proceeding with your experiment.  

Use the following table format to devise your test mixtures:

• Stock [Fe³⁺] = 0.2 M in 0.5 M HNO₃
• Stock [SCN¯] = 2x10^-3 M in 0.5 M HNO₃
• Stock [HNO₃] = 0.5 M

Experiment 2

Perform the lab that you have designed. Absorption λ_max for FeSCN²⁺ is 455nm. Work in groups if specified by your lab instructor.

Part 3: Develop a Plan for Determining Equilibrium Constant, K, of FeSCN²⁺

Complete the Prelab for next week.

Prelab & Results – Week 1

Click here for Prelab worksheet, Week 1: pdf or Word format

There is no Results worksheet for Week 1 (i.e. all of your results will be due after Week 2).

Preparation – Week 2

Questions:

Fill out the prelab worksheet that can be found after the Experiment section for Week 2 (you should have already done this in lab during Week 1). Submit your prelab to your instructor.

Print:

Print and bring to lab Workshop 7 in pdf or Word format. Please print double-sided.

Experiment – Week 2

To print instructions, select the portion that you with to print, choose File/Print, and choose "selection" to prevent printing the entire document.

Safety

• Wear safety glasses
• Pour all waste solutions into the appropriate waste container.

Equipment and Materials

• 0.2 M iron (III) nitrate[Fe(NO₃)₃] 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 x 10^-3 M 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(s) 
• Spectronic 20 with test tube cuvettes
• 1-5 mL volumetric pipets
• 10, 15, & 25 mL volumetric pipets (week 2 only)
• 50 & 100 volumetric flasks
• 10, 25, and 100 mL graduated cylinders

Experimental

Part 4: Determining the Equilibrium Constant, K, of FeSCN²⁺

Workshop 7: Equilibrium

Click here to download Workshop 7: Equilibrium in pdf or Word format

Prelabs & Results – Week 2

Click here for Prelab worksheet, Week 2: pdf or Word format

Click here for Results worksheet, due following Week 2: pdf or Word format

Question for Thought:

1. Discuss the precision of your results based on the solutions and glassware used in lab. [Instead of using AU and RU, please use the number of significant figures for the solutions, glassware, and Spec 20 measurements to determine the precision in each measurement.]  Compare this to your standard deviation.

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