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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 reductaseIndian 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.

smiley graphic 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


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

In HNO3: Fe3+ (aq) + SCN- (aq) arrows FeSCN2+ (aq) (1)

for which the equilibrium constant expression is

equilibrium constant expression (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.



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) arrows [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) arrows [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, Part 2 (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.

Part 1: Creation of a Beer’s Law Relationship for FeSCN2+

Reading Assignment:

  • Theory of experiment—see previous pages.
  • Equilibrium and Le Chatelier’s Principle—Zumdahl, Chapter 6.

1. Define the following terms in your own words:

  • Le Chatelier's Principle
  • Beer's Law (see Vitamin Pill lab)
  • limiting reagent


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:

  • [Fe3+] : [SCN- ] must be at least 650:1. (Blank not included.)
  • [SCN- ] must be between 4 x 10-5M and 2 x 10-4M! [These molarities were determined as the minimum and maximum concentrations detectable by the Spec 20.]
  • Never use two volumetric pipettes to generate a single volume of a solution.  For example, never use volumetric pipets of 10-mL and 8-mL to get 18 mL of Fe3+. The error is too great in this volume.
  • If you need to use volumetric pipettes, use volumes that can be delivered by the pipets you have available.  Always fill to the final volume of the volumetric flask with the last reagent.
  • Use 0.5M HNO3 instead of water when creating your dilutions.  This is to keep the concentration of ions in solution constant for every dilution/reaction. 


Equipment and Materials

  • 0.2 M iron (III) nitrate[Fe(NO3)3] in 0.5 M HNO3
  • 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)
  • Spectronic 20 with test tube cuvettes
  • 0.5M nitric acid (HNO3)
  • 1-8, 10, 25 & 50 mL volumetric pipets
  • 25, 50, 100 & 200 mL volumetric flasks


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)
Blank     0 0   0.5M 50.00  
1.       1.2 x 10-4   0.5M 50.00  
2.           0.5M 50.00  
3.           0.5M 50.00  
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.


Part 2: Determination of the Equilibrium Constant, K of FeSCN2+

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!
  • Try to create a concentration of FeSCN2+ at equilibrium ( X ) that will give an absorbance between 0.2 and 1.0 on the Spec 20. Check out your standard curve from Part 1!
  • The K for this reaction is very roughly 102.  You will need to use this value to approximate the equilibrium concentrations of Fe3+ and SCN-
  • Use the following data table to help you determine the initial diluted solution Fe3+ and SCN- that you will need to use to generate your approximate equilibrium concentration of FeSCN2+ determined in “Hint a.” 
  • Once you determine the molarities you will need to create your first reaction from the data table below, modify them slightly to generate your two other reactions.  The reactants don’t need to be exactly the same concentration, just not in a 650:1 ratio!
  • KEY: It will be very helpful if you create a diluted solution of Fe3+ so that you have your own ‘stock solution’ of Fe3+ that is similar in concentration to the SCN- stock solution.
  • Create a data table, similar to Table 1, to enter your volumes and concentrations used for each reactant (Fe3+ , SCN-, & HNO3).  Show this table to your instructor before proceeding.
  Fe3+ + SCN- arrows FeSCN2+

Initial diluted concentration of reactants




Change(X) due to reaction going to equilibrium

- X

- X

+ X

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).


Equipment and Materials

  • 0.2 M iron (III) nitrate[Fe(NO3)3] in 0.5 M HNO3
  • 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)
  • Spectronic 20 with test tube cuvettes
  • 0.5M nitric acid (HNO3)
  • 1-8, 10, 25 & 50 mL volumetric pipets
  • 25, 50, 100 & 200 mL volumetric flasks



Click here for printable instructions

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+ .

Part 1—Questions:
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.

  1. Calculate the equilibrium constant.  Obtain at least three estimates of K by repeating your experiment at least twice with different initial concentrations of Fe3+, SCN-, or both.
  2. Give the mean value and standard deviation for K from your data.  Report your final answer to the correct number of significant figures.
  3. 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.] 


Place all solutions in the waste beakers or bottles in the hood.


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Created By: Adilia James '07 and Sarah Coutlee '07
Maintained By: Nick Doe
Date Created: July 3, 2006
Last Modified: June 6, 2007
Expiration Date: July 3, 2007