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1212L-Equilibrium Constant for Ferric Thiocyanate-v1

Academic year: 2020/2021
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Determination of the Equilibrium Constant for Ferric Thiocyanate

Please review the background knowledge as: The Equilibrium Constant: chem/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/ Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Equilibria/Chemical_Equilibria/ The_Equilibrium_Constant

Calculating an Equilibrium Concentration: chem/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/ Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Equilibria/Chemical_Equilibria/ Calculating_an_Equilibrium_Concentration

Spectrophotometry: chem/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/ Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Kinetics/Reaction_Rates/Experi mental_Determination_of_Kinetcs/Spectrophotometry

Chemical reactions occur to reach a state of equilibrium. The equilibrium state can be characterized by quantitatively defining its equilibrium constant, Keq. In this experiment, you will determine the value of Keq for the reaction between iron (III) ions and thiocyanate ions, SCN-.

Fe 3 + ( aq ) + SCN −( aq ) → FeSCN 2 +( aq )

The equilibrium constant, Keq , is defined by the equation shown below.

[ ][ ]

[ ]

3

2

  • −

=

Fe SCN

FeSCN

K eq

To find the value of Keq , it is necessary to determine the molar concentration of each of the three species in solution at equilibrium. If we know the initial concentrations of each species, and measure the equilibrium concentration of a single species, we can use the reaction stoichiometry to determine the equilibrium concentrations of all the other species. Please review “ Calculating an Equilibrium Concentration ” on Chemwiki, which will help you carry out the calculation.

Ferric Thiocyanate is a brick-red complex ion. We will use spectrophotometry to determine the molar concentration of Ferric Thiocyanate, FeSCN2+, at the equilibrium. Fe3+(aq) and SCN-(aq) do not interfere with this measurement because they are only weakly colored. Knowing the initial concentrations of Fe3+ and SCN- will allow us to determine their equilibrium concentration via the reaction stoichiometry.

In order to successfully evaluate this equilibrium system, you will: Part I You will prepare a series of standard solutions of FeSCN2+ from solutions of varying concentrations of SCN– and constant concentrations of H+ and Fe3+ that are in stoichiometric excess. The excess of H+ ions will ensure that Fe3+ engages in no side reactions (to form FeOH2+, for example). The excess of Fe3+ ions will make the SCN– ions the limiting reagent, thus all of the SCN– used will form FeSCN2+ ions. The FeSCN2+ complex forms slowly,

taking at least one minute for the color to develop. It is best to take absorbance readings after a specific amount of time, between two and four minutes after preparing the equilibrium mixture. Do not wait much longer than four minutes to take the absorbance readings because the mixture is light sensitive and the FeSCN2+ ions will slowly decompose. Part II You will prepare a new series of solutions that have varied concentrations of the Fe3+ ions and the SCN– ions, with a constant concentration of H+ ions. You will use the standard curve developed in Part I to determine the concentration of FeSCN2+ in this new series of equilibrium system. With the known initial concentration of Fe3+ ions and the SCN– ions, you will calculate the equilibrium concentrations of each species, and calculate the equilibrium constant of the reaction.

OBJECTIVES

In this experiment, you will - Prepare and test standard solutions of FeSCN2+ in equilibrium. - Determine the molar concentrations of the ions present in an equilibrium system. - Determine the value of the equilibrium constant, Keq , for the reaction.

MATERIALS

Computer with Logger Pro 0 M Fe(NO 3 ) 3 solution in 0 M HNO 3 Plastic cuvettes 0 M Fe(NO 3 ) 3 solution in 0 M HNO 3 Pipet Pump and 10 mL pipettes 0 M KSCN solution Beakers Distill water Test tubes on the test tube rack 50 mL volumetric flask Thermometer Kimwipe

  1. Use a USB cable to connect the Spectrometer to the computer. Choose New from the File menu.

  2. Prepare a blank by filling a cuvette 3/4 full with blank solution (Beaker 5). To correctly use cuvettes, remember: a. Wipe the outside of each cuvette with a lint-free tissue. b. Handle cuvettes only by the top edge of the ribbed sides. c. Dislodge any bubbles by gently tapping the cuvette on a hard surface. d. Always position the cuvette so the light passes through the clear sides.

  3. To calibrate the Spectrophotometer, place the blank cuvette into the cuvette slot of the Spectrometer, choose Calibrate → Spectrometer from the Experiment menu. The calibration dialog box will display the message: “Waiting 90 seconds for lamp to warm up.” After 90 seconds, the message will change to “Warm up complete.” Click.

  4. Determine the optimum wavelength for the equilibrium mixture and set up the mode of data collection. a. Empty the blank solution from the blank cuvette. Using the solution in Beaker 1, rinse the cuvette twice with ~1 mL amounts and then fill it 3/4 full. Wipe the outside with a tissue, place it in the Spectrometer. b. Click. The absorbance vs. wavelength spectrum will be displayed. Note that one area of the graph contains a peak absorbance. Click. c. To save your graph of absorbance vs. wavelength, select Store Latest Run from the Experiment menu. d. Click the Configure Spectrometer Data Collection icon, , on the toolbar. A dialog box will appear. e. Select Absorbance vs. Concentration under Set Collection Mode. The wavelength of maximum absorbance (λ max) is automatically identified. Click.

  5. Collect absorbance-concentration data for the five standard equilibrium mixtures. a. Leave the cuvette, containing the Beaker 1 mixture, in the Spectrometer. b. Click. Click , type the concentration of FeSCN2+ (You calculated in Step 2) in the edit box, and click. c. Discard the cuvette contents as directed. Rinse and fill the cuvette with the solution in Beaker 2 and place it in the device. After the reading stabilizes, click , type the concentration of FeSCN2+ (from your pre-lab calculations) in the edit box, and click. d. Repeat Step b to measure the absorbance of the solutions in Beakers 3, 4, and 5. e. Click when you have finished collecting data. Click the Examine button , and write down the absorbance values for each data pair in your notebook.

  6. Click the Linear Fit button,. A best-fit linear regression equation will be plotted for your data. This line should pass near or through the data points and the origin of the graph. Leave the equation in place on the graph and proceed to Step 9. You can use “Print Screen” (Click “Fn” button on the keyboard and “F9” button) function to capture the graph. Paste you screen capture in a Powerpoint file. You can transfer the

Powerpoint file using a USB or email the file to you. You can also record your data and use Microsoft Excel to plot your graph, and get the trendline equation. This is recommended.

Part II Prepare and Test Equilibrium Systems

Test Tube Number

0

Fe(NO 3 ) 3 (mL)

0 M

KSCN (mL)

H 2 O

(mL)

Initial [Fe3+] Initial [SCN-]

1 3 2 5.

2 3 3 4.

3 3 4 3.

Test Tube Number 1 2 3

Absorbance

Initial [Fe3+]

Initial [SCN-]

[FeSCN2+] at equilibrium

[Fe3+] at equilibrium

[SCN-]at equilibrium

Keq

Average Keq

  1. Use the absorbance values, along with the best fit line equation of the standard solutions in Part I to determine the [FeSCN2+] at equilibrium for each of the mixtures that you prepared in Part II. Give an example of your calculations.
  2. Calculate the equilibrium concentrations for Fe3+ and SCN– for the mixtures in test tubes 1-3 in Part II. Give an example of your calculations.
  3. Calculate the value of Keq for the reaction. Explain how you used the data to calculate Keq.
  4. The theoretical value for the equilibrium constant is 138. Calculate the percent error, and discuss the sources of errors.
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1212L-Equilibrium Constant for Ferric Thiocyanate-v1

Course: Foundations Of General Chem Ii (CHEM 7020)

University: Georgia State University

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1
Determination of the Equilibrium Constant for Ferric Thiocyanate
Please review the background knowledge as:
The Equilibrium Constant:
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/
Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Equilibria/Chemical_Equilibria/
The_Equilibrium_Constant
Calculating an Equilibrium Concentration:
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/
Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Equilibria/Chemical_Equilibria/
Calculating_an_Equilibrium_Concentration
Spectrophotometry:
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/
Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Kinetics/Reaction_Rates/Experi
mental_Determination_of_Kinetcs/Spectrophotometry
Chemical reactions occur to reach a state of equilibrium. The equilibrium state can be
characterized
by quantitatively defining its equilibrium constant, Keq. In this experiment, you will
determine the
value of Keq for the reaction between iron (III) ions and thiocyanate ions, SCN-.
)()()(
23
aqFeSCNaqSCNaqFe
+−+
→+
The equilibrium constant, Keq, is defined by the equation shown below.
]][[
][
3
2
−+
+
=
SCNFe
FeSCN
Keq
To find the value of Keq, it is necessary to determine the
molar concentration of each of the three
species in solution at equilibrium. If we know the initial concentrations of each species, and measure
the equilibrium concentration of a single species, we can use the reaction stoichiometry to determine
the equilibrium concentrations of all the other species. Please review “Calculating an Equilibrium
Concentration” on Chemwiki, which will help you carry out the calculation.
Ferric Thiocyanate is a brick-red complex ion. We will use spectrophotometry to determine the molar
concentration of Ferric Thiocyanate, FeSCN2+, at the equilibrium. Fe3+(aq) and SCN-(aq) do not
interfere with this measurement because they are only weakly colored. Knowing the initial
concentrations of Fe3+ and SCN- will allow us to determine their equilibrium concentration via the
reaction stoichiometry.
In order to successfully evaluate this equilibrium system, you will:
Part I You will prepare a series of standard solutions of FeSCN2+ from solutions of varying
concentrations of SCN– and constant concentrations of H+ and Fe3+ that are in stoichiometric
excess. The excess of H+ ions will ensure that Fe3+ engages in no side reactions (to form
FeOH2+, for example). The excess of Fe3+ ions will make the SCN– ions the limiting reagent,
thus all of the SCN– used will form FeSCN2+ ions. The FeSCN2+ complex forms slowly,

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