Determination of an Equilibrium Constant
CH 127 – Chem 2 Lab
Determination of an Equilibrium constant
Goals
The purpose of this experiment is to determine the equilibrium constant for the reaction Fe3+(aq) + HSCN(aq) –>FeSCN2+(aq) + H+(aq). The equilibrium constant expression Kc for Reaction is kc=FeSCN2+[H+]Fe3++[HSCN] Procedure
*Preparation of the Beer’s law plot
Prepare five solutions of FeSCN2+(aq) of known concentrations between 1x10-5M and 1x10-4M by diluting various volumes of 4.62x10-4 HSCN.
Calculate the Final concentration FeSCN2+ for Beer’s law Using the C1V1=C2V2 Initial Concentration (M) | Initial Volume(ml) | Final Concentration(M) | Initial Volume(ml) | 0.000462 | 1.08 | 0.00002 | 25 | 0.000462 | 3 | 0.00005544 | 25 | 0.000462 | 6 | 0.00011088 | 25 | 0.000462 | 8 | 0.0001478 | 25 | 0.000462 | 10.8 | 0.0002 | 25 |
Using a spectrophotometer, the absorbance, A, of a solution measured directly. FeSCN2+ is placed into the spectrophotometer and their absorbances at 447nm are measured.
Beer’s law is then used to fine unknown concentration in the experiment from their measured absorbance.
Calibrate the spectrometer by using 0.500M HNO3 as the blank.
*Determination of an Equilibrium constant
Use the larger micropipette with tip to place 1.00 mL of 5.00 x 10-3 M Fe3+ (in 0.500 M HNO3 (aq)) in a clean, dry cuvette. This is not the same Fe3+ solution used in the earlier Beer’s Law portion of the experiment. Using the smaller micropipette with tip, add 0.20 mL of 0.00100 M HSCN (in 0.500 M HNO3 (aq)) to the cuvette. Stir the mixture by carefully drawing up the solution several times using a clean, plastic transfer pipet, being certain that no solution is lost during the mixing. Measure the absorbance of the solution at 447 nm.
Add another 0.20 mL of HSCN (in 0.500 M HNO3 (aq)) to the cuvette, mix, and measure the absorbance of the solution at 447 nm. Repeat the HSCN additions until a total of at least 1 mL of HSCN has been added.
Data
The following data we
Determination of an Equilibrium constant
Goals
The purpose of this experiment is to determine the equilibrium constant for the reaction Fe3+(aq) + HSCN(aq) –>FeSCN2+(aq) + H+(aq). The equilibrium constant expression Kc for Reaction is kc=FeSCN2+[H+]Fe3++[HSCN] Procedure
*Preparation of the Beer’s law plot
Prepare five solutions of FeSCN2+(aq) of known concentrations between 1x10-5M and 1x10-4M by diluting various volumes of 4.62x10-4 HSCN.
Calculate the Final concentration FeSCN2+ for Beer’s law Using the C1V1=C2V2 Initial Concentration (M) | Initial Volume(ml) | Final Concentration(M) | Initial Volume(ml) | 0.000462 | 1.08 | 0.00002 | 25 | 0.000462 | 3 | 0.00005544 | 25 | 0.000462 | 6 | 0.00011088 | 25 | 0.000462 | 8 | 0.0001478 | 25 | 0.000462 | 10.8 | 0.0002 | 25 |
Using a spectrophotometer, the absorbance, A, of a solution measured directly. FeSCN2+ is placed into the spectrophotometer and their absorbances at 447nm are measured.
Beer’s law is then used to fine unknown concentration in the experiment from their measured absorbance.
Calibrate the spectrometer by using 0.500M HNO3 as the blank.
*Determination of an Equilibrium constant
Use the larger micropipette with tip to place 1.00 mL of 5.00 x 10-3 M Fe3+ (in 0.500 M HNO3 (aq)) in a clean, dry cuvette. This is not the same Fe3+ solution used in the earlier Beer’s Law portion of the experiment. Using the smaller micropipette with tip, add 0.20 mL of 0.00100 M HSCN (in 0.500 M HNO3 (aq)) to the cuvette. Stir the mixture by carefully drawing up the solution several times using a clean, plastic transfer pipet, being certain that no solution is lost during the mixing. Measure the absorbance of the solution at 447 nm.
Add another 0.20 mL of HSCN (in 0.500 M HNO3 (aq)) to the cuvette, mix, and measure the absorbance of the solution at 447 nm. Repeat the HSCN additions until a total of at least 1 mL of HSCN has been added.
Data
The following data we