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Experiment 22: Colorimetric determination of an equilibrium constant

Experiment 22: Colorimetric determination of an equilibrium constant. PURPOSE To determine the value of the equilibrium constant for the equilibrium system involving Fe 3+ ( aq ), SCN – ( aq ) and FeSCN 2+ ( aq ) using colorimetric analysis. THEORY.

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Experiment 22: Colorimetric determination of an equilibrium constant

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  1. Experiment 22: Colorimetric determinationof an equilibrium constant PURPOSE To determine the value of the equilibrium constant for the equilibrium system involving Fe3+(aq), SCN–(aq) and FeSCN2+(aq) using colorimetric analysis.

  2. THEORY • In aqueous solution, Fe3+ ions react with SCN– ions to form the blood red coloured FeSCN2+ ion: • Fe3+(aq) (pale yellow) + SCN–(aq) (colourless)  FeSCN2+(aq) (red) • The colour of the solution is directly proportional to the concentration of FeSCN2+ ions present. The concentration of these ions can be determined by measuring the absorbance of the solution and comparing it with the absorbance of a solution of known concentration. If the initial concentrations of Fe3+ and SCN– are known, a value for the equilibrium constant, K, for the reaction can be calculated. • Colorimeters measure the amount of light that is transmitted or absorbed by a solution. A description of how they work can be found in your text book

  3. SAFETY PROCEDURES • Follow all instructions for using the equipment in this activity. • 2. Wear safety glasses and a laboratory coat for this experiment. • 3. Potassium thiocyanate is irritating to the skin and eyes. Avoid contact. • 4. The iron(III) nitrate solution used in this experiment contains nitric acid. The solution is irritating to skin and body tissues.

  4. Part 1. Preparation of Equilibrium Solutions • Label six 50.00 ml flasks 1-6 • Pipet 10.00 ml of .2M Fe(NO3)3 into each flask • Then pipet 1.00, 2.00, 3.00, 4.00, and 5.00 ml of .002M (NaSCN)- into flasks 2 – 6 respectively • Add .1M nitric acid to each flask to make a total of 50.00 ml per flask and stopper each flask

  5. Part 1. Preparation of Equilibrium Solutions • Calculate the final Fe(NO3)3 concentration and report them in part A • Calculate the final Fe(NCS)+2 concentration in each flask

  6. Part 1. Preparation of Equilibrium Solutions • Obtain 2 cuvettes, rinse one cuvette with soln 1, and discard in waste jar • Fill the rinsed cuvette with soln 1, insert cuvette into spectrophotometer and adjust 10 0 absorbance and 100 transmittance. • Rinse the second cuvette with soln 2 and discard in waste jar and measure the absorbance and transmittance of soln 2 at 447nm • Repeat procedures for soln 3 -6

  7. calibration curve • Prepare calibration curve by plotting absorbance vs concentration

  8. Part B: Determination of the calibration curve • Label six clean dry test tubes 1- 6 • Pipet 5.00 ml of .002M Fe(NO3)3 into each test tube • Then pipet 1.00, 2.00, 3.00, 4.00, and 5.00 ml of .002M (NaSCN)- into flasks 2 – 6 respectively • Add .1M nitric acid to each flask to make a total of 10.00 ml per flask and stopper each flask

  9. Part B: Determination of the calibration curve • Measure and record the absorbance and transmittance of these solns at 447nm • From your calibration curve determine the equilibrium concentration of Fe(NCS)+2

  10. Part B: Determination of the calibration curve

  11. Part B: Determination of the calibration curve

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