Table 1: Grid indicating the solutions to add for each initial rate experiment Initial rate experiment 1 Initial rate experiment 2 Initial rate experiment 3 10.0 mL KI(aq 5.0 mL KI(aÄ ) Add first 5.0 mL NapsOsta 5.0 mL NazSO4(ag)5.0 mL NaClaÄ ) 1.0 mL starch(aq)︶ | 1.0 mL starch(aq) 1.0 mL starch(aq) Add last, 10.0 mL K2S,Os(ag) start timer Part A: Determination of rate law Collect the following solutions in separate small, clean, and dry beakers (label them!): 15-20 mL of 0.10 M Na2SO 15-20 mL of 0.10 M NaCl 15-20 mL of starch solution 130 mL of 0.10 M KI 130 mL of 0.10 M K2S20 Obtain three large (25 x 200 mm) test tubes, and make sure that they are clean and as dry as possible. You will also need a digital thermometer and a timer You will perform three initial rates experiments. Table 1 provides the final contents for each of three initial rates experiments. You will set up and run three trials for each initial rates experiment in the three test tubes. For each initial rates experiment: ã£Add the amount of KI solution indicated in Table 1 using a digital pipette to the three test tubes. Rinse the pipette tip with distilled water. b. Where required, add the indicated amounts of NaCI and Na2SOs solutions using the c. Add 1.0 mL of starch solution using the digital pipette to each one of the three test d. Add 5.0 mL Na2S Os solution using the pump dispenser to each one of the three test digital pipette to the three test tubes. Rinse the pipette tip with distilled water betweer solutions. tubes. Rinse the pipette tip with distilled water tubes. Do the following one test tube at a time for all three test tubes: Start the timer as you deliver the indicated amount of K S,Os solution using the digital pipette. Stir the solution at a constant rate using a digital thermometer -consistency in your rate of tirring for all initial rates experiments is important! Stop the timer at the first sign that the solution has turned blue-black, and record time and the temperature at this point. Clean and dry as well as possible the test tubes after each set on initial rates experiments 1sk your TA to check your data for Part A before starting Part B 0ä¸
3. You have been told that the reaction you are observing is:
Fe3+ (aq) + SCN-â (aq) â FeSCN2+ (aq)
Another possibility is that the reaction is:
Fe3+ (aq) + 2 SCN-â (aq) â Fe(SCN)2+ (aq)
Write the equilibrium equation for this new reaction.
Using your data from the mixtures in test tubes 1, 3, and 5 of Part 2, recalculate three new equilibrium constants for this new possible reaction. Hint: Notice that the change in concentration of SCN- in your equilibrium table will change because of the new stoichiometry. Also, the exponent on [SCN-] in the equilibrium constant expression will change for the same reason.
What are the three new equilibrium constants?
Why do the results of your calculations show that the reaction you are observing is more likely to be the first one (one mole of SCN- with one mole of Fe3+)?
4. In Part 1 of this experiment, you made the assumption that the FeSCN2+concentrations in your standard solutions were equal to the initial concentrations of SCN- ions. Now that you know the value of the equilibrium constant, calculate the actual equilibrium concentrations of Fe3+, SCN-, and FeSCN2+ in mixture in test tube 2. How good was the assumption? Did this satisfy the 5% rule (the loss of reactant is less than 5% of its initial concentration)?