Repeat for the remaining kinetic trials. Mix and time the test solutions for the
ID: 1003566 • Letter: R
Question
Repeat for the remaining kinetic trials. Mix and time the test solutions for the remaining seven kinetic trials. If the instructor approves, conduct additional kinetic trials, either by repeating those in Table 5.1 or by preparing other combinations of KI and H_2 0_2. Make sure that the total diluted volume remains constant at 10 mL. Disposal: Dispose of the solution from the kinetic trails in the waste lodide salts container. CLEANUP: Rinse the beakers or test tubes twice with tap water and discard in the Waste Iodide Salts container. Dispose of two final rinses with deionized water in the sink. Perform the calculations, carefully one step at a time. Appropriate and correctly programmed software would be invaluable for completing this analysis. As you read through this section, complete the appropriate calculation and record it for each test solution on the Report Sheet. Moles of I_3^- produced. Calculate the moles ot S_2 O_3^2- consumed in each kinetic trial. From equation 5.11. the moles of I_3^- that form in the reaction equals one-half the moles of S_2 O_3^2- that react. This also equals the change in the moles of I_3^- starting with none at time zero up until a final amount that was produced at the time of the color change. This is designated as "Delta (mol I_3^-)"produced. Reaction rate. The reaction rate for each kinetic trial is calculated as the ratio of the moles of I_3^- produced. Delta (mol I_3^-), to the time lapse. Delta t, for the appearance of the deep-blue color.^5 Compute these reaction rates Delta (mol I_3^-)/Delta t, and the logarithms of the reaction rates (see equations 5.7and 5.8)for each kinetic trial and enter them on the Report Sheet. Because the total volume is a constant for all kinetic trials, we do not need to calculate the molar concentrations of the I_3^- produced. Initial iodide concentrations. Calculate the initial molar concentration.[I^-]_0 and the logarithm of the initial molar concentration,log.[I^-]_0,of iodide ion for each kinetic trial.^6 Sec Prelaboratory Assignment, question 4d. Initial hydrogen peroxide concentrations. Calculate the initial molar concentration [H_2 O_2]_0 and the logarithm of the initial molar concentration log [H_2 O_2]_0 of hydrogen peroxide for each kinetic .^7 see prelaboratory assignment, question 4e. C. Determination of the Reaction Order, p and q, for Each Reactant Appendix C Appendix C Determination of p from plot of data. Plot on top half of a sheet of linear graph paper (available at the end of this lab manual) using appropriate software log (Deltamol I_3/Delta t), which is log (rate) (y-axis), versus log[I^-]_0(x-axis) at constant hydrogen peroxide concentration. Kinetic trails 1, 2, 3 and 4 have the same H_2O_2 concentration. Draw the best straight line through the four points. Calculate the slope of the straight line. The slope is the order of the reaction p, with respect to the iodide ion. Determination of q from plot of data. Plot on the bottom half of the sheet of linear graph paper or preferably by using appropriate software log Delta mol I_3/Delta t) (y-axis) versus log [H_2O_2]_0 (x-axis) at constant iodide ion using kinetic trials I, 5, 6, and 7. Draw the best straight line through the four points and calculate its slope. The slope of the plot is the order of the reaction with respect to the hydrogen peroxide. Approval of graphs. Have your instructor approve both graphs. Determination of k', the Specific Rate Constant for the Reaction Appendix B Substitution of p and q into rate law. Use the values of p and q (from Part C) and the rate law, rate = Delta(mol I_2)/Delta t = k'[I^-]^p [H_2O_2]^q, to determine k' for the seven solutions. Calculate the average value of k' with proper units. Also determine the standard deviation and relative standard deviation (%RSD) of k'from your data. Class data. Obtain average k' values from other groups in the class. Calculate a standard deviation and relative standard deviation (%RSD) of k' for the class. Determination of Activation Energy Appendix C Prepare test solutions. Refer to Table 5.1, kinetic trial 4. In separate, clean 150-mm test tubes prepare two additional sets of solution A and solution B. Place one (solution A/solution B) set in an ice bath. Place the other set in a warm water backsim35degree C) bath. Allow thermal equilibrium to be established for each set, about 5 minutes. Test solutions prepared at other temperatures are encouraged for additional data points. Mix solutions A and B. When thermal equilibrium has been established, quickly pour solution B into solution A, START TIME, and agitate the mixture. Place it back in the bath. When the deep-blue color appears, STOP TIME. Record the time lapse as before. Record the temperature of the water bath and use this time lapse for your calculations. Repeat to check reproducibility and for the other set(s) of solutions. The reaction rates and "new" rate constants. The procedure for determining the reaction rates is described in .2. Calculate and record the reaction rates for the (at least) two trials (two temperatures) from Part E.2 and re-record the reaction rate for the (room temperature) kinetic trial 4 in .5. Carefully complete the calculations on the Report Sheet. Use the reaction rates at the three temperatures (ice, room and backsim35degree C temparature) and the established rate law from Part C to calculate 'the rate constants k', at these temperatures. Calculate the natural logarithm of these rate constants. Plot the data. Plot In k' versus MT(K) for the (at least) three trials at which the experiment was performed. Remember to express temperature in kelvins and R = 8.314 J/mol middot K Activation energy. From the data plot, determine the slope of the linear nln. (= -E_a/R) and calculate the activation energy for the reaction. You may need to seek the advice of your instructor for completing the calculations on the Report Sheet. The Next Step The rate law for any number of chemical reactions can be studied in the same manner - for example, see Experiment 4, Parts B, C, and F. Research the Internet for a kinetic study of interest (biochemical?) and design a systematic kinetic study of a chemical systemExplanation / Answer
As you have not given all the data (dt in the 1st experiment and 2nd experiment), I am describing the methods you should follow for the post-experiment calculations.
Experiment 1:
Step (a): The reactions are:
3I- + H2O2 + 2H+ ---> I3- + 2H2O (1)
I3- + 2S2O32- ---> 3I- + S4O62- . (2)
From the 2nd eqn, Rate of the reaction : d[I3-]/dt = d[S2O32- ]/2dt
No of moles of S2O32- present in each solution = 0.02 moles * 1 mL/10 mL = 2*10-3 moles
So, No of moles of I3- present in each solution = 2*10-3 moles/2 = 10-3 moles
Calculate d[I3-]/dt = 10-3 moles/dt for each experiment.
Take logarithm of this for each kinetic experiment.
Step (b) :
Calculate initial concentration of [I-]0 for each experiment.
For kinetic trial 1: [I-]0 = 1 mL * 0.3 M /10 ml = 0.03 M
Calculate the same for rest of the trials.
Take log of each.
Step (c):
Calculate initial concentration of [H2O2]0 for each experiment.
For kinetic trial 1: [H2O2]0 = 3 mL * 0.1 M /10 ml = 0.03 M
Calculate the same for rest of the trials.
Take log of each.
Step (d):
Plot d[I3-]/dt vs log[I-]0 for the kinetic trials where [H2O2]0 are constant (1,2,3, 4 here).
The slope of the plot is order(p) with respect to I- ion.
Step(e):
Plot d[I3-]/dt vs log[H2O2]0 for the kinetic trials where [I-]0 are constant (1,5,6,7 here).
The slope of the plot is order(q) with respect to H2O2 ion.
The rate equation is:
d[I3-]/dt = k' [I-]p[H2O2]q.
Exp 2:
The Arrhenius equation is: rate constant, k' = Ae-Ea/RT
or, lnk' = lnA - (Ea/RT)
So if we make a plot of lnk' vs 1/T at different temperature for a particular reaction, we can calculate activation energy(Ea) from the slope(=Ea/R) of the graph.
You can calculate the rates (d[I3-]/dt) by the calculations described in the 1st experiment. Calculate initial concentration of I- and H2O2 in the new set of solutions. You know the value of p and q from the 1st experiment. From the equation d[I3-]/dt = k' [I-]p[H2O2]q, calculate k' for the reactions at different temperature.
Plot k' vs (1/T) and calculate activation energy(Ea) from the slope(=Ea/R) of the graph.