1 Purpose: To measure reaction rate at different temperatures for the reaction between persulfate ions, S2O8-2, and iodide ions, I -, and thereby determine the activation energy and frequency factor for the reaction Background The reaction you will study in this experiment is the same one you examined last week for the purpose of getting the rate law for the reaction, including reaction orders and rate constant. Once again, the primary reaction is S2O8-2 (aq) + 3 I - (aq) 2SO4-2 (aq) + I3 - (aq) (eq. 1) You may recall that sodium thiosulfate is added to prevent the triiodide ion from reacting immediately with starch to form the deep blue complex, and that the thiosulfate thereby acts as a chemical stopwatch, and allows determination of the reaction time as the time required to consume a constant amount of thiosulfate and turn the solution blue. The secondary reaction is thus 2S2O3-2 (aq) + I3 - (aq) S4O6-2 (aq) + 3 I - (aq) (eq. 2) In this experiment, you will carry out the reaction in exactly the same way you did in the last experiment, only you will do it at three temperatures other than room temperature. Measuring the reaction rate at different temperatures allows determination of the Arrhenius parameters Ea, activation energy, and A, the frequency factor. Theory You might want to first review the theory behind calculation of reaction rates and orders of reaction that you used in the last experiment, as you will need to do the same kinds of calculations in this lab. Refer to the experiment entitled Determining the Rate Law for a Chemical Reaction. The Arrhenius equation is -E a RT k = A e (eq. 3) where k is the rate constant, A is called the frequency factor, Ea is the activation energy (in J/mole), R is the gas constant 8.314 J/K-mole, and T is the Kelvin temperature. Taking the natural logarithm of both sides of eq. 3 yields
2 E a ln k = ln A - RT (eq. 4) Rearranging slightly on the right side of the equation gives; ln k = - E a RT + ln A (eq. 5) Then factoring the constants out of the first term on the right; ln k = - Ea 1 R T + ln A (eq. 6) which is of the form y = m x + b which, of course, you recognize as the equation for a straight line. So, if rate constants, k, are determined at various temperatures, T, the data can be plotted as 1/T on the x axis, and ln k on the y axis, with the resulting graph being a straight line whose slope is equal to Ea/R and whose y-intercept is equal to the natural logarithm of the frequency factor, ln A. Values of Ea and A can thus be easily determined from the graph. In this experiment, you will determine rate constants for the reaction described by eq. 1 at temperatures about 10 o above room temperature (i.e. about 35 o C), 20 o above room temperature (i.e. about 45 o C), and about 10 o below room temperature (i.e. about 15 o C). These rate constants, along with the room temperature rate constant you determined in the last experiment, will be plotted as described above in order to determine values for the activation energy, Ea, and frequency factor, A, for the reaction. Procedure SAFETY NOTE!! The reagents used in this experiment are irritants. Avoid contact with your skin, eyes, and clothing. If you do spill some on yourself, wash immediately with cold water. Notify your instructor. You will do only three determinations in this experiment, and they will use exactly the same reagents and the same volumes of those reagents as determination 2 in the last experiment, Determining the Rate Law for a Chemical Reaction. The three determinations in this lab will be labeled as determinations 5, 6, and 7 so as not to confuse them with the determinations carried out in the last lab. Table 1 gives the
3 reagents and volumes to be used. For comparison, determination 2 has been included in this table as a shadow. You do not have to repeat determination 2! You will use the data you collected in the previous experiment. determination number (temperature) Table 1 - Determining the Effect of Temperature on Reaction Rate FLASK A FLASK B starch, 1.2x10-2 M 0.20 M 0.20 M 0.20 M ml Na2S2O3, KI, ml KNO3, ml (NH4)2S2O8, (use 1 ml ml ml (use 2 ml pipet) pipet) 0.20 M (NH4)2SO4, ml 2 (room T) 1.00 2.00 4.00 4.00 2.00 6.00 5 (15 o C) 1.00 2.00 4.00 4.00 2.00 6.00 6 (35 o C) 1.00 2.00 4.00 4.00 2.00 6.00 7 (45 o C) 1.00 2.00 4.00 4.00 2.00 6.00 1. Thermostatically controlled constant temperature baths will be set up for use in obtaining the temperatures about 15 o C, 35 o C and 45 o C (determinations 5, 6 and 7, respectively). These baths will automatically maintain a constant temperature. For all three baths, the temperature does not have to be exactly 15 o or 35 o or 45 o, but you will need to record in Data Table 3 exactly what the temperature is as measured by the thermometer immersed in the bath. 2. You will follow exactly the same procedure you did in the last experiment, Determining the Rate Law for a Chemical Reaction, to carry out the three determinations. The only difference in this experiment is that you must put flasks A and B separately into the constant temperature bath, and allow both flasks and their contents to thermally equilibrate, that is, come to the same temperature as the bath, before mixing the contents of the flasks. There should be small (150 ml) beakers about half full of water, sitting on the bottom of the constant temperature baths. The purpose of these beakers is to allow you to set your 50 ml flasks A and B in the beakers in the bath without the flasks tipping over and spilling their contents in the bath water. Label your flasks with your initials, and allow the flasks to sit in the bath for at least 5-7 minutes to equilibrate before mixing. Do not walk away from the water bath while the equilibration occurs. Either you or your lab partner should remain there the entire time the equilibration is taking place. 3. Once the flasks have equilibrated, follow the same procedure as in the last experiment to mix the contents back and forth three times, as one person begins the timer. As soon as the contents have been poured back and forth several times, the flask containing the mixture must be returned to the constant temperature bath to remain at that temperature while the reaction proceeds. Stop the timer when the reaction mixture turns deep blue. Record your reaction times in Data Table 1. Also record the bath temperature, in o C, in Data Table 3.
4 NOTE: The reaction at 15 o C will occur very slowly be patient! 4. Repeat each determination at least one more time until you have reaction times for each determination that agree to within 10% of each other. Make sure you have recorded all reaction times in Data Table 1. 5. Before cleaning your burets, show your data to your instructor to confirm that the results are satisfactory. If yours is not the last section of the lab this week, cover the tops of the burets with parafilm and leave the solutions in them for the next class. If yours is the last section of the lab this week, rinse your burets thoroughly (including the tip) with lots of water and store them in the buret stands upside down with the stopcocks open. All waste from this experiment can be poured down the drain along with lots of water. Calculations 1. Calculate the initial concentrations of iodide ion, persulfate ion, the average reaction time, and reaction rates just as you did in the experiment Determining the Rate Law for a Chemical Reaction under Calculations steps 1, 2, 3, and 4. Record your results in Data Table 2. 2. Using eq. 17 from the previous experiment, calculate values of the rate constant k for each of the three determinations. Use the reaction orders you determined in last week s lab. Record your results in Data Table 3. 3. Convert the Celsius temperatures at which you carried out the reaction into Kelvins, and record the results in Data Table 3. 4. Using the Microsoft Excel spreadsheet in the Science Learning Center, make a plot of ln k on the y-axis versus 1/T on the x-axis. Go to the Science Learning Center website at http://www.montgomerycollege.edu/departments/scilcgt/, click on Online Study Material, then choose CHEM132, and then click on Excel Graphs for Lab. Make sure the Excel window is maximized, and choose the appropriate tab at the bottom of the window. 5. Use the linear regression analysis on the graph to obtain the slope (m) and the y- intercept (b) of your plot. Use those values to calculate Ea and A. Remember, the slope = -Ea/R, and the y-intercept = ln A. Record the results on Data Table 4. 6. Include a copy of the printout of your graph and data table with your lab report.
5 Data Table 1 Reaction Times (s) determination trial 1 trial 2 trial 3 5 6 7 Data Table 2 determination initial [S2O8-2 ] initial [I - ] 5 6 7 average time (s) rate (include correct units) Data Table 3 determination temperature, o C temperature, K k (include correct units) 2 5 6 7 Data Table 4 Activation energy, Ea Frequency Factor, A
6 Post-lab Question 1. Using your values of the activation energy, Ea, and frequency factor, A, that you calculated in today s experiment, calculate a) the rate constant and b) the reaction rate for the reaction S2O8-2 (aq) + 3 I - (aq) 2SO4-2 (aq) + I3 - (aq) at a temperature of 75 o C. Show all calculations.
7 Pre-laboratory Assignment 1. Do Calculation 1 on page 4 for [I - ] and [S2O8-2 ] initial concentrations, and record the results in the appropriate spaces in Data Table 2. (Refer to Table 1 on page 3 for concentrations and volumes.) 2. Draw an energy diagram for an endothermic reaction, and label both axes, the activation energy, and ΔH for the reaction. (You may wish to consult your text.) 3. A reaction is first order in reactant A and second order in reactant B. Starting with [A] = 0.175 M and [B] = 0.00250 M, it is found that the reaction rate is 2.83 x 10-4 Ms -1. What is the value of the rate constant, k? Include correct units. Last revised 8/21/2017 DN