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Reaction Rate (Kinetics) of the Iodine Clock Reaction Some chemical reactions such as an explosion occur so rapidly that they cannot be easily studied. Other reactions such as rusting of iron require more time and have a slower measurable rate of reaction. The rate of a chemical reaction between A, B, and C is influenced by concentration of the reactants (or partial pressure if they are gases), the presence of a catalyst, and of course the temperature as predicted by the collision and kinetic molecular theories. The rate can be described by the following relationship: rate = k [A]a[B]b[C]c The quantities in brackets are read as moles/liter of each reactant and are raised to an exponent. Multiplied together with the constant (k), they give the rate of the reaction. The numerical values of a, b, and c must be determined by experimentation. These numbers determine the order of the reaction. Added together they give the over-all order of the reaction. It is the purpose of this experiment to determine the rate and order of an iodine clock reaction. The overall chemical reaction and the rate determining process is: 2 IO3-1 (aq) + 5 HSO3-1 (aq) + 2 H+1 (aq) → I2 (aq) + 5 HSO4-1 (aq) + H2O Iodine, I2, reacts quickly with a bisulfite ion and is converted into colorless iodide: I2 (aq) + HSO3-1 (aq) + H2O → 2 I-1 (aq) + SO4-2 (aq) + H+1 (aq) The second reaction is very fast and as long as any HSO3-1 (aq) remains the triiodide ion can not form the blue complex with starch. The time required, which is inversely related to the rate, for the formation of the blue color is dependent upon the concentration of HSO3-1 (aq). If HSO3-1 (aq) is in excess the blue-black starch-iodide complex will not form. This reaction is being studied with IO3-1 (aq) as the excess reactant and the HSO3- (aq) as the limiting reactant. Materials: 1. 0.0187 M KIO3 (A) 2. 0.016 M NaHSO3 with acidified starch (B) 3. DI water 4. 3 micro-pipets 5. 8 micro-well strips (may be broken into 4-unit strips) 6. 3 small beakers 7. thermometer Procedure: Be careful to not contaminate the solutions in the vials marked A, B, and water. The lids must be put back on the proper vial. Mark the pipets and place them on a piece of marked paper towel. Squeeze the pipet before (not after) you put it into a vial to acquire some liquid. Only use a pipet in the solution it is marked to dispense. You can assume that if you hold the pipet straight up and do not let the drops touch the side of the micro-well strip before they are released that they will all form the same size drop. Use an electronic balance and dispense drops of DI water from each pipet until approximately 1.00 g is reached, comparing the mass of each, calculate the number of drops in 1.00 mL. Without touching the bulb of the thermometer, record the temperature of the room. Share your temperature data so that an average room temperature can be determined. You may assume that each of the solutions are at room temperature. If you have time to investigate this reaction at different temperatures as described in Part III, you should measure the temperature of the water in the water bath just as you remove the micro-pipets from the waterbath and assume that A and B are the same temperature. Good scientific technique is very helpful here so think and be careful to eliminate sources of error. Create a data table in your journal before the experiment starts (if possible) to organize the data and observations you collect.. Part I: Determining the effect of varying concentration on the rate of reaction 1. React eight drops of A with eight drops of B without any dilution with water. This must be performed at least three times and should be completed five times if there is not good agreement in the first three trials. Add eight drops of A and B into separate well strips. Invert one strip over the other and shake or tap them to transfer on liquid into the other. Time immediately until there is a distinct color change. 2. Use C1V1 = C2V2 to calculate the initial concentrations of IO3-1 and HSO3-1 at the instant the solutions are mixed together. If the initial concentration of IO3-1 (aq) in 8 drops of A is 0.0187 M, then what will be the concentration after the addition of 8 drops of another liquid making a total of 16 drops. Do the same calculation for HSO3-1. Assume that the HSO3-1 (aq) ion is completely consumed in each reaction and calculate the rate for each of these reactions. Record your results in Table 1. Rate = - [HSO3-1] = [HSO3-1]f - [HSO3-1]i or Rate = - [IO3-1] = [IO3-1]f - [IO3-1]i t t t t 3. What was the average rate of the reaction when 8 drops of each liquid was reacted? Table 1: Reaction Time and Rate at Initial Concentrations Trial # A KIO3 (drops) B (NaHSO3) [HSO3-1]i Time (s) 1/time (sec-1) Rate (M/s) (drops) 1 8 8 2 8 8 3 8 8 4 5 Part II: Varying concentration by dilution with DI water (16 drops total) Add the following to two strips and combine and react as in part I. Record the time for each reaction. Repeat each for best results. Table 2: Reaction Time and Rate at Various Concentrations Trial # A KIO3 B (NaHSO3) [IO3-1]i [HSO3-1]i Time (s) 1/time (sec-1) Rate (M/s) (drops) (drops) 6 8 6 + 2 (DI) 7 8 5 + 3 (DI) 8 8 4 + 4 (DI) 9 8 3 + 5 (DI) 10 6 + 2 (DI) 8 11 5 + 3 (DI) 8 12 4 + 4 (DI) 8 13 3 + 5 (DI) 8 14 2 + 6 (DI) 8 Points to consider: What effect did the dilution of B have on the rate of the reaction? What effect did the dilution of A have on the rate of the reaction? Part III: Varying temperature without dilution (16 drops total) Your group may be assigned one temperature to check several times and share with the rest of the class. Place some A in a pipet and B in a different pipet. Place one of each into a small beaker that has water of the temperature you were assigned; that is approximately 15oC (tap), 27oC (slightly above room temp), 35oC (max). Allow the liquids in the pipets to warm for 3-5 minutes. Measure the temperature of the water in the beaker just before you dispense them and assume that A and B are that temperature. Adjust by adding hot or cool water. Quickly dispense 8 drops of each into strips that are held in water on a small paper plate that is the proper temperature, quickly invert the strips to mix the solutions, and time the reaction. Repeat at least twice for good results. You may use the room temperature A and B reactions (trials # 1-3) as another temperature. Table 3: Reaction Time and Rate at Different Temperatures Trial # A KIO3 B (NaHSO3) [HSO3-]i Time (s) 1/time (sec-1) Rate (M/s) (drops) (drops) 12 (15 oC) 8 8 Average 13 (27 oC) 8 8 Average 14 (35 oC) 8 8 Average 15 (room temp) 8 8 Average Questions and Calculations: 1. Use the balanced reaction(s) that show the overall stoichiometry for the reaction and notice that the two major reactants, iodate and hydrogen sulfite, do not react in a 1:1 ratio. When HSO 3-1 was consumed iodine could form the I3- ion and combine with the starch. 2. Calculate the initial concentration of each reactant, the IO3-1 and the HSO3–1. The initial concentration is understood to be the concentration of each reactant in the instant after the solutions are mixed, but before any reaction actually takes place. Because the concentration of a solution is determined in part by the total solution volume, when two solutions are mixed, the concentration of each solution decreases because the total solution volume increases. You must include the dilution factor when calculating the change in concentration, C1V1 = C2V2. You need to calculate the change in reactant concentration. The change a reactant concentration is the same in some of the trials since the amount of IO 3-1 and HSO3–1 added is identical in different trials. The reaction rate varies when the reaction time is different for those trials. Plot concentration hydrogensulfite vs 1/time (sec-1) and plot concentration iodate vs 1/time (sec-1). 3. Be sure to calculate the reaction rate for each trial. The reaction rate is being defined as – [HSO3–1]/ t. Thus, in order to determine the reaction rate, you will need to determine the amount of HSO 3–1 that reacted in the measured time period. This reaction involves some complicated processes but it will be assumed that when the solution turns blue–black all of the HSO3–1 had reacted, allowing the accumulation of I2 molecules and I3-1 ions in the solution. The initial concentration is understood to be the concentration of each reactant in the instant after the solutions are mixed, but before any reaction actually takes place. When two solutions are mixed, the concentration of each component decreases because the total solution volume increases. This dilution factor should be considered when calculating the change in concentration. 4. From the reaction rates calculated above, examine how the rate changes as the concentration of the reactants changes. 5. From the data you collected, you should be able to determine the values of m and n in the rate law expression, rate = k[HSO3–1]m[IO3–1]n? Clearly explain your reasoning and/or show your calculations for determining both m and n. 6. What is the order of the reaction with respect to NaHSO 3? What about the order of the reaction with respect to KIO 3? What is the overall order of the reaction? 7. You should graph your data to help determine or verify the order of the reaction as shown in your chemistry text. If a plot of ln [HSO3–1] vs time gives a straight line the reaction is first order in respect to HSO 3–1and the negative slope of the line is related to the rate constant. If the plot of 1/[HSO3–1] vs time gives a straight line the reaction is second order in respect to HSO3–1. These plots can be used to determine the reaction order for both reactants. 8. Calculate the rate constant, k, for each trial. It should be (about) the same for the five reactions where HSO 3–1 is constant, be sure to include the correct units, so average the values that appear to be useful. ***** Something Extra ***** 9. Use the data from the trials at different temperatures to prepare a plot of temperature (oC) vs time (sec), temperature (oC) vs 1/time (s-1), and then ln k vs 1/T. The last graph can be used to estimate the activation energy for the reaction as defined by the Arrhenius equation, k = A e-Ea/RT, and the slope of the line is equivalent to –Ea/R. For any credit on this part, include your graph and your calculations in your lab notebook and/or report.
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