Chemical Kinetics Ziyue Zhu 2015/2/4
Introduction: The purpose of this experiment was to measure the effects of changes in the concentrations of reactants and temperature on the rate of a reaction. The definition of reaction rate is: the amount of reagent (conventionally reported in units of molarity) consumed or produced per unit of time Rate = - Δ[reactant]/Δt = Δ[product]/Δt (1) Effects of concentration on reaction rate can be mathematically summarized by a rate law: Rate = -Δ[A]/Δt = k[a] 1 [B] 1 (2) The rate law often has the form: Rate= k[a] x [B] y [C] z (3) Effects of a temperature Change on the Reaction Rate: K = Ae -Ea*/RT (4) The first step of this experiment was to run three qualitative tests before the measurements began. In the first step, a few drops of H2O2 was added into KI solution and the following reaction would occur: H2O2 + 2I - +2H + I2 + 2H2O (5) The second test was to divide the mixture into three portion and add a few drops of starch solution into the second portion. Finally, thiosulfate solution and a few starch was added into the third portions and the following reaction will occur: 2S2O3 2- + I2 2I - + S4O6 2- (6) Take record of observations after each test was finished. The second step of this experiment was to measure the effects of concentration and temperature on the rate of reactions. The first step was to mix D.I. Water, Buffer, Na2S2O3, KI and starch for A-H measurements with different given volumes. Then poured H2O2 and take record of time from the moments H2O2 was poured to the moment the solution began to change color because the reaction (5) and (6) happened. In that case, the time could be used to calculate the rate of reaction by using equation (2). In addition, measurements E and H would need to record temperature and the temperature of measurements H should always be 5 o C lower than the temperature of measurements E. In that case, it was easy to tell the effects of temperature on the rate of reaction by comparing measurements E with H.
Data and Calculations Table 1: Molar concentrations of reagent solutions: Reagents S2O3 2- I - H2O2 0.019M 0.06M 0.11M Concentrations Table 2: Data summary table Sol n Temp [I - ] Log[I - ] [H2O2] log[h2o2] Δt min Rate: -Δ[H2O2]/Δt log(rate) A 0.021-1.678 5.5x10-3 -2.26 9.3 5.10753x10-5 -4.2918 B 0.015-1.8239 5.5x10-3 -2.26 11.55 4.1126x10-5 -4.2859 C 9x10-3 -2.0458 5.5x10-3 -2.26 18.15 2.617x10-5 -4.5822 D 6x10-3 -2.2218 5.5x10-3 -2.26 21.97 2.162x10-5 -4.665 E 22 o C 6x10-3 -2.2218 0.011-1.9586 14.217 3.341x10-5 -4.476 F 6x10-3 -2.2218 0.0165-1.7825 14.33 3.315x10-5 -4.48 G 6x10-3 -2.2218 0.0275-1.5607 5 9.5x10-5 -4.0223 H 17 o C 6x10-3 -2.2218 0.011-1.9586 18.5 2.568x10-5 -4.59
Calculations: 1. a. Moles S2O3 2 in each solution. 9.5x10-5 mol (a). b. Number of moles of I2 required to exhaust the moles of S2O3 2 from 4.75x10-5 mol c. Number of moles of H2O2 required to produce number of moles of I2 in (b). 4.75x10-5 mol d. Change in H2O2 concentration, Δ [H2O2] The number of moles of H2O2 lost (from c above) IS the negative change in moles of H2O2 ( Δ moles H2O2). You just need to change moles to molarity (what is final solution volume?) 4.75x10-4 M (Note: Answers to a, b, c, and d are identical for all eight solutions.)
2. Reaction Order (attach graphs): Graph 1 : determining r, the order in HxOx: Graph 2: determining s, the order in I -: According to the two graph above, it is easily to tell the value of r and s. r = 0.8374 s = 0.7132
3. Value of rate constant, k Run k A 0.062665 B 0.064143 C 0.05876 D 0.06482 E 0.056058 F 0.0396 G 0.074 Average 0.00600657 The units should be mol -0.5506 x L 0.5506 x min -1 Using the data above, Ea* = 134.38 and units should be J/ (Δmol x o C) All the calculations are on the carbon copy of the calculation page.
Results and discussion: First of all, the results of the qualitative tests were: only the second portion turned to blue color. The reason why the first test did not change color was because that the H2O2 can change the I - into I2. However, I2 did not has much color in water when it was diluted. In the second test, the starch was added into the solution and OH group in starch reacted with I2 and formed some clathrate solution which could only relfect blue light so that the solution showed blue color. In the third test, the thiosulfate solution was added first and the S2O3 2- reacted with I2 and consumed all of I2 so that when the starch was added after the first step, no I2 was left to react with starch. In that case, the solution did not change into blue. The results of r value was 0.8374 and the s value was 0.7132,according to graph 1 and 2.The two graph shows that the log(rate) has a linear relation with log(h2o2) and log(i - ). These means the reaction order in H2O2 and I - was both in 1 st order. The expected value of r and s are both 1 so the error percentage of r was 16.26% and s was 28.68%. The results of k value was 0.00600657 in average. This means that the reaction rate constant was 0.00600657 mol per L per min. Because the expected k value was not found, the error percentage can not be calculated. The value of Ea* was 134.38J per Δmoles per minute. The most possible reason of the errors was that during the F trials, the starch was not added before the time was recorded. In that case, the Δt of trial F was higher than actual Δt. Another possible reason was when doing several trails, the actual compounds added into the solution was not as accurate as the given table. Post-Lab Questions: 1. The reason of the errors caused the actual value of r and s lower than 1 because if the starch was added later, it would let the Δt increase so that the rate of reaction will slow down. In that case, the log(rate) will be lower than expected value and the r and s value would also be lower. 2. According to the data of experiment, the rate of reaction was increased when the concentration of the reactants was increased. The rate of reaction was decreased when the temperature of the reaction was decreased. This two conclusion happened when the k value was determined to be 1 or close to 1.