Dr. Charles J. Horn
Abstract: This two part experiment is designed to determine the rate law of the following reaction, 2I-(aq) + H2O2(aq) + 2H+I2(aq) + 2H2O(L), and to then determine if a change in temperature has an effect on that rate of this reaction. It was found that the reaction rate=k[I-]^1[H2O2+]^1, and the experimental activation energy is 60.62 KJ/mol.
The rate of a chemical reaction often depends on reactant concentrations, temperature, and if there’s presence of a catalyst. The rate of reaction for this experiment can be determined by analyzing the amount of iodine (I2) formed. Two chemical reactions are useful to determining the amount of iodine is produced. 1) I2(aq) + 2S2O32-(aq) 2I-(aq)+S4O62-(aq)
2) I2(aq) + starch
Reaction 2 is used only to determine when the production of iodine is occurring by turning a clear colorless solution to a blue color. Without this reaction it would be very difficult to determine how much iodine is being produced, due to how quickly thiosulfate and iodine react. However this reaction does not determine the amount of iodine produced, it only determines when/if iodine is present in solution. Reaction 1 is used to determine how much iodine is produced. To understand how the rate constant (k) is temperature dependent, another set of data is recorded in week two’s experiment using six trials and three different temperatures(two trials per temperature change). Using the graph of this data we determine the energy required to bend of stretch the reactant molecules to the point where bonds can break or form, and then assemble products (Activation Energy, Ea). Methods
To perform the experiment for week 1, we first prepare two solutions, A and B, as shown in the data. After preparing the mixtures, we mix them together in a flask and carefully observe the solution, while timing, to see how long it takes for the solution to change from clear to blue. We use this method for all 5 trials, and record the time it takes to change color, indicating the reaction has taken place fully. This data is used to find p (trials1-3) and q (trials3-5), to use in our rate law. This experiment concluded that both p and q are first order. The rate constant average of all five trials is used as just one point on the Arrhenius Plot. In week 2, we perform the experiment to test the relation of temperature to the rate of reaction. We start by again, preparing six solutions. We prepared two trials/solutions at 0 degrees Celsius, two and 40 degrees Celsius, and two at 30 degrees Celsius. Again, for each trial we mixed solution A with B, and carefully timed the reaction to look for a color change that indicates the reaction is complete. The interpretation of this data indicated out results of whether temperature has an effect on the rate of this reaction.
Results- It is determined that the rate of reaction is dependent on the temperature in which the reaction occurs. The solutions observed at 40 degrees Celsius reacted at a quicker rate, than those at lesser temperatures, in a linear manor.
Data Week 1
Table 1: Solution Concentrations Week 1- Room Temperature
trial #| solution A| | | | | Solution B| | | | | | buffer| 0.3MKI| starch| 0.02MNa2S2O3| Distilled water| 0.1MH2O2| time(s)| total volume(mL)| | 1| 5.01| 2.0| 0.4| 5.0| 21.68| 6.0| 585| 40.01| | | 2| 5.0| 4.0| 0.4| 5.0| 19.60| 6.0| 287| 40.00| | | 3| 5.02| 6.0| 0.4| 5.0| 17.60| 6.0| 131| 40.02| | | 4| 5.0| 6.0| 0.4| 5.0| 13.62| 10.0| 114| 40.02| | | 5| 5.0| 6.02| 0.4| 5.0| 9.60| 14.0| 80| 40.02| | |
Calculations Week 1
1. Find the moles of S2O3-2
Take the value from NaS2O3 *(0.2)/1000
(5)*(0.2)/1000= 0.001 mol of S2O32-
2. Find moles of I2
Take S2O32- /2
3. Find I2
Mol I2*1000/vol mL
(0.0005)*1000/40)= 0.000799885 mol