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Determination of Rate Law Essay

Abstract: A solution containing potassium iodide was mixed with a hydrogen peroxide solution to determine the rate law for the reaction between hydrogen peroxide and potassium iodide at room temperature and the activation energy, Ea, of the above reaction. Potassium iodide and hydrogen peroxide react according to the following equation: We found the experimental rate law for this reaction to be and the activation energy of the reaction was calculated to be and the Arrhenius pre-exponential factor (A) of . Introduction All chemical reactions require some minimum amount of energy to transform the reactants into products.

The rate of the reaction is the rate at which the products are formed from reactants. At constant temperature, the rate of a chemical reaction is constant and can be determined experimentally using the general rate law . Varying the temperature at which a particular reaction takes place changes the activation energy of the reaction at the different temperatures. Using a transformed version of the Arrhenius equation the value of for the reaction and the Arrhenius pre-exponential factor can be determined graphically. Methods There were two solutions involved in this experiment: Solution A comprised of 5.

OmL buffer (to stabilize [H+]), 0. 3M KI (a source of I-), starch (indicator for I2), 0. 02M sodium thiosulfate (source of thiosulfate ion), and distilled water (to bring the total volume to 40. 00mL), while solution B contained 0. 1M hydrogen peroxide. In the first part of the experiment, we determined the rate law as follows: We prepared solutions A and B for each trial using the recommended volumes in Table 2 of the lab manual.

After preparing the solutions, we used separate thermometers to record the temperature of each solution to the nearest 0. 1, ensuring that both solution temperatures did not deviate by more than 0. 5. The data obtained was recorded as Table 1. After recording the temperatures, my partner started the timer on her phone while I poured solution B into the flask containing solution A. The end of the reaction was signaled by the formation of a blue iodine-starch complex in the flask. The amount of iodine produced was calculated using the amount of thiosulfate (limiting reagent in the thiosulfate-iodine reaction) in the solution.

After performing all five trials, the values obtained for the first three trials were used to create Table 1a below. These values were then plotted using Graphical Analysis and curve fitted to determine the order of the reaction with respect to iodide as shown on Figure 1a. Table 1b was also created using the values for the last three trials, then plotted on a graph (as shown on Figure 1b) to determine the order of the reaction with respect to hydrogen peroxide and two values for the rate constant, .

The values for p and q were rounded to the nearest integer and the average of the two values was then calculated to be resulting in the rate law for the of In the second part of the experiment, we determined the activation energy for the decomposition of hydrogen peroxide using potassium iodide by performing runs similar to part 1, but varying temperatures at which the reaction takes place. For each run, solution A and B were prepared using the recommended values from the lab manual.

We then place both solutions in an ice bath (for the first 2 runs) and in a water bath (for the remaining runs) to get their temperatures to the same values as that of the water/ice in the bath. We also used the temperature values suggested in the lab manual. When needed, we increased the temperature by heating the hot plate on which we placed our bath of adding ice cubes into the bath. Once the thermometers in each solution and that in the bath reached the desired value, I simultaneously noted the time on the lab clock and poured solution B into the flask containing solution B.

I recorded the time from when I poured solution A into B to when I noticed a color change from colorless to purple. I then used the data obtained to plot a graph of ln(k) against the reciprocal of the temperature for all the six runs, plus the average value of k and temperature calculated from the first part of the lab. This graph was then used to determine the activation energy, Ea and the Arrhenius pre-exponential factor, A. We report an A value of and an activation energy value of 56. 80kJ/mol. This compares to theoretical value of 56. 5kJ/mol at 0. 53% difference.


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