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In conclusion it was found that the activation energy of the decomposition of hydrogen peroxide with the help of the catalyst KI is . When compared to an actual value there was a percent difference of 43%. In fact the actual value that was used is the activation energy of hydrogen peroxide in the absence of a catalyst. So in reality it is likely that a catalyst would cause the activation energy to be even smaller, and the percent difference would be even greater. The only reasonable explanation as to why the calculated activation energy is so much greater is that somewhere during the experiment a random error occurred but since only two trials were done, it is impossible to pinpoint where it exactly occurred.
When looking at graphs 1-4 it can be seen that as time went on the pressure increased exponentially. This makes sense because over time the amount of gas increased, and so the pressure would increase too. In addition in graphs 1-4, if a curved line of best fit were to be plotted, than the y-intercept would show the pressure of the room at the moment the experiment was conducted.
There were quite a few weaknesses and limitations in this lab. The biggest weakness in this lab was probably the fact that the experiment was conducted at only two different temperatures. This in turn only allowed for two points to be graphed on graph 5 (ln K vs 1/T). Having only two points on a graph is very inaccurate and imprecise. This is because if for some reason a random error occurred during the procedure, than it would be impossible to tell, as there are no other points to compare with. Two trials are not sufficient enough to collect results from; a minimum of five trials should have been done at different temperature increments. Another limitation was the fact that the two different temperatures only had a difference of 10. As seen on graph 5, the slight change in temperature caused the slope to be extremely steep.
A weakness in the lab was the fact that the lab was conducted over a period of two days. This in turn caused the temperatures to be different for part two and part three of the experiment when the temperature should have been kept the same. This is seen in the results, during part two the temperature of the water bath was 21.5 and the initial rate was 0.062kPa/s. But since the rate order of was determined to be one, if the molarity of the were to be halved, and the molarity of the catalyst KI were to be doubled, than ideally the initial rate should have remained the same.
But it didn’t, since part three was conducted at a different day, the temperature of the water bath that day was 22.7 and that slight change in temperature is what caused the initial rate to increase to 0.076kPa/s. Lastly, when the test tube was put in the water bath, the temperature of the water bath was constantly changing because the temperature of the water was relatively lower than the room temperature. This in turn is what caused fluctuation in the temperature readings. This in turn could have also caused the initial rate to vary because as discussed earlier, even the slightest change in temperature causes the initial rate to change.
This experiment can be improved in many ways. One of the main things that can be done is that the experiment should be conducted at different temperatures so that at least a minimum of 5 k values against 1/T can be plotted on graph 5. In addition, the temperature increments should have a relatively broad range, which in turn will make the data and the trends clearer. In addition, the experiment should be conducted on the same day so as the temperature can be kept constant for certain k values. In order to keep the temperature of the water bath constant, the experiment should be conducted in a closed environment so that less energy is lost to the environment.
“The Catalytic Decomposition of Hydrogen Peroxide.” Purdue University College of
Science Welcome. Web. 12 Oct. 2009.