Kinetics of the Decomposition of Hydrogen Peroxide
Autor: Mikki • June 3, 2018 • 851 Words (4 Pages) • 744 Views
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Experiment 3: Trial 2
Volume (mL)
Time (sec)
Total Time (sec)
2.0
0
0
4.0
39.71
39.71
6.0
44.96
84.67
8.0
43.50
128.17
10.0
43.98
172.15
12.0
43.41
215.56
14.0
37.90
253.46
Table 2: Rate of Reactions and Results
Trial 1
Trial 2
Experiment 1
0.0298
0.0306
Experiment 2
0.0594
0.0510
Experiment 3
0.0506
0.0467
Determining m [H2O2]
1.00
0.740
Determining n [KI]
0.764
0.614
Calculations:
Experiment 1
1st Trial - Rate of Reaction: 0.0298
2nd Trial - Rate of Reaction: 0.0306
Ave. Rate of Reaction for Experiment 1: 0.0302
Experiment 2
1st Trial - Rate of Reaction: 0.0594
2nd Trial - Rate of Reaction: 0.0510
Ave. Rate of Reaction for Experiment 2: 0.0552
Experiment 3
1st Trial - Rate of Reaction: 0.0506
2nd Trial - Rate of Reaction: 0.0467
Ave. Rate of Reaction for Experiment 3: 0.04865
Determining m: Calculation
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m = 1.00
Determining n: Calculation
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n = 0.764
Complete expression for the Rate of Reaction (Rate Law)
Rate = k[H2O2]1 [KI].764
Discussion:
In this lab, the formation of oxygen in a reaction was measured. The experiment consisted of 0.10 M of KI, 3% of H2O2, and H2O. There was three experiments conducted to determine the effect of concentration of reactants. When conducting the first experiment 10 mL of 0.10 M of KI , 15 mL H2O, and 5 mL of 3% H2O2 were added together in a flask. As the solution was being stirred and timed as the oxygen level increased, the rate of oxygen formation was averaged to be about a minute. The second experiment remained the same but consisted of an increase of H2O2 of 10 mL and decrease of H2O, 10 mL. This experiment measured a more rapid formation of gas. The average time of gas formation in this experiment was 33.5 seconds. Experiment three doubled the amount of concentration of KI to 20 mL and decreased the amount of H2O2 and H2O to 5 mL. When measuring the rate of oxygen the rate was somewhere in between experiment one and two. The average time of gas formation was 40 seconds.
These results showed the affect of concentration of reactants. The amount of concentration determined the rate of oxygen formed. When graphing the results, the linear graphs were consistent with time. The graphs compared volume versus the total time that was measured of gas. The volume of gas increased as time went by.
Post Lab Questions:
- Throughout each experiment the water level in the buret was kept the same as that in the leveling bulb. If the leveling bulb had remained in a fixed position, the effect upon the rate would be incorrect. The reason for this is that, it is essential that the leveling bulb and the buret remain in the same position because the matched water level assures the pressure inside the buret is the same as the outside pressure which is the same as atmospheric pressure.
- Another term in the rate law that affects the rate of reaction that was shown in this experiment was temperature. Because the solutions used were at 18.0°C, it allowed for the rate of the gas formation to be measured at a certain rate as opposed to if it was warmer. If the temperature of the solution was warmer the production of gas would have been more rapid.
- If the rate of H2O formation would have been monitored instead of O2, the rate of reaction for the entire process would have taken a very long time, there would be no sufficient time in class to finish the experiment. H2O is hydrophilic which makes it more difficult to extract from the solutions and difficult to be measured.
- If the concentration of H2O2 is increased three folds from experiment 1 to experiment 2, this would affect the rate of the reaction of experiment 2 with respect to that of experiment 1, the O2 formation would increase rapidly. The majority of oxygen comes from H2O2 so in adding more concentration, this
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