The kinetics of the decomposition of hydrogen peroxide
Introduction
Kinetics refers to the rates of reaction. Thus, this experiment constituted measuring the rate of decomposition of hydrogen peroxide. This rate was measured by taking the volume of gas evolved as a function of time. The moles of dry gas produced were measured at different times after the reaction had been initiated. The volumes and moles are related in accordance with the Ideal Gas Equation, which states that the hydrogen peroxide concentration in the solution decreases as the volume of gaseous product formed per unit time decreases. When all the hydrogen peroxide was decomposed, its concentration was zero, and no more product was formed.
The purpose of the experiment was to determine the products of the reaction that took place during the decomposition of hydrogen peroxide, also to investigate the effect of changing concentrations and catalysts on the rate of reaction. The rate of the reaction is expressed using the rate-law expression which contains more information about the reaction, other than just the rate. The rate-law expression for iodide catalyzed decomposition of hydrogen peroxide is generalized as:
R = rate = k H2O2 n I m
Where the values for n and m were determined experimentally, and are referred to as the orders of the reaction with respect to hydrogen peroxide and iodide. The sum of the exponents is the overall order of the reaction. K is the rate constant, and the rate of the reaction was determined by the slowest step in a multi-step sequence of reactions. The experiment used stoichiometric concepts, gas collection equipment, and extensive data handling.
Procedure
Prepared 200 ml of a 0.15M solution of KI (yellow pebbles) as shown in the diagram below
Reaction of hydrogen peroxide with potassium took place. Place 30 drops of H2O2. Stood the test tube in the beaker and added KI solution to cover top of the spatula. Observe oxygen because it lit when burnt. Obtain and record the parametric pressure, and temperature. Poured 125 ml of distilled water into the bulb, and got 0.15 M KI to 20ml mark and poured into 125 ml elementary flask, and fell it with distilled water in a graduated cylinder. Observe the water level in the burette, dropped and stopped at level above the bulb. Record time and volume of the liquid after every 20 seconds. Repeat the procedure a second time, by adding 20 ml 0.15 M KI, and 10 ml H2O, 20 ml H2O.
Repeated a third time by adding 40 ml 0.15 M KI, and no water, 10 ml H2O2.
Results
Kinetics Data: Run 1 (20 ml KI, 20 ml H2O, 10 ml H2O2
Time in seconds 20 40 60 80 100 120 140 160 180 200 220 240
Volume 25.9 27.5 28.6 30 32.4 33.2 34.6 35.2 39.5 43.0 46.5 48.5
The volume increased gradually with time.
Kinetics Data: Run 2 (20 ml KI, 10 ml H2O, 20 ml H2O2)
Time in seconds 20 40 60 80 100 120 140 160 180 200 220 240
Volume 26.1 28.5 30.2 32.2 34 38.8 40.3 43.9 48.5 49.4 50
The volume increased gradually with time
Kinetics Data: Run 3 (40 ml KI, 0 ml H2O, 10 ml H2O2)
Time in Seconds 20 40 60 80 100 120 140 160 180 200 220 240
Volume 26.5 26 28.9 30.3 33.9 36.0 39.0 40.8 44.7 45 50.0
The volume increased gradually with time
In all the three runs, the volume is linearly related to the time. For the third run without water, the volume increase is lower, as compared to using 10 ml water.
Rate law
The correct rate law of the experiment is k[H2O2][KI].
Rate = 0.00726 [H202]^-1 [KI]^1/2
From the experiment, we can determine the order of reactions of the two catalysts by comparing the concentration of the solution to the first rate. We find that the concentration of KI is almost twice in part 2.
If the vegetable-extract-catalyzed decomposition of hydrogen peroxide were to give the same rate law expression, what would be the value of k for that reaction?
0.00325 = k(0.706)^-1 (0.1)^1/2
k = 0.00726
Suggest a mechanism that is consistent with the rate law you have determined
This mechanism resembles like that of the reaction of hydrogen peroxide and hydrazine
What relationship did you see between rates and concentration between run 1 and run 2, how did the concentration of hydrogen peroxide affect the rates
The reinforcement of the hypothesis of the experiment is clear from the tables. These are the best reference for the actual rate of reaction of the hydrogen peroxide. It is important to note that there was variation in the rates of reactions of hydrogen peroxide with the changes in the quantity of hydrogen peroxide, potassium iodide and variation in temperature. An increase in the amount of hydrogen peroxide lead to a decrease in the rate of reaction, while a decrease in the amount of the hydrogen peroxide lead to an increase in potassium iodide and a further increase in the rate of reaction. On the other hand, the a decrease in the amount of hydrogen peroxide lead to a decrease in the amount of potassium iodide and a further increase in the rate of reaction. Reducing the amount of potassium iodide increases the rate of reaction. However, when heat is introduced (by use of heated water at degree Celsius), the rate of reaction increased. This is because, the molecules gain more energy and move in the solution at increased speed, the rate of collision of the molecules leads to increased rate of reaction
What relationship did you see between rates and concentration between run 1 and run 2, how did the concentration of KI affect the rates
the rates of the reaction was affected by the amount of KI (concentration of KI), this is because, when the concentration of KI increased, the rate of reaction was increased, however, if the rate of concentration of KI decreased, the rate of reaction decreased. However, the rate of reaction between hydrogen peroxide and KI was directly proportional as an increase in the concentration of hydrogen peroxide. Increase in the concentration of KI decreases the rate of reaction while a decrease in the amount of KI increases the rate of reaction.
Generalisation concerning the rate concentration of Hydrogen peroxide and KI
From the experiment we can conclude that the hydrogen peroxide was more concentrated than the KI, because the amount of hydrogen peroxide has an impact on the rate of reaction of KI, while KI has no impact on the rate of reaction of hydrogen peroxide. The experiment can be likened to the decomposition of hydrogen peroxide with catalyse enzyme because it resembles an enzymatic reaction.
Higher temperature has a tendency of accelerating the rate of reactions. When the then amount of hydrogen peroxide is reduced and the amount of KI is kept constant, the rate of reaction increases. This may be because the amount of KI has the capacity to accelerate the reaction of only half of the hydrogen peroxide. This takes less time. However, if the amount of KI is reduced, the rate of reactions slows down remarkably, because the amount of catalyst used was very low to catalyze the reaction of hydrogen peroxide. We find that the rate expression is in line with the formula bellow for the rate of reaction
Rate of reaction = R = k [A]^m x [B]^n
4) Conclusion about the two catalysts
The two catalysts in the experiment have equal molarities and their density are sale the same, the first catalyst may be an enzyme catalyse as the reaction is zero order, however, hydrogen peroxide is very reactive as compared to potassium iodide. It is also healthy to conclude de that the reaction rate of the decomposition of hydrogen peroxide can increase with the increase in the temperature. It is also right to conclude that that potassium iodide and hydrogen peroxide are example of the first order reactants
