This experiment is revolves around the decomposition of H peroxide. There were several aims for this experiment: Monitor the rate of a chemical reaction, find the kinetic order, happen the energy activation and understand how molecular constituents map in a rate restricting measure of the reaction. Two reagents were examined were, and KI. 8 tests were carried out with 4 differing concentrations for each reagent and two different temperatures were used. The force per unit area and temperature was recorded utilizing Logger Pro. The experiment found that the decomposition of H peroxide is increased from the add-on of the accelerator KI. Specific computations found the kinetic rate order, m, n, rate ( m/s ) , rate invariable and energy of activation.
Hydrogen peroxide is a vitally of import oxidising agent in the industry. It is used for assorted things and interruptions down easy over many old ages. The decomposition of H peroxide generates oxygen as a merchandise. . This O can so be analyzed utilizing a closed vas by mensurating the alteration in force per unit area.
To analyze the chemical dynamicss behind the decomposition of ; a major oxidising agent in industry. By presenting a accelerator, such as KI, the decomposition will be sped up. This reaction releases. The alteration in force per unit area is so analyzed utilizing a force per unit area proctor on Logger Pro. The kinetic order of both and KI and activation energy to get down the reaction can besides be measured.
The experiment was based on several aims. First, supervise the rate of chemical reaction. Besides, the experiment will concentrate on how to find the kinetic order of a reaction from the dependance of the rate on reactions. Additionally, the experiment will assist one understand how the molecular constituents of the rate restricting measure of the reaction are determined by the kinetic order of the reaction. Last, to happen the activation energy of the reaction from the temperature dependance of the reaction rate.
The intent of the experiment is to find the kinetic order for both and KI. Additionally, the experiment will look to find the energy of activation for the reaction.
Before the experiment could get down solutions were prepared. 10 milliliter of 3 % was placed into a trial tubing. A big beaker was so filled with 300-500 milliliter of room temperature H2O. The trial tubing was so placed in the beaker and a investigation inserted to mensurate H2O temperature. 50 milliliter of.5M KI was so measured out in a 100 milliliter graduated cylinder. The KI was transferred into a 125 milliliter Erlenmeyer flask. 10 milliliter of KI was measured and placed into a trial tubing which was placed into the same H2O bath as above. The staying KI was covered and set aside.
A 25 milliliter Erlenmeyer flask was obtained and secured in the H2O bath. 2 milliliter of was pipetted from the trial tubing into the flask. 1 milliliter of.5M KI was rapidly pipetted into the flask. The flask was swirled, connected to the force per unit area proctor and placed into the H2O bath. Using the logger pro device, information was collected for 500 seconds. Temperature and force per unit area were monitored during the experiment. The experiment was repeated 8 times with changing sums of both reagents. A sum of 4 different concentrations for each reagent. Water was added into the flask to accomplish different concentrations. The last experiment changed the H2O bath temperature by 10 K. The information was so graphed on logger pro, plotting force per unit area by clip.
Before the experiment could get down, the sum of H2O added to the flask to thin the reagent, needed to be calculated. This was done utilizing the equation. Where M is the molar concentration and V is the volume. The gave the concentration needed for the reagent being examined. Multiplying showed the sum of H2O needed to be added for the right concentration to be achieved. Trial 1-4 changed the concentration of KI while test 5-8 changed the concentration of. The information was so examined on logger pro and an mean temperature was taken for each test. A additive arrested development line calculated the mean force per unit area on the graphs. With this force per unit area, the equation was used to cipher the M for. The reaction order for, expressed as m, and the reaction order for KI, expressed as N, were found diagrammatically. The m value was found by plotting log rate ( M/s ) as a map of the log of concentration, while the n value was found by plotting log rate ( M/s ) as a map of the log of KI concentration. The rate jurisprudence expresses m and N in the equation. After m and Ns were determined diagrammatically, the equation, was used to happen the the rate invariable ( K ) . The standard divergence and mean of the rate invariables were besides calculated for tests run near room temperature. Finally, the activation energy was calculated for the two tests performed at different temperatures. The equation used was where R= 8.314 J/mol*K.
After all the information was collected. Consequences were analyzed. The undermentioned several pages outline the consequences of the dislocation of Hydrogen Peroxide by usage of a accelerator, Potassium Iodide. ( Graph 1 ) shows the force per unit area vs. clip secret plan of for test one of the experiment. The following figure ( Table 1 ) outlines the rate of the reaction in M/s with partnered temperature and the concentrations of both and KI. The two graphs accompanied close to each other ( Graph 2, Graph 3 ) item the kinetic order of the reaction with regard to each reagent. These graphs calculated incline on a additive patterned advance line which analyze the m and n value for and KI. The rate jurisprudence is so shown where m and N were taken from ( Graph 2, Graph 3 ) . The following tabular array ( Table 2 ) gives the rate invariable ( K ) for each test. A standard divergence and mean rate changeless were calculated from tests 1-7. The last equation, energy of activation shows the energy of activation for the reaction run at two different temperatures
A secret plan of force per unit area vs. clip for a individual test is shown below ( Graph 1 ) . This graph shows that as the clip increased so did the force per unit area. This relationship seems to be additive with a positive correlativity of.9974. The incline of the graph is.07959 kPa/s and a y-intercept of 99.51 kPa.
Graph 1: Pressure vs. Time of.88 ( mol/L ) for test figure one of the information set. had a concentration of.88 mol/L and KI was.5 mol/L
The rates for each test were calculated below ( Table 1 ) . Trial one is the foundation of all tests with the stock solutions used for the reactant and accelerator a concluding concentration of Hydrogen Peroxide and KI was.59 mol/L and.17 mol/L. For tests 2-4 the concentration of decreased while the concentration of KI stayed the same. Tests 5-7 saw a reduced concentration of KI while stayed changeless.59 mol/L. The rates for each test are detailed on the far right column. The temperature is besides shown in the 2nd to last column in Kelvin. For tests 1-7 the temperature was kept around room temperature of 294 K. In test 8 the temperature was increased by 10 K to 304 K. The concentration of both solutions is given in the in-between two columns. The rate neared zero as concentrations of either or KI decreased except when concentration was reduced to.15 mol/L. This shows that in order for the reaction to happen at that place needs to be a minimal concentration of either solution.
Concluding Concentration ( mol/L )
KI Final Concentration ( mol/L )
Temperature ( K )
Rate ( M/s )
Table 1: Concluding Concentrations of both and KI for each of the 8 tests. Temperature is given in Kelvin with all but 1 test around room temperature. Rate of the reaction was calculated for each test.
The log ( rate ) vs. log ( concentration ) of for tests 1-4 was used to happen the m value for for ( Graph 2 ) . The m value is the kinetic order of the reaction for Hydrogen peroxide. The incline of the graph is.947. With regard to the kinetic order ; this was rounded to 1. This means that the m value for is 1. The correlativity for this graph was really high significance that the relationship between the two is interdependent on one another.
Graph 2: Log ( Rate ) vs. Log ( Concentration ) of refering to trials 1-4. The incline of the graph gives the m value for the kinetic order of the equation
The log ( Rate ) vs. log ( Concentration ) of KI for tests 1, 5-7 is shown below in ( Graph 3 ) . This graph was used to happen the n value for KI with regard to the kinetic order of the rate equation. The graph has a positive incline of.6406. This was rounded to 1. The 1 value represents the n value in the rate equation. As the log ( concentration ) increased so did the log ( rate ) demoing a positive correlativity between the two. The correlativity was really strong at.9062 portraying a strong relationship between the two
Graph 3: Log ( Rate ) vs. Log ( Concentration ) of KI for tests 1, 5-7. The incline gives the n value for KI with regard to kinetic order of the reaction
The rate jurisprudence for the decomposition of catalyzed by KI was found diagrammatically from ( Graph 2, Graph 3 ) . The equation is. This means that it is a 1:1 reaction order for both reagents.
The rate invariables, K, for each test are given in the tabular array ( Table 2 ) . The mean rate changeless and standard divergence of the rate invariable are given on the underside of the tabular array. They pertain merely to the tests carried out near room temperature of 294 K. Rate invariable is highest in test 5. The starting concentrations for and KI were.88 mol/L and.25 mol/L severally. At this point, was kept to its stock concentration while KI was halved. The standard divergence of the rate invariable is 1.26E-04 is really little, intending the computations and experiment was carried out with really small mistake.
Rate Constant ( K )
Temperature ( K )
Average of Changeless
Standard Deviation of Constant
Table 2: The deliberate rate invariable ( K ) for each test. The mean rate changeless and standard divergence of rate invariable for the room temperature values for tests 1-7. Trial 8 had a temperature of 304 K
The energy of activation was 28,252 J/mol and 28.252 kJ/mol. The equation was rearranged to happen. . This is the minimal sum of energy required to convey about the reaction
KI increases the rate of decomposition of H peroxide as seen in this experiment. It does non undergo a chemical alteration as first idea. Since KI is fundamentally recycled it ne’er has to undergo a net alteration in its chemical composing. It is consumed in the reaction but reproduced each clip. The ion is the catalyzer in the reaction of H peroxide. The ion is consumed in one reaction and reproduced in a series of stairss in the following reaction. The rate jurisprudence helps to explicate the phenomenon of rate restricting measure of the reaction.
Since the rate order of the reaction is first order, it can be explained that the first measure is the rate restricting measure of the reaction. Before hydrogen peroxide can break up it must be catalyzed by the KI. By definition, the activated composite is a species at energy extremum that must fall apart to organize merchandises or reform reactants. In order for the rate restricting measure to happen, the H peroxide and K iodide must fall apart to reform reactants. This measure is the slowest because in order for this to go on a buildup of energy must happen. Once the energy reaches its extremum, the reaction picks up velocity and the reactants are reformed over and over until the reaction exhausts itself. At this point the merchandises are formed, H2O and O.
In footings of possible chemical stairss for this contact action that are consistent with the ascertained dynamicss ; the simple stairss of the reaction coincide with the slow decomposition of H peroxide, reformation of KI and chemical exhaustion of both reagents to organize H2O and O.
This experiment was interesting due to the fact that it involved assorted stairss and equations. I was able to get the hang several equations and learn legion constructs. I learned how to cipher the energy of activation, which is needed in order for the reaction to happen. Additionally, the alteration in concentration of either the reactant or accelerator can greatly impact the rate at which a reaction occurs. Besides, I found out how the kinetic order ties into the basic foundation for the reaction and that this can be found diagrammatically by plotting the log ( rate ) vs. the log ( concentration ) .
I think that this experiment was really helpful non merely in maestro certain accomplishments within the research lab but besides wining in understand the footing for all chemical reactions. I did non cognize chemical dynamicss was such an of import portion of reactions and it will decidedly assist me to farther break my apprehension of chemical science in general at a molecular degree.