CHE113 experiment 14 PDF

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Taylor Dorsey CHE113-052 TA: Sarah Kriger Tuesday, October 6th 2020

Using Rate Laws to Determine Reaction Rate and Decomposition of Hydrogen Peroxide Introduction The purpose of this experiment was to determine the decomposition of hydrogen peroxide. This was done by determining the general rate law of a reaction, the rate constant for the reactions, and the activation of a reaction from two or more trials (French et al. 81). The functions and properties of catalysts needed to be understood as well as dilution calculations and calculations to convert between mass percent and molarity (French et al. 81). In order for a chemical reaction to occur, molecules within the reaction must collide with enough kinetic energy to create vibrational energy that breaks existing bonds, or else they would just bounce off one another. This minimum required energy to initiate a reaction is called the reaction’s Activation Energy (Ea) (French et al. 82). The rate of a chemical reaction is defined as the speed at which the reaction proceeds. The rate of the reaction can be influenced by many factors. The concentration of a reaction can change the reaction’s rate due to the amount of molecules colliding. The more molecules present within the reaction, the more molecules are going to collide. This creates a higher kinetic energy and will speed up the reaction. The kinetic energy of a reaction can also be influenced by the temperature at which the reaction takes place. When temperature increases, molecules move faster which causes a raise in kinetic energy and a higher amount of molecular collisions. A

catalyst is another factor that can influence a chemical reaction. A catalyst lowers a reaction's activation energy, causing less energy to be needed to initiate the reaction. While a catalyst takes part in the reaction, it is a homogeneous or heterogeneous substance that does not get consumed during the reaction and always returns to its original composition (French et al. 83). The rate of a reaction can be shown by the equation below, where the rate equals the change in concentration of the product over the change in time. rate = Δ[C]/Δt As previously stated, the rate of a chemical reaction can change as the reaction’s concentration changes. The relationship between reaction rate and concentration is defined by the rate law. The rate law can be shown by the equation below, where k is the rate constant of the reaction at a given temperature and m and n are defined as the order of the reaction with respect to A an B (French et al. 83). The sum of m and n give the order of the overall reaction, which must be determined experimentally. Rate = k [A]m [B]n The Purpose of this lab is to determine the decomposition of hydrogen peroxide (H2O2) which is defined by the chemical equation below. 2 H2O2 (aq) ⇋ 2 H2O (l) + O2 (g) The decomposition of hydrogen peroxide is very slow, therefore potassium iodide (KI) was used as a catalyst. The rate of the reaction was determined by measuring the pressure increase produced by oxygen gas (O2) during the reaction. This was done by using the ideal gas law. The equation the represents the ideal gas law is defined as; PV = nRT This equation was then rearranged to present in units of mol/L;

n/v = P/RT Whereas n/v is equal to molarity, the following can be stated; M = P/RT When both sides of the above equation are divided by seconds, the molar rate constant can be determined (French et al. 85). The rate law of the reaction was determined by varying the initial molar concentrations of the hydrogen peroxide and the concentrations of potassium iodide. Trials were run at different temperatures to determine the activation energy of the reaction. The Arrhenius Equation was used to show how the rate constant will vary depending on the temperature of reactions with the same concentrations (French et al. 85). The natural logarithm was then taken to produce an equation for a linear relationship where the slope can be used to calculate activation energy (French et al. 85). K = AeEa/RT ln(k) = (-Ea/R)(1/T) + ln(A) This equation can be rearranged to calculate activation energy without the need to graph; ln(k2/k1) = Ea/R [1/T1 - 1/T2] Mass percent was used to calculate the molarity of each solution in order to find the rate constant. The following equations show mass percent and molarity are calculated. Mass percent = mass of solute / total mass of solution Molarity = moles of solute / L of solution The mass percent of the hydrogen peroxide is given as 3%. These equations were then used to convert the mass percent to grams of hydrogen peroxide. The grams were then converted to moles using the molecular weight and density of the compound, and the moles were then used

to find the molarity. The density was assumed to be 1.00 g/ml due to the solutions being dilute aqueous (French et al. 85). In order to determine the rate law, multiple solutions of varying concentrations had to be used during the experiment. Dilutions were made by using the following dilution equation. MiVi = MfVf Mi and Mf represent the initial and final concentrations, Vi and Vf represent the initial and final volumes.

Methods 1. Copy table 14.1 from the lab manual into the lab notebook. 2. Obtain the following materials; a. Gas pressure probe kit and measureNet b. Thermometer c. 3% hydrogen peroxide solution d. 0.5 M potassium iodide solution e. Four test tubes f. Pipet and pipet bulb g. 3 beakers of varying sizes h. 400mL beaker i. Graduated cylinder j. Wash bottle 3. Set up the MeasureNet with the gas probe kit; a. “Main Menu” → “Pressure” → “Pressure v. Time” → “Calibrate” →

“760mmHg” → “Display” 4. In the three small beakers, collect approximately 16mL of hydrogen peroxide, 5mL of potassium iodide, and 4mL of DI water. Collect each solution in its own clean and dry beaker, make sure to rinse the graduated cylinder after each measurement. 5. Set up the four clean and dry test tubes in the test tube tray. Label the test tubes according to their trial. 6. Prepare the solution for the first trial in the appropriate test tube according to the volumes given in table 14.1. Immediately after mixing the solutions, place the gas probe pressure cap in the test tube in a way no air can get through and start collecting data by pressing “start/stop” on the MeasureNet. Keep the thumb delicately placed over the cap to avoid the cap popping off, be careful not to put pressure on the cap whereas this could alter results. Collect data for 100 seconds then save the file as “001” for trial one. 7. Repeat step six for trials two and three using the appropriate volumes according to each trial on table 14.3. Save trial two as “002” and trial three as “003”. 8. For trial four, gather 400mL of hot water in the 400mL beaker prior to mixing the solution. Make sure the water is at least 10℃ hotter than room temperature. 9. Repeat the process for step six with the appropriate volumes of each solution for trial four and keep the test tube placed in the hot water beaker for the entirety of the trial. Run this trial for only 50 seconds and be aware of the cap popping off whereas the pressure may increase faster. Save the file as “004”. 10. Be sure to clean instruments with DI water in between each use, record all exact volumes and temperatures in the lab notebook. 11. After each set of data has been approved, clean materials and lab bench.

Discussion The purpose of this experiment was to determine the decomposition of hydrogen peroxide. This was done by determining the general rate law of a reaction, the rate constant for the reactions, and the activation of a reaction from two or more trials (French et al. 81). The functions and properties of catalysts needed to be understood as well as dilution calculations and calculations to convert between mass percent and molarity (French et al. 81). Using the relationships between mass percent and molarity, the molarity of the 3% hydrogen peroxide was determined to be 0.882M. Using this data, the molarities of each solution in each trial were able to be calculated. The rate formation of O2 in torr/s was then derived from the slope of the Pressure vs. Time graph that was measured during the lab. The rate formation of O2 was then converted to mol/L•s-1 in order to find the rate of decomposition of hydrogen peroxide. Results are listed in the table below. Trial

Decomposition of H2O2 (mol/L•s-1)

1

0.00052

2

0.000151

3

0.000233

4

0.00182

It was then determined that the order of the reaction with respect to H2O2 was 1 and the order of the reaction with respect to KI was 2, creating the following rate law; Rate = k [H2O2]1 [KI]2 This equation was then used to determine the rate constants for each trial. The results are

shown in the table below. Trial

Rate Constant k (M-1•s-1)

1

0.074

2

0.067

3

0.069

4

0.24

After all calculations were made, the Arrhenius Equation was used to calculate the activation energy of the reaction. According to the data and calculations, the reaction’s activation energy was 45 kJ/mol. Overall, the purpose of the lab was supported because there was success in determining a rate law for the reaction, the decomposition of hydrogen peroxide, the rate constant (k) for each trial, and the activation energy. The data also supports the hypothesis of the reaction rate increasing with temperature with all other variables being held constant. Although the results supported the original purpose and hypothesis of the lab, potential sources of error may have hindered some of the data creating numbers that may have been slightly off. One potential source of error was the amount of pressure that escaped the reaction prior to putting the cap on the tube or prior to starting the MeasureNet. This would have altered the Pressure v. Time graph, giving inaccurate results for the decomposition of the hydrogen peroxide. This could have been prevented by having two people working together to accurately time the capping of the test tube. Another potential source of error includes cross contamination due to the pipet bulb. While other instruments can be cleaned with DI water in between trials, the pipet bulb is not so easily cleaned and could have caused cross contamination that would have affected the concentrations of solutions. This could have been prevented by using a new bulb

each time a solution was measured using the pipet. A third source of error may have been inaccurate volume measurements during the collection of each substance. This could have been due to bubbles in the pipet during collection or inaccurate reading of the numbers on the pipet. This can be prevented by tapping the pipet to rid it of all bubbles and double checking each measurement. Conclusion In this experiment I learned how to determine the activation energy of the decomposition of hydrogen peroxide through determining the general rate law of a reaction and the rate constant for the reactions. I also learned how to utilize the functions and properties of catalysts as well as dilution calculations and calculations to convert between mass percent and molarity. Chemical Kinetics are important for real world applications as well. This concept is important in the pharmaceutical profession and is utilized when testing the stability of certain drugs and knowing how long a drug lasts in the body after ingesting. On the other hand, in the food industry, many bakers will use catalysts such as baking soda or increase the cooking temperature to increase the speed at which some foods cook.

Work Cited French, April, et al. General Chemistry II Laboratory Manual. Plymouth, MI: Macmillan Learning Curriculum Solutions, 2020. pp 81-86....