Temperature Dependence

Source: Smaa Koraym at Johns Hopkins University, MD, USA

1. Temperature Dependent Decomposition of Hydrogen PeroxideExpand

   We suggest that students work in pairs for this experiment. Equipment controls may vary.

   In this lab, you'll perform a decomposition reaction where a single compound breaks down into two or more simpler products. You will observe the decomposition of hydrogen peroxide into water and oxygen. This decomposition happens very slowly, so you will use iron (III) nitrate as a catalyst to lower the activation energy.

   During this process, iron undergoes a redox reaction and then returns to its starting oxidation state. You will be able to see this as a color change in your solution during the reaction. You'll perform the same reaction at four different temperatures and track the rate of the reaction by recording the pressure inside the flask. This way you can measure how quickly oxygen was produced in each experiment. You'll use this to calculate the reaction's activation energy.

   Table 1. Estimate the apparent activation energy for the decompostion of hydrogen peroxide

   Trial	Temperature (℃)	ΔP (kPa/s)	1/T	ln (ΔP)
   1
   2
   3
   4

   Click Here to download Table 1  

   • Before starting this experiment, put on a lab coat, safety glasses, and nitrile gloves.
   • Ensure that your hot plate is off and then place a 600-mL beaker on the hot plate.
   • Connect vacuum tubing to the barbed sidearm of a 125-mL Büchner filter flask, and carefully clamp the filter flask in the 600-mL beaker so that the sidearm is just above the top of the beaker.
   • Fill a 400-mL beaker with deionized water., and pour the water into the 600-mL beaker until the water level is about 2–3 cm below the side arm.
     NOTE: Make sure that the water is not able to get into the flask or vacuum line.
   • Immerse the thermometer in the water, clamping it in place so that the end is in contact with the outer wall of the flask. Ensure that you can read the current temperature, as well as the 40, 60, and 80 °C marks.
   • Make sure that each hole of the 2-hole adapter is fitted with a tapered locking adapter. Check that all the adapters are firmly seated, since any air that escapes around them will affect your data.
   • Lock a 2-way stopcock into one of the stopper adapters. Then, lock one end of the flexible tubing into the second adapter with the other end into the connector of the gas pressure sensor.
   • Turn on the data acquisition device for the gas pressure sensor and ensure that the pressure is displayed in kPa.
   • Set the acquisition rate to two samples per second and the duration to 300 s. Then, display the pressure in real-time.
   • Label a 400-mL beaker as ‘waste’, a 50-mL beaker as ‘0.5 M Fe(NO₃)₃’, and a 100-mL beaker as ‘3% w/w hydrogen peroxide’.
   • Place a few paper towels on the work surface as a clean area for glassware that you will reuse. Keep a supply of paper towels on hand for later.
   • Then, bring the 50-mL and 100-mL beakers to the stock solution area. Pour about 30 mL of 0.5 M Fe(NO₃)₃ in the 50-mL beaker and 100 mL of 3% w/w hydrogen peroxide in the 100-mL beaker.
   • Back at your workspace, setup a 20-mL volumetric pipette and fill it to the mark with 3% w/w hydrogen peroxide. Dispense the hydrogen peroxide into the filter flask and set the volumetric pipette aside.
   • Adjust the thermometer as needed so that it contacts the flask below the hydrogen peroxide solution level.
   • Fit the rubber stopper into the mouth of the filter flask, being careful not to loosen the connections to the sensor and the 2-way stopcock.
   • Check that the 2-way stop is closed. Open the vacuum line and monitor the pressure as it decreases. This will seal the stopper in the flask. Once the pressure reaches 10 kPa, close the vacuum.
   • Monitor the pressure for at least 1 min to confirm there are no slow leaks.
     NOTE: If the pressure immediately starts rapidly increasing, there is a leak in your setup, so tighten the connections and try again until the pressure holds at 10 kPa when the vacuum line is closed.
   • Draw 5 mL of 0.5 M Fe(NO₃)₃ into a 20-mL syringe. Expel all air from the syringe so that it only contains the solution.
   • Lock the syringe into the top of the 2-way stopcock. You are now ready to start the room temperature trial, so record the water temperature in your lab notebook.
   • Start acquiring gas pressure data. Allow the device to record data for about 15 s, then open the stopcock and quickly close it once all of the Fe(NO₃)₃ has entered the flask. The observed increase in pressure is from the evolution of oxygen gas produced by the decomposition of hydrogen peroxide.
   • Once the data collection finishes, save your data. Then, disconnect the syringe and open the stopcock to vent the flask.
   • Take out the rubber stopper, carefully remove the thermometer from the clamp and the flask from the beaker, and disconnect the vacuum tubing from the filter flask.
   • Empty the flask into the waste beaker. Try not to get liquid in the side arm. Dry the outside of the flask with paper towels.
   • Rinse the interior of the flask with deionized water and pour the rinse into the waste beaker. If any liquid gets into the side arm, remove it with paper towels.
   • Reconnect the flask to the vacuum line and clamp it in the 600-mL beaker in contact with the thermometer.
   • Turn on the hotplate and heat the water around the flask until the thermometer reads 80 °C. Then, turn off the heat.
   • Then, add 20 mL of 3% w/w hydrogen peroxide to the flask.
   • Dry the stopper with paper towels. Do this after each trial to ensure that the stopper has a snug fit in the neck of the beaker and fit the stopper in the flask.
   • Check that the stopcock is closed and evacuate the flask to about 10 kPa. Then, close the vacuum and confirm that there are no leaks.
   • Draw 5 mL Fe(NO₃)₃ solution into the syringe, expel the air from the syringe, and connect it to the stopcock.
   • Record the temperature shown on the thermometer in your lab notebook, then, start collecting data.
   • Wait about 15 s and introduce the Fe(NO₃)₃ solution in the same way as before.
     NOTE: As the pressure approaches 150 kPa, the stopper may pop off!
   • After data collection finishes, save the data, vent and clean the filter flask, and set up for the third trial.
   • Fill the 400-mL beaker with ice and add some to cool the water to about 60 °C.
   • Perform the third trial in the same way as before. Remember to dry the stopper and the inside of the flask and to record the temperature prior to data collection.
   • After the data is collected for that trial, cool the water to about 40 °C and perform the fourth trial.
   • Once you have finished all four trials, empty the filter flask and rinse it into the waste beaker one last time.
   • Dispose of any excess 0.5 M Fe(NO₃)₃ and the contents of the waste beaker in a container labeled for aqueous iron waste.
   • Then, disassemble the apparatus and pour the leftover ice, water, and hydrogen peroxide in the sink. Wash your glassware following your lab's standard protocol.
2. ResultsExpand

   The decomposition of hydrogen peroxide with iron is a complex, multi-step process that we can't easily describe in a single equation. However, we can estimate the apparent activation energy from the rate of oxygen production and compare it to the apparent activation energy of the un-catalyzed process.

   • Find the rate of pressure change, which is directly proportional to the rate of oxygen production. For each experiment, make a plot of pressure with respect to time and find the point where the reaction started.
   • Identify the maximum pressure attained and determine the slope between the two points.
   • Once you've determined the slopes and the corresponding temperatures in Kelvin for all of the temperatures, use the Arrhenius equation to estimate the apparent activation energy of this reaction.
   • Take the reciprocal of the temperatures in Kelvin and the natural log of the rate of change in pressure. Remember that the rate constant k is essentially equal to the change in pressure.
   • Make an Arrhenius plot and find the slope of the line.
   • The slope is equal to the negative activation energy over the universal gas constant, so multiply the slope by the negative universal gas constant to get the apparent activation energy of the iron catalyzed decomposition reaction. You'll see a value in the range of 35–60 kJ/mol because the iron catalyst made the decomposition take less energy. The apparent activation energy of the uncatalyzed decomposition of hydrogen peroxide is around 78 – 88 kJ/mol.

In this lab, you'll perform a decomposition reaction where a single compound breaks down into two or more simpler products.You will observe the decomposition of hydrogen peroxide into water and oxygen.This decomposition happens very slowly, so you will use iron nitrate as a catalyst to lower the activation energy.During this process, iron undergoes a redox reaction and then returns to its starting oxidation state.You will be able to see this as a color change in your solution during the reaction.You'll perform the same reaction at four different temperatures and track the rate of the reaction by recording the pressure inside the flask.This way you can measure how quickly oxygen was produced in each experiment.You'll use this to calculate the reaction's activation energy.Before starting this experiment, put on a lab coat, safety glasses, and nitrile gloves.Now, ensure that your hot plate is off and then place a 600-milliliter beaker on the hot plate.Next, connect vacuum tubing to the barbed sidearm of a 125-milliliter B