Why is iron (III) nitrate used as a catalyst in the hydrogen peroxide decomposition experiment?

Iron (III) nitrate lowers the activation energy of the decomposition reaction, allowing hydrogen peroxide to break down into water and oxygen much faster than it would naturally. The iron undergoes a redox reaction during the process and returns to its original oxidation state, which you can observe as a color change in the solution.

How does temperature affect the rate of hydrogen peroxide decomposition?

Higher temperatures increase the reaction rate, producing oxygen gas more quickly. By performing the decomposition at four different temperatures and measuring pressure changes, you can quantify this relationship and calculate the apparent activation energy using the Arrhenius equation to determine how temperature influences reaction kinetics.

What does the pressure change inside the flask indicate during the experiment?

The pressure increase is directly proportional to the rate of oxygen production from hydrogen peroxide decomposition. By recording pressure changes over time using a gas pressure sensor, you can determine the reaction rate at each temperature and track how quickly the decomposition occurs.

How do you calculate apparent activation energy from experimental data?

Plot the natural log of the pressure change rate against the reciprocal of absolute temperature to create an Arrhenius plot. The slope of this line equals negative activation energy divided by the universal gas constant. Multiply the slope by the negative gas constant to obtain the apparent activation energy value.

Why is it important to evacuate the flask to 10 kPa before each trial?

Evacuating the flask removes air and creates a sealed system where pressure changes accurately reflect only the oxygen produced from the reaction. This ensures that your pressure measurements are reliable and not affected by atmospheric air or leaks in the apparatus.

What is the expected range for the apparent activation energy of the catalyzed decomposition?

The iron-catalyzed decomposition of hydrogen peroxide yields an apparent activation energy between 35 and 60 kJ/mol, significantly lower than the uncatalyzed reaction at 78–88 kJ/mol. This difference demonstrates how effectively the iron catalyst reduces the energy barrier for the reaction.

Temperature Dependence

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The measure of how fast a reaction proceeds is called the reaction rate. The rate of a chemical reaction is defined by the rate law, which describes the relationship between the speed of the reaction and the reactant concentrations. In this equation, k is the rate constant, A and B are the two reactants, and m and n are their respective reaction orders.

The rate constant converts the relationship to the proper units of rate, moles per liter per second. Thus, the rate constant has different units, depending on the overall order of the reaction. However, the rate constant holds more significance than simply unit conversion. The rate constant is related to the minimum amount of energy required for a chemical reaction to occur - called the activation energy.

In a reaction, the reactants are at an initial state of potential energy. As the reaction proceeds, it must overcome a certain potential energy, the activation energy, before reaching its final state. The net energy of the reaction is the difference between the initial and final states. This difference can be negative, meaning that the reaction releases energy, or positive, meaning that it absorbs energy.

If there is not enough energy available to overcome the activation energy, the reaction will not proceed. In some cases, energy can be supplied in the form of heat. This provides additional energy to overcome the barrier to activation, and the reaction can proceed. A catalyst may also be added, which provides an alternative lower activation energy pathway between the reactants and products.

Catalysts are not consumed in the reaction and, therefore, do not affect the net energy of the reaction. The activation energy is determined experimentally, and it is related to the reaction constant k by the Arrhenius equation where A is the pre-exponential or frequency factor, R is the universal gas constant, and T is the absolute temperature at which the reaction occurs.

From this equation, we know that increasing the reaction temperature or decreasing the activation energy increases the rate constant. Going back to the rate law equation, it follows that a higher rate constant results in a higher reaction rate. This makes sense because as temperature increases, molecules move faster and collide more frequently, resulting in an increased fraction of molecules with higher energy than the activation energy.

In this lab, you will learn how to measure the activation energy of a reaction experimentally using the decomposition of hydrogen peroxide as the model reaction.

Print Procedure Steps

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

Temperature Dependent Decomposition of Hydrogen Peroxide
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.
Results
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.