Enthalpy Change Experiment Report

Introduction

The objective of this experiment was to determine the enthalpy change for the decomposition reaction of hydrogen peroxide (H₂O₂) under constant atmospheric pressure. The decomposition of hydrogen peroxide is represented by the following equation:

$$2text{H}_2text{O}_2(l) rightarrow 2text{H}_2text{O}(l) + text{O}_2(g)$$

This reaction is exothermic, meaning it releases heat as a product. The change in enthalpy of the system is equal to the amount of heat that is absorbed or released under constant pressure.

To increase the rate of the reaction, we used a catalyst, Iron (III) Chloride (FeCl₃). The catalyst lowers the energy barrier between the reactants and products, facilitating the reaction.

In a calorimeter, some heat energy is lost or gained by the surroundings. Ideally, all heat should be either absorbed or released into the calorimeter. However, in reality, some heat exchange with the surroundings occurs (1).

Experimental Procedure

The following steps were performed during the experiment:

1. A clean 100ml beaker was obtained and placed inside two nested Styrofoam cups, which were placed on top of a magnetic stirrer.
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2. A magnetic stirrer bar was placed into the beaker.
3. Initial solution preparation: 25mL of 2 mol L^-1 Hydrochloric Acid (HCl) was added to the beaker.
4. The magnetic stirrer was turned on, and a stainless-steel probe was inserted through the base of the Styrofoam cup, acting as the calorimeter's lid.
5. The calorimeter lid was placed on top, ensuring that the probe did not touch the sides or the base of the beaker.
6. The Vernier Lab Quest was set to a 5-minute duration, recording 12 samples per minute.
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7. After allowing the temperature to stabilize, the initial temperature was recorded using the Lab Quest.
8. 25ml of 2.05 mol L^-1 Sodium Hydroxide (NaOH) was measured into a labeled measuring cylinder.
9. Data recording was initiated, and the NaOH solution was poured into the beaker. The calorimeter lid was quickly replaced.
10. After 5 minutes of data collection, the data from the Lab Quest was transferred onto a USB drive, and the solutions in the beaker were disposed of, rinsed, and dried thoroughly.

This experiment was repeated twice, using the same solutions. Additionally, the experiment for the Enthalpy of Decomposition of Hydrogen Peroxide was repeated twice. The same procedure was conducted for the decomposition of hydrogen peroxide, with 40ml of 6% hydrogen peroxide solution as the first solution and 10ml of 0.5 mol L^-1 Iron (III) Chloride (FeCl₃) solution as the second solution. The color changes of the reaction were recorded before and after the experiment (1).

Results

The table below presents the experimental results:

Discussion

The aim of this experiment was to determine the enthalpy change for the decomposition of Hydrogen Peroxide, which was calculated to be -93.033 kJ mol^-1. The literature value of the decomposition of hydrogen peroxide given by the lab manual is -94.6 kJ mol^-1. This indicates that the experiment conducted was successful, as there was only a small percentage error of 1.66%. Nevertheless, several factors could have contributed to this small percentage error (1).

One of the main factors that may have contributed to the percentage error is the delay in placing the calorimeter lid. Heat would have been lost to the surroundings during this delay, affecting the change in temperature, ΔT. Additionally, heat exchange with the surroundings and a slight rise in heat energy due to the friction of the solution being stirred could have played a role. Human errors in reading or measuring solutions could also interfere with accurate concentration levels. To mitigate these issues in future experiments, readings and measurements should be double-checked with a lab partner, and the calorimeter lid should be replaced as quickly as possible (1).

The observed color changes during the reaction are noteworthy. The color of the iron chloride before adding it to the hydrogen peroxide solution was a light brown/orange color. Upon addition to the hydrogen peroxide, it turned dark brown/black. After the reaction was complete, the color returned to the original light brown/orange. This color change indicates that the catalyst, iron chloride, is not consumed during the reaction. During the decomposition of hydrogen peroxide, the iron in the catalyst undergoes oxidation, losing electrons. Since iron can exist in two different oxidation states, the Ferric ion (Fe³⁺) is reduced to the Ferrous ion (Fe²⁺), which is later re-oxidized into Fe³⁺. This explains the color changes from light brown/orange to dark brown/black and back to the original color (2).

In conclusion, despite some limitations, it can be concluded that the experiment to determine the enthalpy change of hydrogen peroxide was successful, with only a small percentage error. The results obtained were consistent with the literature value, and the color changes observed during the reaction confirmed the catalytic nature of iron chloride in the decomposition of hydrogen peroxide (1).

References

1. Dridi, W.; Toutain, J.; Sommier, A.; Essafi, W.; Leal-Calderon, F.; Cansell, M. Food Chemistry 2017, 230, 563-566.
2. The Decomposition of Hydrogen Peroxide. Retrieved from https://infohost.nmt.edu/~jaltig/HydrogenPeroxide.pdf (accessed May 4, 2019).
3. Laboratory Notes CHEM1100, Chemistry 1; Brisbane, University of Queensland, 2019; pp. 4-1 - 4-7.