Rate of Decomposition of Hydrogen Peroxide (H2O2) using various catalysts | KI, MnO2 and Fe(NO3)3 | Kinematics of Reactions

What is the effect of KI, MnO₂ and Fe(NO₃)₃ as a catalyst on the rate of decomposition of Hydrogen peroxide (H₂O₂) at room temperature?

INTRODUCTION

Chemical kinetics discusses about the rate of a reaction and how various parameters such as temperature, catalyst, concentration and pressure affect the rate at which the reactant gets converted into a product in a chemical reaction. The decomposition of Hydrogen peroxide (H₂O₂) is a reaction used as an example to explain the effect of catalyst on the rate of a reaction. In the past H₂O₂ was used to clean wounds because the enzyme present in the blood, catalase, can decompose H₂O₂ into water (H₂O) and oxygen (O₂). The nascent oxygen released during the process is highly effective against bacterial growth in the wound ^[1]. This shows how catalase acted as a catalyst that facilitated the increased rate of decomposition of H₂O₂. 

The decomposition of hydrogen peroxide using catalyst has widespread application in the domestic and industrial sector^[2]. Seeing this immense application of decomposition of hydrogen peroxides using catalysts ignited the curiosity in me to know how different catalyst behaved with hydrogen peroxide and how the speed of the reaction varied with each catalyst. This ultimately led to the research question, What is the effect of KI, MnO₂ and Fe(NO₃)₃ as a catalyst on the rate of decomposition of Hydrogen peroxide (H₂O₂) at room temperature. From the literature survey, it was seen that transition metals and iodide were used for catalytic decomposition of H₂O₂^{[3], [4]}. Therefore, MnO₂, KI and Fe(NO₃)₃ were selected as a catalyst to decompose H₂O₂ (1 mol dm^-3). The amount of O₂ released in a given span of time is measured and the obtained graph and data will be used to compare the rate of reaction. During the investigatory period, it was learnt that a high concentration of H₂O₂ was a dangerous chemical to experiment with^[5]. Therefore, a lesser concentration of H₂O₂ was used for the experiment and proper safety measures were taken into consideration while conducting the experiment. From the literature survey, it was found that a lower concentration of H₂O₂ was used to conduct the experiment to get favourable and proper data for the study

BACKGROUND

Catalyst is a chemical substance that is used to increase the rate of a reaction without being consumed in a reaction^[6]. It provides with a different reaction pathway by altering the activation energy which is responsible for the initiation of a reaction.

Rate of a reaction is defined as the change in concentration of the reactant or the product with respect to time^[6]. The unit used for the rate reaction is mol dm^-3 s^-1. For a given reaction:

the rate at which the reactants are consumed will be equal to the rate at which the product is formed. That is:

HYPOTHESIS

I predict the order of the effective catalyst to decompose H₂O₂ as: . This order is based on the assumption of higher the oxidation state of the element reacting with hydrogen peroxide, the better it decomposes. 

VARIABLES

·       Independent variables: The catalyst used. Three different types of catalyst will be used for decomposing H₂O₂ and all the three catalyst will be in solid form. The three catalyst are MnO₂, KI and Fe(NO₃)₃.

·       Dependent variables: Amount of oxygen (O₂) produced. This was taken as a dependent variable because the change in concentration of the product with respect to time is very much required to quantify the rate at which hydrogen peroxide decomposes, in general, rate of the reaction. The concentration of the evolved oxygen will be in terms of percentage (%).

·       Controlled variables:

1.     Temperature: H₂O₂ is a self-decomposing chemical substance due to its unstable peroxide bond (oxygen-oxygen single bond) and at higher temperatures it decomposes faster^[2]. If exposed to sunlight, self-decomposition may happen which will alter the concentration and bring in errors to the data collected. Therefore, the experiment was performed at constant room temperature.

2.     Concentration of H₂O₂: Concentration is one of the parameters that affect the rate of a reaction since collision frequency of particles increases with concentration which can affect the rate of the reaction^[6].

3.     Amount of catalyst used: The amount of catalyst added into H₂O₂ was kept constant (1g) for each trial using the three catalysts. A low amount of catalyst was used such that the progress of the reaction was easy to observe.

4.     Volume of H₂O₂ used: Low concentration and lesser amount of H₂O₂ (10 cm³) was used for the experiment for observing the progress of the reaction more effectively.

METHOD:

Materials required

1.     Hydrogen peroxide, H₂O₂ (0.1 mol dm^-3)

2.     Manganese (IV) dioxide, MnO₂

3.     Potassium iodide, KI

4.     Ferric Nitrate, Fe(NO₃)₃

5.     250 cm³ conical flask

6.     Weighing balance

7.     10 cm³ pipette

8.     Measuring cylinder

9.     Vernier O₂ gas sensor

10.  Vernier Lab quest

11.  Wax 

Experimental Procedure

Note: This experiment was conducted at room temperature and in a room where there wasn’t much exposure of sunlight.

1.       0.1 mol dm^-3 of standard solution of hydrogen peroxide (H₂O₂) was prepared by diluting a 30% W/V H₂O₂. The prepared 0.1 mol dm^-3 of H₂O₂ was stored in a dark room because exposure to sunlight can lead to self-decomposition. 

2.       1g of the selected catalyst was weighed using a weighing balance and transferred into a 250 cm³ conical flask.

3.       The vernier O₂ sensor was activated by connecting it into the lab quest. It was made sure that the data collection time is set for 180 seconds and 4 seconds per sample.

4.       10 cm³ of the prepared 0.1 mol dm^-3 of H₂O₂ was measured using a pipette and later transferred into a 10 cm³ measuring cylinder. The solution was transferred into a measuring cylinder because the transferring of solution into the conical flask which contains the catalyst will be easy. Pipette takes time to transfer the solution and this can affect the data since the data collection using the O₂ sensor should be started the exact moment when the solution starts reacting with catalyst.

5.       Now the H₂O₂ solution present in the measuring cylinder was transferred into the conical flask which contains the catalyst.

6.       The O₂ sensor was immediately inserted into the mouth of the 250 cm³ conical flask and the data collection was started. Some amount of wax was applied around the mouth of the conical flask in order to seal the smallest of the smallest gaps which can lead the liberated O₂ to escape. A 250 cm³ conical flask was used because its diameter was proper to fit the O₂ sensor tightly thereby reducing the chance of oxygen to escape.

7.       After 180 seconds (3 min), the lab quest produced an O₂% v/s time graph which showed the evolution of oxygen inside the conical flask with respect to time. The data and graph was saved for further analysis and trials.

8.       The O₂ sensor was removed from the mouth of the conical flask and was left open such that it gets matches with the O₂% present in the room. 

9.       4 trials for the selected 3 catalyst were conducted by repeating steps 2-7. It was made sure that a new 250 cm³ conical flask which is clean and dried is used for further trials. This was done to reduce the possibility of error due the presence of any chemical stain or water around the surface of the conical flask.

Safety measures

1.       H₂O₂ is a strong oxidizer and it may cause fire. It can cause severe skin burns and eye damage if it gets in contact with skin or any other part of your body^[5]. Therefore, wear a lab coat, use goggles and safety gloves while conducting the experiment and especially during dilution.

2.       H₂O₂ is harmful if swallowed. It can cause changes in liver weight, enzyme inhibition, induction and changes in blood or tissue levels^[5].

3.       H₂O₂ is harmful if inhaled. It can cause nausea, headache and shortness of breath^[5].

4.       Keep the H₂O₂ solution away from oxidizing agents and store in a cool and dry conditions^[5].

5.       MnO₂, KI and Fe(NO₃)₃ are harmful if consumed and inhaled. MnO₂ is a very fine powder which is black in colour and so it easy for it to enter into the respiratory system and cause harm such as inability to breathe properly. Their contact with skin and eyes can cause irritation^{[5], [8]}.

6.       Wash hands thoroughly after handling these above-mentioned chemicals

Environmental precautions

There were no environmental considerations to be taken into account. 

Ethical consideration

There were no ethical considerations to be taken into account

Signal word^{[5], [8]}

1. Hydrogen Peroxide, H₂O₂: Danger
2. Manganese (IV) dioxide, MnO₂: Warning
3. Potassium Iodide, KI: Warning
4. Ferric Nitrate, Fe(NO₃)₃: Warning

Data Collection

The below shown TABLE 1, TABLE 2 and TABLE 3 are the raw data’s of the amount of oxygen produced due to the decomposition of H₂O₂.  These data’s were obtained with the help of Vernier O₂ gas sensor and a software interface, Logger Pro.

DATA PROCESSING

The values which came under the time frame between 52 to 68 seconds were taken for calculating the rate because during this time period the rate of the reaction is constant for all the three curves. The rate can be determined using the equation:

ANALYSIS

The deviation of values from the mean rate value for MnO₂ is quite high and it results in an error of 50% and so the calculated mean value (true value) won’t make sense. The value for the 4^th trial is comparatively very low with respect to other trials. This may have happened due to any random error. The value obtained from the 4^th trial will be excluded from further calculation because it is an anomalous point.  

Observations

1.     There was a change in colour for H₂O₂ solution. Using MnO₂, the solution became black in colour. With Ferric nitrate, it gave an orange colour. Using Potassium Iodide, it gave a yellow colour (Fig.1)

2.     The initial O₂% reading during each trail was different.

3.     Towards the end of the reaction, slight moist was formed around the neck of the conical flask.

4.     The reaction of H₂O₂ with MnO₂, was almost completed at the 100^th second and there was no liberation of O₂ gas after then. Whereas, for the other catalysts, O₂ was liberating even after 180 seconds.

Best comparison graph:

As mentioned in the observation, the initial oxygen concentration was different for each trial and since I had to merge all the three reaction graphs for the comparison, I standardized the values slightly such that the three curves begin from a common point on y-axis. The least O₂% value, 18.01% was kept as the common point from which the curve evolves. For this, a factor of 0.20 was subtracted from the MnO₂ Trial 1data-set and factor of 0.21 from the KI Trial 1 data-set. The corresponding data-set obtained was used to plot the graph shown above. The red curve is of MnO₂, KI is green and Fe(NO₃)₃ is blue.

B.      Graphical Analysis:

From GRAPH 1 it can be observed that the curve produced by MnO₂ is smooth but not with the other two catalyst. This made the rate determination of MnO₂ easy by finding the tangent of the curve and difficult for other catalysts since the curves produced weren’t smooth and uniform. The rate is varying at different points on the curve, which is expected to happen unlike with the reaction using MnO₂ as catalyst where the rate at which the oxygen evolving remains constant. The graph portrays which catalyst results in higher rate and which has a lower rate. MnO₂ as the catalyst decomposes hydrogen peroxide at a higher rate compared to KI and Fe(NO₃)₃. Comparing KI and Fe(NO₃)_{3 ,} KI decomposes hydrogen peroxide at a faster rate. It is observed that the graph produced Fe(NO₃)₃ was not smooth. Initially, there was a decrease in the oxygen concentration and then an increase. This shows that the oxygen present inside the conical flask was consumed during the reaction. The decomposition of H₂O₂ using Fe(NO₃)₃ can be explained using two different reaction mechanism, ‘Kremer Stein mechanism’ and ‘Haber-Weiss mechanism’^[7].

With respect to ‘Kremer Stein mechanism’, the ferric ion creates a complex with oxygen and will have an oxidation state of +5 due to reaction between Fe³⁺ and H₂O₂. The complex will further react with a molecule of H₂O₂ in presence of water which will result in giving back (Fe³⁺) ^[7].

From this I assume that the initial decrease in the oxygen concentration is due the oxidation process of Fe³⁺ to Fe⁵⁺ which consumed the oxygen present inside the conical flask to oxidize.

Looking into this reaction from the ‘Haber-Weiss’ perspective, it is seen that the ferric ion (Fe³⁺) which initiates the radical reaction get reduced into a ferrous ion (Fe²⁺) during the reaction and again converts back to Fe³⁺ by oxidation. The change of oxidation number from iron from +3to +2 and again back to +2 can be termed as a chain reaction and it is the chain reaction that consumes the hydrogen peroxide^[7].

From this reaction mechanism I assume the decrease in the initial concentration of oxygen is due to the chain reaction which consumes the oxygen inside flask in order to oxidize Fe²⁺ and convert it into Fe³⁺.Both reaction mechanisms justify the decrease in the initial oxygen concentration. However, my experiment follows the ‘Haber-Weiss mechanism’ because (e.q.3) shows the final reaction that will happen in which the ferrous ion (Fe²⁺) is formed in the product side along with oxygen and water. Presence of ferrous ion (Fe²⁺) in a solution gives out an orange colour where as ferric ion (Fe³⁺) gives out yellow colour. I have mentioned in the observation that the colour of the solution after the reaction was orange in colour. I also assume that rate of decomposition of H₂O₂ was less compared to other catalyst due to the various steps involved in the reaction mechanism. In general, there involved three steps The reduction of Fe³⁺ ion to Fe²⁺, Fe²⁺ converting back to Fe³⁺ and Fe³⁺ ion decomposing H₂O₂ into O₂ and H₂O. The reaction between H₂O₂ and KI involves two steps^[9].

The number of steps involved in the reaction mechanism justifies the difference in the rate of decomposition of hydrogen peroxide using different catalyst. However, the physical nature of MnO₂ also helped to give out a higher rate. MnO₂ in solid for is a very fine powder. Its surface provides a proper environment for the reaction to happen more effectively.  TABLE 5 shows the rate of the reaction which is determined graphically (from best-comparison graph) by finding the tangent in the time frame between 52s and 68s.

Error calculation:

CONCLUSION

The rate of decomposition of hydrogen peroxide (0.01 mol dm^-3) using MnO₂, KI and Fe(NO₃)₃ was successfully determined through this investigation and MnO₂ came as the most effective catalyst to decompose H₂O₂ in a lesser span of time  followed by KI and Fe(NO₃)₃ in the second and third respectively.

However, the obtained result falsifies my proposed hypothesis. This shows that my assumption based on idea of oxidation state is wrong. Therefore, oxidation sate of the reacting element/ion won’t determine how efficient the catalyst is for decomposing H₂O₂. The physical nature of the catalyst and as well the number of steps involved in the reaction mechanism certainly affected the rate of the reaction. MnO₂ having the apt physical nature and less number of reaction steps, it was able to decompose H₂O₂ much faster than the other catalysts. Therefore, the physical nature and the number of steps involved in the reaction mechanism determine how efficient the catalyst is. So, it can be concluded as, 

EVALUATION

Despite of high uncertainties in the values calculated, this investigation was successful in determining the effective catalyst among the selected ones to decompose hydrogen peroxide (H₂O₂). Fe(NO₃)₃ gave consistent rate values for all the four trials which shows the precision of results. However, the values obtained using MnO₂ and KI were not precise and so the obtained mean rate values needn’t be the true value. The potential reason for this error could be, MnO2 and KI are basic compounds and H₂O₂ being both acidic and basic. Taking consideration of this case, during the reaction either neutralization reaction or hydrolysis reaction could have taken place which led to the varying rate values. With MnO₂ as catalyst, there were only three values to calculate the mean rate value and four values for the other two catalysts. This limited the accuracy of my results. Multiple trials can be conducted to reduce the random errors. One reason for which the experiment was limited to four trials is because there was only limited number of conical flasks. Even if the conical flask was reused after cleaning, there will be possibility of left-over stains around the surface of the conical flask. Proper cleaning of the flasks will consume a lot of time which may affect the concentration of H₂O₂ since it is a self-decomposing chemical. The data collection all were automated using the probe and the software interface. To some extent, this eliminates the possibility of random errors that can occur. One potential extension to this study could be determining the change in the activation energy (E_A) using various catalysts.

BIBLIOGRAPHY

[1]      R. Melina, “Why Does Hydrogen Peroxide Fizz On Cuts?,” LiveScience, 2011. [Online]. Available: https://www.livescience.com/33061-why-does-hydrogen-peroxide-fizz-on-cuts.html. [Accessed: 23-Feb-2020].

[2]      P. Pędziwiatr and A. B. Filip Mikołajczyk, Dawid Zawadzki, Kinga Mikołajczyk, “Paulina Pędziwiatr, Filip Mikołajczyk, Dawid Zawadzki, Kinga Mikołajczyk, Agnieszka Bedka,” Acta Innov., no. 26, pp. 45–52, 2018.

[3]      I. Dalmázio et al., “The iodide-catalyzed decomposition of hydrogen peroxide: Mechanistic details of an old reaction as revealed by electrospray ionization mass spectrometry monitoring,” J. Braz. Chem. Soc., vol. 19, no. 6, pp. 1105–1110, 2008.

[4]      J. Abbot and D. G. Brown, “Stabilization of iron-catalysed hydrogen peroxide decomposition by magnesium,” Can. J. Chem., vol. 68, no. 9, pp. 1537–1543, 1990.

[5]      D. Bp, J. Continental, D. Hydrochloride, G. Simone, and D. Tpo, “Safety data sheet Safety data sheet,” Carbon N. Y., vol. 1173, no. i, pp. 1–8, 2005.

[6]      S. Owen, C. Ahmed, C. Martin, and R. Woodward, Chemistry for the IB Diploma, Second Edi. Cambridge University Press, 2014.

[7]      X. It, “Decomposition of Hydrogen Peroxide by Various Catalysts,” Msds, no. Iii, pp. 2–4.

[8]      M. Dioxide, M. Details, S. Details, M. Dioxide, and O. N. Classification, “Safety Data Sheet Manganese Dioxide,” pp. 1–7, 2014.

[9]      RUGERS, “Catalytic Decomposition of Hydrogen Peroxide by Potassium Iodide.” [Online]. Available: https://chem.rutgers.edu/cldf-demos/1019-cldf-demo-elephant-toothpaste. [Accessed: 23-Feb-2020].