Rate of Decomposition of Hydrogen Peroxide - Lab Report Example

The rate of decomposition of hydrogen peroxide at room temperature is so small that measurement cannot be performed. Therefore, hydrogen peroxide appears to exhibit a metastable characteristic in solution or its pure state (Egan and Nills, 2005). Oxygen evolution is rapid at room temperature and in concentrated solutions, the released heat may increase the temperature so that an explosion occurs. On heating, hydrogen peroxide decomposes and it may be explosive. The stability of hydrogen peroxide at room temperature is attributed to the fact that the first step in its thermolysis involves the splitting of the molecule into hydroxide radicals whose formation demands much heat. Catalysts such as silver, gold, platinum, manganese dioxide act as heterogeneous catalysts whereas ions such as I-, IO-, OH-,Fe3+, or copper act as homogenous catalysts.

            Catalysis of hydrogen peroxide decomposition by iron ions is important in redox catalysis. The steps in the process as demonstrated by Evgenil, Oleg, and Gerts (2005), give data on the mechanisms of redox catalysis. This decomposition is also important in processes in living organisms. The decomposition may be represented as below;

2H2O2   (l) → 2H2O(l) + O2(g)

Method

Apparatus

            200ml Dewar flask, Standard potassium permanganate solution, 2M sulphuric acid solution, Fe3+ solution catalyst, Pipette.

Procedure

            25 mL of the ten‐volume hydrogen peroxide was diluted to about 200 ml with distilled water. A Dewar flask was rinsed with distilled water and then with a few mL of the dilute H2O2 solution. Hydrogen peroxide solution was added into the Dewar flask, and the solution temperature was noted at a constant value. 10 ml of Fe3+(a) catalyst solution was added into the beaker while stirring gently. A stopwatch was used to keep track of time.

Thirty seconds after the addition of the catalyst, a 10 mL aliquot of the reaction solution was pipetted and transferred into an Erlenmeyer flask containing 18 mL of 2 M H2SO4. The acid provides an acid medium for the titration and also quenches the decomposition reaction. 10 mL samples were drawn from the reaction mixture at 5 min intervals and titrated with potassium permanganate. The titration reaction follows the equation;

2MnO4 (aq) + 5H2O2 (aq) + 6H+(aq) → 2Mn2+(aq) + 5O2(g) + 8H2O(l)

            The volumes of potassium permanganate were noted and recorded as per table 1(Laboratory Manual Semester 1, 2014).

Results

Table 1 Results from the titration

mins

Titre volume(Vt)

ln Vt

[A]

ln[A]

0

24.1

3.18221

1.24481

0.21899

5

36.3

3.59182

0.82645

-0.1906

10

45.8

3.82428

0.65502

-0.4231

15

6.6

1.88707

4.54545

1.51413

20

10.9

2.38876

2.75229

1.01243

25

13.6

2.61007

2.20588

0.79113

30

16

2.77259

1.875

0.62861

35

18

2.89037

1.66667

0.51083

40

19.8

2.98568

1.51515

0.41552

45

21

3.04452

1.42857

0.35667

Discussions

From the results obtained in table 1, the plot of the volume Vt versus the time t is a smooth curve from 15minutes to 45minutes (Egan and Nills, 2005).

Figure 1 The plot of Vt against time t

Reaction rate at 5 minutes

Reaction at 25 minutes

Reaction rate at 45minutes

                    The reaction rate increased with time but eventually remained constant during the decomposition. This can be attributed to the fact that the rate of decomposition of hydrogen peroxide and the removal of its by-products was efficient therefore the reaction mixture maintained a constant reaction rate

Figure 2 The plot of [H2O2] against time(minutes

Figure 3 Plot of ln Vt against t (mins)

Rate Constant k= = (3-2.5)15 = 0.03333min-1

            From the results obtained the plot of ln Vt against time t is linear. This conforms to the expected first-order reaction of graphs. The iron catalyst does not initiate the decomposition of hydrogen peroxide. Hydrogen peroxide decomposes even in the absence of the Fe3+ through the amounts of O2 liberated is negligible for measurement. The Fe3+ provides adsorption sites thereby increasing the rate of reaction (Evgeny T. Denisov, Igor B. Afanas'ev).

            The temperature has a high influence on the rate of reaction. Depending on whether the reaction is exothermic or endothermic the temperature will either increase or decrease the rate of reaction respectively. In the case of hydrogen peroxide, an increase in temperature will increase the rate of reaction. Egan and Nills, (2005), explain that at room temperature the thermolysis of H2O2 takes place at high temperatures as it follows the formation of hydroxide radicals.

 Conclusion

            The rate of reaction can be estimated using an elaborate laboratory experiment. In this experiment, by measuring the concentration of hydrogen peroxide at specific time intervals it has been possible to find the rate of reaction and the reaction constant. Also, graphs that comply with the first-order reaction have been obtained.