Zinc Oxide (ZnO)
Cuprous Oxide (CuO)
Safety Precautions
- Wear goggles.
- Make sure that the hydrogen peroxide does not make contact with skin or hair.
- Wipe bench thoroughly after experiment.
Procedure:
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Set up your apparatus as shown below measuring the hydrogen peroxide solution as being 48cm3 H2O and 2cm³ H2O2. Also make sure the burette is filled to a sensible volume; one which will make recording the volume of oxygen released less problematic.
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Next measure out 0.5g of manganese oxide. This will be added to the H202 solution.
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Add the manganese oxide to the solution and quickly plug the conical flask with the bung. The quicker this procedure is completed the smaller the probability of any O2 being lost to the atmosphere.
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In synchronisation with part 3, start the stop clock and proceed to take readings of the volume of O2 released every 15 seconds until the reaction appears to have completed.
- Reset the burette, replace the hydrogen peroxide solution and run through parts 2 to 4 again only this time using one of the other metal oxides.
Results.
Conclusion:
From the above results it is clear that the only oxide of the ones tested to show any appreciable increase in volume of O2 produced is Manganese Dioxide. The other two oxides tested do not show any change at all and can therefore be ruled out as being catalysts to the decomposition of hydrogen peroxide.
Discussion:
All of the substances tested share a likeness in that they are all oxides, they do not however all act the same when put with hydrogen peroxide. Only the manganese dioxide increased the rate of the reaction: 2H2O2 → 2H2O + O2.
The manganese dioxide has therefore acted as a catalyst and as with all catalysts has lowered the activation energy required in order for this reaction to take place. In doing so the rate of reaction has increased.
As you can see in the diagrams the level of energy required for reactants to become products is greatly reduced in the presence of a catalyst.
Although the rate of reaction has been increased the total volume of O2 produced, (were the H2O2 left to fully decompose without the catalyst) is not increased only gained in a shorted period of time. So a catalyst therefore does not create something that would not form anyway it just makes the reaction happen faster. In certain circumstances a reaction would not happen at all were it not for the presence of a catalyst.
The way a catalyst works is by forming covalent bonds with the reactants, which in turn breaks the bonds between the reactants. The bonds between reactant and catalyst then begin to break as new bonds are formed and the subsequent product is created. The diagram below shows this:
A catalyst therefore will not be affected by the reaction it catalyses and can be used (in theory) an infinite number of times unchanged. Another way to have confirmed the Manganese Dioxide as a catalyst would have been to weigh it after the reaction had finished. If the Manganese Dioxide were a catalyst then there would be no loss of mass. We could also have conducted the experiment a second time using Manganese dioxide recovered from the first experiment and measuring the volume of O2 released, the result should be very similar if not identical. The practicalities of removing the catalyst from the liquid in this case would be extensive. A way around this could be to obtain a figure for the density of O2 and using the figures recorded for the O2 released calculate the weight of the O2 released. If the total mass (conical flask, water, H2O2, catalyst and bung) is recorded at the beginning and then again following the experiment, and if the weight of the O2 is subtracted from the mass recorded initially then there should be no loss of mass. This in itself would prove complicated and without extremely accurate apparatus and the assurance that no O2 was lost during the experiment we could not prove whether or not any loss of mass was down to the “catalyst” or just inferior practice.
