The Decomposition of Hydrogen peroxide - investigation in to the effect of concentration on the rate at which hydrogen peroxide decomposes.

• Concentration – Increasing the concentration of a reactant also increases the speed of a reaction because there are more particles in the same volume.

                Lower concentration                                higher concentration

The more crowded the particles the more collisions will occur( and therefore the more effective collisions) and the faster the speed of the reaction.

• Surface Area – The more the surface area, the more particles available to collide and react with one another, therefore the more effective collisions and consequently the faster the speed of the reaction. Smaller pieces of solids, especially powders would react faster than larger pieces, this is due to the smaller particles having a larger surface area compared to the bigger pieces.
• Catalysis – Catalysts are substances which, either man made or natural, speed up a reaction without itself being used up itself in the reaction. It works by giving the reacting particles a surface to stick to where they can bump into each other. Therefore increasing the number of collisions. They lower the activation energy for the reaction so therefore increase the speed.

Hydrogen peroxide is found in hair dyes and bleaches and is found on the earth naturally. Hydrogen peroxide (H2O2) decomposes naturally at a very slow rate (weeks), this is due to its instability so reacts with the air slowly.

H2O2                                H2O     +  O2

         2H2O2(l)                                     2H2O(l)   +  O2(g)

I think that as the concentration of hydrogen peroxide to water is increased the rate of reaction will increase as there will be more acid particles in the same volume which are more reactive and therefore will react quicker when they collide with the manganese dioxide (catalyst). The catalyst will not only be used to speed up the slow reaction to a measurable rate, but the concentration will determine the actual rate of reaction. When the peroxide concentration is increased more collisions will take place as more acid particles will react quicker concentration will determine the rate of the reaction. When the acid concentration is increased more collisions will take place as more acid particles will react and as they are more reactive than water it will mean the reaction rate will increase.

There will be an increase in the rate of oxygen produced when the rate of reaction is faster therefore when the concentration of acid is higher. This is again to do with the more reactive particles colliding. I predict that the graph produced from my results will become steeper as the concentration of acid to water is increased. This is on the next page.

Rate of Reaction

                

Concentration

The equation for this experiment will be:

2H2O2                                2H2O  +   O2

This is because the reaction is a decomposition equation as opposed to a regular reaction equation.

Plan

I plan to find the rate of reaction by doing an experiment to see how long it would take for the hydrogen peroxide to produce 50ml of oxygen, when 1.00 gram of catalyst (manganese dioxide) was also used to help speed it up. I will be looking at how the different levels of concentration will effect the rate. To do this I will be measuring the rate at which  oxygen is produced by the reaction at different concentrations of H2O2. To then go on to work out the rate I will use the following formula:

Rate of production        =        Volume of O2 (dm3)

      of Oxygen                            Time taken (s)

I will plot the graphs of the following once I have taken my measurements and processed the results.:

1. Concentration and Time
2. Concentration and 1/Time
3. Concentration and rate of reaction

Equipment List::

• Hydrogen peroxide at different concentrations
• Conical flask with bung
• Manganese dioxide (catalyst)
• Glass syringe
• Top pan balance
• Measuring cylinder
• Timer
• Water
• Tub
• Spatula
• Funnel
• Tissue

Method:

1. Collect all equipment and set up glass syringe and conical flask with bung as shown in diagram below.
2. Measure out 1.00 grams of manganese dioxide using the top pan balance, spatula and tub.
3. Repeat step 2 for 4 more tubs of manganese dioxide.
4. Using the measuring cylinder measure out 50ml of hydrogen peroxide and use the funnel to stop any spillage.
5. Pour  the acid into the conical flask.
6. Whilst pouring in the manganese dioxide start the timer and close the bung tightly.
7. Time how long it takes for it to produce 50ml of oxygen then stop the timer and remove the bung.
8. Record results.
9. Using the measuring cylinder measure out 40ml of hydrogen peroxide and top the remaining 10ml with water.
10. Repeat steps 5-8.
11. Using the measuring cylinder measure out 30ml of hydrogen peroxide and top with 20ml water.
12. Repeat steps 5-8.
13. Using the measuring cylinder measure out 20ml of hydrogen peroxide and 30ml of  water.
14. Again repeat steps 5-8.
15. Using the measuring cylinder measure out 10ml hydrogen peroxide and 40ml water.
16. Repeat steps 5-8.
17. Now repeat steps 1-16 two more times to get accurate results

Fair Test

In my experiment the independent or control variable will be the concentration of the Hydrogen peroxide; in which the catalyst (Hydrogen Peroxide) will be put into. The dependant variable will be

• To make it a fair test I kept the following features identical:

• The volume of the Hydrogen Peroxide {50ml}
• The weight of the catalyst (Manganese Dioxide) when it was put into the Hydrogen Peroxide.
• The type of Manganese Dioxide in the experiment (powder)
•  The volume of gas which the reaction has to produce (50cm3)

Sampling

I will do this experiment three times. I will however for my results calculate the average for all three experiments.  I will do this experiment once for the preliminary and a further three times for the investigation.

Conclusion:

These results generally  support my original prediction that as the concentration of hydrogen peroxide is increased, the rate of reaction also increases. They also show that overall there is an approximate  doubling of rate for a doubling of concentration

For example at 0.2 moldmˉ³ of H2O2 the rate was 1.27cm³/s and at 0.4 it was 3.24 cm³/s. A similar relationship is shown in comparing 0.4 mol dmˉ³ and 0.8 mol dmˉ³. I can therefore deduce the rate equation for the reaction to be as follows;

                                Rate α [H2O2]

This relationship between rate and concentration is characteristic of a first order reaction and from this I can deduce that the decomposition of hydrogen peroxide is a first order reaction with respect to the Hydrogen peroxide.Since it is the only reactant then the reaction is also a first reaction overall. So…

                        Rate = k[H2O2]

From  this and my results I am also able to calculate the rate constant for this reaction (since all the measurements were taken at the same temperature – it will not change).

For example.  At 0.4 moldmˉ³                     rate = 3.24cm³/s

                                        [H2O2] = 0.4 moldmˉ³

                Rate = k[H2O2]

So:            3.24 cm/s = k[0.4mol]

        G            k  = 0.00324 dm³

                        0.4 moldmˉ³

Therefore:             k  = 0.08 dm6molˉ¹

Evaluation

There were a number of  possible sources of error in the investigation as well as ways in which the it could be improved and extended. There are also some anomalous results which are interesting.

Errors may have entered through stages in the procedure as follows;

• weighing out of the manganese dioxide – the mass of the catalyst would have made a difference since above 1g and the rate would appear too high and below 1g and the rate would appear too low.
• Measuring out of 50cm³ of the hydrogen peroxide was done crudely using a measuring cylinder – it would have been better to use a graduated pipette. This is because if the amounts are not exactly what we think they are it can have an affect on the final concentration after dilution.
• Diluting of the hydrogen peroxide was done again rather crudely using measuring cylinders with the appropriate volume of water. It would have been better to use volumetric flasks to produce diluted concentrations of the hydrogen peroxide.
• The act of pouring in the peroxide and then putting the stopper on the flask would have introduced an element of error since it is very difficult to do it in exactly the same time for each reading. This would have had an effect on the time reading for 50cm³ of gas to be collected.
• It is possible that the gas syringe itself after a few measurements starts to “stick|” a little bit so the  plunger does not move  equally smoothly each time.

The results are generally consistent between the sets of readings which  have been averaged out. However, all the measurements for the highest concentration (1.0mol/dm³) appear anomalously high and do not fit with the line of best fit which I have attempted to draw. Either we can explain this in terms of a large scale error due to one or more of the sources listed above  OR  it may  be something to do with the reaction itself.

Obviously the rate appears faster at the highest concentrations than it should do according to my prediction. It might be that since this is the only one which has no added water = the presence of water might be having an inhibiting effect on the reaction so that the more there is the slower than expected the rate is. I know from my knowledge of the equilibrium of other reactions we have studied ( e.g. The Haber process) that the presence of the product can cause the  forward reaction to be inhibited so since water is a product of this reaction then it could be a possibility here.  This would explain the apparent anomaly.

To improve the  quality of this investigation I would

• Use graduated pipettes to measure volumes
• Use volumetric flasks to carry out dilutions of the hydrogen peroxide
• Use  a  better system for delivering the hydrogen peroxide to the manganese dioxide ( using a thistle funnel)
• Find out the exact concentration of the hydrogen peroxide by titration experiments to make the results more reliable.
• Repeat the measurements a larger number of times

To extend this investigation I would like to  follow up on my idea about the reason for the high rate at the higher concentration. I would use a wider range of concentrations both below the lowest we used and above the highest e.g. 0.01M to 0.1M and 1.0M to 3.0M or 4.0M hydrogen peroxide.