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

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Introduction

The Decomposition of Hydrogen peroxide Aim: to investigate the effect of concentration on the rate at which hydrogen peroxide decomposes. Prediction I predict that as the concentration of the hydrogen peroxide is increased to water the rate of reaction will increase. This is because the concentration leads to more collisions between particles as is shown in the hypothesis. Furthermore I also predict that there will be a linear increase in the rate of reaction when I increase the concentration. This is because when you double the concentration you double the number of particles in the same volume. If the number of particles with the activation energy is in the same proportion, you will double the number of particles with the activation energy. I also predict that this reaction will be first order with respect to the hydrogen peroxide. This is because it is a decomposition reaction and most decomposition reactants are first order, so I predict that the rate equation will be this: Rate = k[H2O2] Hypothesis/Background knowledge Chemical reactions can proceed at different speeds, this is called the "rate of reaction". The rate of reaction means how fast or slow the reaction is going. Another way of representing this is how quickly the following occurs. ...read more.

Middle

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. ...read more.

Conclusion

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. ...read more.

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