Aim: Various metal oxides will be added to the hydrogen peroxide and the production of oxygen of the reaction mixture will determine catalysis.
The volume of oxygen evolved will be observed and recorded to measure the reaction rate and the reaction rates of the different metal oxides will be compared.Scientific Background: Catalysis is the process by which the activation energy is lowered to allow the reaction to occur at less extreme conditions, during the process the catalyst does not under go any overall change. The catalyst reduces the activation energy by using a chemical route with activation energy less then the route, which would otherwise be taken in the absence of the catalyst. During catalysis the reacting substance usually undergoes a change or changes in oxidation state, therefore the catalyst must also be able to change its oxidation state. The s & p block metals possess or exhibit only one oxidation state. The reason being that their oxidation state depends on the removal of electrons from their outermost shell.The further removal of electrons will result in the penetration of stable inner shells that are filled with electrons. This would require an excessive amount of energy.
As a result of this, the catalysts are not able to enter their different oxidation states and this therefore does not allow them to successfully act like a catalyst. Transition metals can form ions, which have D orbitals, which are partially filled with electrons. D orbitals are the outer most shells and they can hold up to 10 electrons and they also have similar energy levels, which allows them to overlap within each other. It is this process which allows the transition metals to have various oxidation states. The orbital possess the same energy, which enables transition metal ions to enter their different oxidation states.
This therefore allows them to act as catalysts. As there are more spaces for electrons to be lost and gained the reaction can take place faster and better.The general equation for the experiment is:CATALYST2H2O2 O2+2H20From the equation it can be seen that Oxygen is produced in the reaction and this is what is being collected and measured in the gas syringe.Prediction: The transition metal oxides MnO2 (Manganese oxide), ZnO (Zinc Oxide) and CuO (Copper Oxide) will be compared with SiO2 (Silicon Oxide), A12O3 (Aluminium Oxide) and PbO (Lead Oxide) which are transitional metal oxides to see their action on hydrogen peroxide and therefore see if only the transition metal oxide allow catalysis or that all metal oxides allow catalysis for this reaction.Apparatus:In the experiment, 20cm3 of water and 20cm3 of hydrogen peroxide will be used.The above apparatus is easy to use and will give quick results; a stop clock will be used to see how much gas is collected in a given time.
The conical flask is used to keep the solution in and the rubber bung is placed on top firmly to prevent any loss of gas, which would obviously affect results.The following oxides will be chosen for the experiment: SiO2, A12O3, PbO, MnO2, ZnO and CuO. 1 gram of each oxide shall be usedSiO2, A12O3 and PbO are not transition metal oxide, whereas MnO2, ZnO and CuO are transition metal oxide. They will be used to compare the catalytic action if any of transition and non-transition metal oxide on hydrogen peroxide.1 gram of each oxide allows a sensible amount of oxide to be used without causing a dangerously vigorous reaction to occur.
Variables:The following variables will be controlled throughout the experiment:Temperature can affect the reaction rate of H2O2 decomposition. Two molecules can only react if they have enough energy, by heating a mixture the energy levels will be raised of the molecules involved in the reaction. Therefore, by increasing the temperature the molecules will move faster and collide with each other more quickly. This will increase the rate of reaction and so the temperature will be kept constant throughout the experiment by carrying it out under room temperature.Concentration and volumes of the hydrogen peroxide must be kept constant because if there are more molecules of a particular substance in a certain volume then there is more of a chance of the molecules colliding with each other. The frequency of collisions is increased which increases the rate of reaction. Thus the hydrogen peroxide must be kept constant for each run of the experiment with a metal oxide. The amount of oxide must also be kept constant for the same reasons.
Particulate size could affect the rate of reaction as if the particle sizes are larger then the rate of reaction will be slower because there is a smaller surface area for the substance to react on. By crushing the catalysts to have the same particle sizes this can be controlled.The key variable is the one, which will be the one, which is varied in the experiment. In this case it will be the metal oxide used in the experiment, which will be varied.Preliminary Experiment:A preliminary experiment was carried out using the same apparatus and the catalysts used were manganese oxide, iron oxide and lead (iv) oxide. The manganese oxide worked well as a catalyst but no results could be collected as the reaction went to quickly. This is most probably due to the amount of oxide being too much. No results were collected for iron oxide, and lead (iv) oxide because they did not produce any gas and therefore did not work as catalysts.
The apparatus seemed appropriate however, a clamp stand was not used to hold the syringe up and so the apparatus can be modified so that the gas syringe is supported by the clamp stand. The other modifications, which need to be made, are the catalysts being used and amount of oxide. The amount of oxide used in the preliminary experiment was 2 grams and so this can be reduced to 1 gram to decrease the rate of the reaction so that a good set of results can be collected.Method:1 – Measure out using a measuring cylinder 20cm3 water and 20cm3 hydrogen peroxide and add to a conical flask.2- Using the balance to weigh out accurately 1 gram of MnO2.3- Using the filter paper used to weigh the oxide add the oxide to the hydrogen peroxide, and immediately stopper the flask with the connecting tube attached to the bung.4- Immediately start the stop clock and see how much gas is collected the gas evolved will push the syringe itself; therefore do not touch the syringe.
Readings should be taken every 10 seconds.5- Record how much gas if any is collected until the syringe stops moving itself.6- Repeat the experiment with MnO2 so three sets of results can be collected to find an average.7- Do 1-6 with the remaining oxides and tabulate the results.Safety:* Hydrogen peroxide is a corrosive liquid and should therefore be used sensibly and eye and skin contact should be prevented.* Manganese IV Oxide is harmful and again skin contact should be avoided.* An overall should be worn when handling hydrogen peroxide.* Goggles should be worn when adding the oxide to the peroxide, although they are not needed throughout the rest of the experiment.
Reliability of Results:* When collecting the gas, the attachment of the rubber bung should be very quick to minimise any loss of gas.* The hydrogen peroxide should be measured out accurately. The oxides should also be weighed accurately using an accurate balance.* The oxides should all be in solid powdered form and not in any other form, as this will effect how the catalysts react with the hydrogen peroxide. If the surface area is different on the catalysts then they will react more/less violently.* The apparatus should be cleaned after each run with water to remove any impurities that may affect the next experiment.* The timing of each experiment should be accurate and a digital stop clock should be used.
All theses steps should ensure that reliablePrecision and Accuracy:* The balance is a precision instrument with an accuracy of 0.01 grams. This is sufficient for the task at hand.
* The syringe is accurate to 1cm3, however it can be seen when the syringe marker is halfway through the points so the results can be recorded to an accuracy of 0.5cm3.Results: The following tables show the amount of oxygen evolved from each catalyst and the three readings taken are also listed.Time from start of exp. (s)Volume of gas produced by MnO2 (cm3)Reading 1Reading 2Reading 3Average109.56.
02001009590.595.2The line graph on the next page shows the average amount of oxygen produced by manganese oxide against time in seconds. Adding up all the readings and dividing by all the readings taken i.e.
3 have found the average.The following table shows the amount of oxygen evolved in a certain amount of time from the catalyst PbO.Time from start of exp. (s)Volume of gas produced by PbO (cm3)Reading 1Reading 2Reading 3Average104554.72055.55.55.
0On the next page is a graph with a line of best fit representing the results in the last column above, which represents the average results.The following table shows the amount of oxygen evolved in a certain amount of time by the catalyst Al2O3.Time from start of exp. (s)Volume of oxygen produced by Al2O3 (cm3)Reading 1Reading 2Reading 3Average103232.7203232.
220055.555.2On the following page the graph for the results above has been included. Only the average results are shown.The following table shows how much gas is evolved from the catalyst Copper Oxide (CuO):Time from start of exp. (s)Volume of oxygen produced by CuO (cm3)Reading 1Reading 2Reading 3Average102322.3203433.3304544.
0180101169.0190101169.0200101169.0The graph on the following page shows the average amount of gas produced by Copper oxide in a certain amount of time.The following table shows the amount of oxygen produced by Silicon Oxide (SiO2) on every run of the experiment.Time from start of exp. (s)Volume of gas produced by SiO2Reading 1Reading 2Reading 3Average101121.3202232.
54.3The graph on the next page shows the average amount of gas produced by Silicon Oxide.The table shows the amount of oxygen evolved from the hydrogen peroxide solution and the Zinc oxide catalyst.Time from start of exp. (s)Volume of gas produced by ZnO (cm3)Reading 1Reading 2Reading 3Average106444.
019016121414.020016121414.0On the next page is a graph to show the average amount of oxygen evolved from zinc oxide and hydrogen peroxide solution.Analysis:It can be seen that there is a pattern in the results and this pattern indicates that as the reaction starts more and more gas is produced but as it continues the reaction begins to slow down and not as much gas is produced as the volume of reactants is decreasing. The reaction finishes when no more oxygen is produced. There is a directly proportional relationship between the time and the volume of gas produced; as time increases more gas is produced however, this is only up to a certain point, when the reaction is running at its highest potential it can no longer produce oxygen.
Giving the example of manganese oxide, at 20 seconds it was at 15cm3 and then at 70 seconds it was 49cm3. This re-enforces the conclusion gathered above. However, from the graphs of the other catalysts, it can be seen that they differ from manganese oxide.Each graph has to be analysed individually as they each follow a slightly different pattern. Below each graph has been analysed to further show what they allude to.Graph 1 shows the amount of oxygen evolved with the transition metal catalyst manganese oxide. It can be noticed instantly that as times increases the volume of gas being produced increases however, the gradient of the curve decreases i.e.
the volume of oxygen being evolved rapidly decreases. It can be seen that at the beginning the line of best fit is steepest and so this shows that the reaction is running at its best towards the beginning. It can be seen that more gas is being produced at the beginning, if one looks between 20 and 30 seconds the difference between the amounts of gas evolved is 6.8cm3 and between 180 and 190 seconds the difference is 3.4. This shows that the rate of reaction is decreasing because the catalyst can no longer produce anymore oxygen and the reaction is running at its highest.
The maximum amount of gas evolved from manganese oxide is 95.5cm3.Graph 2 shows the amount of oxygen evolved with the metal oxide Lead oxide. It can be seen that at first the amount of gas is increasing by a large amount each time but as time goes on the difference in the amount of gas produced is decreasing. This shows that the reaction is slowing down and the reactants have been used up in the reaction so there is no more gas being produced.Graph 3 shows the amount of oxygen evolved with the metal oxide aluminium oxide.
It can be seen that initially there is a steep increase in the volume of gas evolved over the first 10 seconds. There is then a gradual increase over the next 80 seconds after which point the graph begins to plateau. After 160 seconds no more gas is evolved from the reaction as it has produced as much as it can. The reaction has ended as all the reactants have reacted together.Graph 4 shows the amount of oxygen evolved with the transition metal catalyst copper oxide. It can be seen that there is a steady increase in gas that is being evolved. There are very few anomalous results on the graph and they have been circled.
They could have occurred due to misreading the gas syringe and recording incorrect results.Graph 5 shows the amount of oxygen evolved with the metal oxide silicon oxide. The gas evolves quite slowly and the reaction completes in 60 seconds.Graph 6 shows the amount of oxygen evolved with the transition metal catalyst zinc oxide. The reaction starts off very quickly and completes in between 130-140 seconds.As the prediction made was to compare the transition metal oxides to metal oxides it could be seen from the graphs that MnO2, CuO and ZnO work best as catalysts as they produce the most amount of gas. The other catalysts did not work as well and did not produce very good results. The reason for the transition metals to work so well is because transition metal atoms have the space and energy to allow the catalysts to have variable oxidation states whereas the other metal oxides are not capable of doing this as they do not have the same set of sub-shells as the transition metal oxides.
The transition metal oxides have d-orbitals, which is a sub-shell with similar levels of energy, which allows the electrons to overlap and jump from being in one oxidation state to another. However, the metal oxides do not have d-orbitlas but just S and P sub-shells.Evaluation:The experimental techniques that were used can be said to be suitable for the task at hand. It can be said at with some catalysts the techniques may have been unsuitable however the reason for them being suitable are:1. The results obtained for some of the catalysts shows clear trend and patterns eluding to the fact that the procedure used was suitable.2. The trends and patterns identified directly correlate with the hypothesis at the beginning and the outcome of the procedure expected.3.
The procedure allowed the volume of oxygen evolved to be measured.The reliability of results was high as the difference between repeated readings was small which re-enforces the reliability of the results. The procedure allowed all variables to be controlled, the temperature and concentration were maintained well and the particulate size was maintained by crushing the catalysts to the same size.
The apparatus allowed precise results to be collected. Evidence for this is that the difference between repeated readings is small and the amount of anomalous results is small thus the number of errors must be small.The plan and method that was used for this experiment can be said to be very good. The reason being that both fulfilled their intended tasks. Before the investigation, it was desired to find how different metal catalysts affect the rate of reaction and to draw conclusions, which did or did not cohere to the prediction made.
This was achieved as it could be seen whether transition metal catalysts and metal catalysts had an affect on the rate of reaction.As in any investigation, errors were present despite the fact that prior to the investigation all such errors were minimised. Below the errors and problems encountered are discussed along with the improvements that could be made to avoid these problems:1. The volume of gas on the syringe could have been misread which would lead to inaccuracy of results. This could be one reason for why some anomalous results would have appeared. This can be avoided by2. The stopper was not being placed on the conical flask quick enough and so much gas was lost during this time.
By the time the stopper was placed onto the conical flask, the reaction was coming to an end. To prevent further loss of gas in future experiments the following could be done:* Another individual could be asked to place the stopper on, however in this process some gas will still be lost.* A divider flask could be used which is a conical flask that is split into two sections.
If the catalyst is placed in one section and the solution is another the divider flask will allow the reactants to be separate for long enough so that the stopper can be placed on. Once this has been done, the conical flask can be tipped over slightly to let the reactants mix.