Manganese (IV) Oxide MnO2
Hypothesis
I predict that the rate of reaction will increase with the greater the volume of the catalyst used until it comes to a certain limit when after that limit the rate of reaction will stay the same. I think this because the collision theory states that if there are a greater number of molecules in a substance or a greater concentration of a solute then the rate of reaction will be faster, so with a greater mass of the catalyst there are more atoms so it should react faster.
Scientific Background
A catalyst is very important in this investigation. A catalyst is a substance that speeds up the rate of a reaction without getting involved with the reaction.
The collision theory is also vital to this investigation. The collision theory says that a chemical reaction can only occur between particles when they collide (particles may be atoms, ions or molecules). There is a minimum amount of energy which colliding particles need in order to react with each other. If the colliding particles have less that this minimum energy, then they just bounce off each other and no reaction occurs. This minimum energy is called the Activation Energy.
There are four things that can increase the likelihood of a collision:
- TEMPERATURE - The faster the atoms move the more likely they are to collide with each other. Speed depends on energy and energy depends on temperature.
In a cold reaction mixture the particles are moving quite slowly – the particles will collide with each other less often, with less energy, and less collisions will be successful.
If we heat the reaction mixture the particles will move more quickly – the particles will collide with each other more often, with greater energy, and many more collisions will be successful.
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The more atoms there are, the more likely they are to collide i.e. PRESSURE of a gas or CONCENTRATION of a solute.
A dilute solution is one in which only a small amount of solute particles are dissolved. A concentrated solution is one which lots of solute particles are dissolved.
The more concentrated a solution is, the closer together the particles are. This makes them more likely too bump into one another resulting in faster reactions.
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The better they are mixed the more likely they are to collide i.e. SURFACE AREA affects rate.
Large particles have a small surface area, in relation to their volume, and so react more slowly.
Small particles have a large surface area, in relation to their volume, and so react more quickly.
A large surface area means that more particles are exposed and available to collisions – this means more collisions and so a faster reaction.
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A final way of changing the rate of a chemical reaction is by adding a CATALYST. A catalyst is a substance, which speeds up a chemical reaction without being involved in the reaction.
A catalyst works by lowering the amount of energy needed for a successful collision (activation energy) – so more collisions will be successful and the reaction will be faster. Also it provides a surface for the molecules to attach to, thereby increasing their chances of bumping into each other.
Most reactions are exothermic i.e. the chemicals involved are going to a lower energy state, but initially they need energy to start the reaction. This can be shown on an energy profile:
We can increase the number of molecules able to get over the barrier in 3 ways:
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Give the molecules more energy → temperature
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Lower the barrier by adding a catalyst
- Some reactions require a specific trigger such as a spark, a physical shock or a particular form of radiation.
Preliminary Work
Apparatus
Side-Arm Conical Flask (250ml),
Gas Syringe (100cm3),
Glass Stopper,
Clamp and Stand,
Weighing Scales (accuracy 0.01g),
Stop Clock (to 1 sec),
Rubber tubing,
Different catalysts (Copper Oxide, Iron Oxide, Manganese Oxide),
Hydrogen Peroxide (H2O2),
Measuring Cylinder (25ml),
Teat Pipette.
Method
- Set up the apparatus as shown in the diagram
- First measure out 0.30g of the first catalyst, which is copper oxide and place it into the side-arm conical flask.
- Then measure out 20ml of the hydrogen peroxide with the measuring cylinder.
- Before pouring the hydrogen peroxide into the side-arm conical flask, make sure the stop clock is ready.
- Pour the hydrogen peroxide into the flask taking care not to spill any onto the side of the flask. As soon as it is all in, put the rubber bung on the flask and start the stop clock.
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Every five seconds, check how many cm3 Oxygen has entered the gas syringe and record it on a chart up to sixty seconds.
- Repeat this experiment again so you can create an average out of the two sets of results.
- Then repeat the experiment with 0.30g of Iron Oxide and 0.30g of Manganese Oxide doing each experiment twice.
Results
The results to the preliminary work were as follows:
After the preliminary work we decided that we should do the second part of the investigation with Manganese Oxide. We decided to change the mass of the Manganese Oxide in 0.05g steps starting at 0.05g and ending at 0.30g. We also decided to change the amount of Hydrogen Peroxide in the investigation down to 10ml from 20ml because there was too much Oxygen being produced to be able to measure the amount accurately every five seconds.
Fair Testing
In this investigation I will keep constant the following to make a fair test:
- The volume of Hydrogen Peroxide in each experiment.
- The mass of catalyst used in the preliminary work but this will be the independent variable in the second part to the investigation.
- The same equipment as there may be minor changes in the size of each piece of equipment or the readings on the side of the equipment.
- Use the same method for each experiment so there will not be any major differences. Only change the mass of catalyst used in the second part of experiment.
- Measure temperature in the room with a thermometer.
- Ensure that the side-arm conical flask is washed out properly every time you finish an experiment to stop contamination.
- Use the substances from the same bottle because if you use a different bottle the concentration maybe slightly different which could affect your results.
Obtaining Evidence
Safety
Whilst this experiment is relatively safe, there are still some safety considerations to be adhered to:
- We wore eye goggles when we used Hydrogen Peroxide in case some of it splashed up into our eyes because it is dangerous and can cause burns.
- We had to be careful that if any Hydrogen Peroxide got onto our hands we had to wash it off straight away with water because it is corrosive.
- We had to be careful how we held the gas syringe because there was a danger that the sliding bit could slip out. We also had to make sure that it did not fall out of the end of the syringe when the experiment was going on. We also had to be careful with all equipment that was made out of glass in case it breaks.
Results
Analysing Information
The table above shows that the most Oxygen was produced with 0.3 grams of Manganese Oxide. It also shows that the amount of Oxygen produced increased gradually with the mass of the Manganese Oxide.
With these results we were able to work out the rate of reaction for each set of results:
Rate of Reaction = volume of gas
Time
The rate of reaction also increased with an increase in the mass of Manganese Oxide. The rate of reaction gradually increased with each different mass of the catalyst except from 0.15g to 0.20g where there is a large jump in the results. Also the rate of reaction goes down from 4.5cm3/s to 4.0cm3/s from 0.2g to 0.25g. This is a mistake in the results and could have occurred in many ways. Probably the most likely answer is that we were slow in putting the bung on the conical flask so some Oxygen might have escaped out of the top. Another possible answer could be that we made mistakes in measuring the amount of Oxygen given out at exactly every five seconds.
Most of the results are as we expected from the beginning in that the rate of reaction will increase with the more mass of Manganese Oxide. This is because there is more surface area of the catalyst for the Hydrogen peroxide to react with as said in the collision theory.
Evaluating Information
The experimental procedure is simple, but effective. It worked very well but we could have made a few improvements in the procedure. One improvement we could have made was that we could have some sort of electrical device which would measure the amount of Oxygen given off and display it every five seconds digitally. Another problem we were having is that it was difficult to get exactly the right mass of Manganese Oxide every time. Another thing is that sometimes it would take quite a lot of time to pour the Hydrogen Peroxide into the conical flask and then by the time we put the bung on, the reaction had started and we would loose some Oxygen through the top. This is important because the results at the beginning are the most important to work out the rate of reaction.
All together our results were very successful but we did have on anomalous result with 0.25g of Manganese Oxide, which was slower than the 0.20g result. The reasons for this I have already discussed above.
We could have got better results by repeating the experiments maybe a few more times so we could get lots of results and then get a better average from those results. This would give us more accurate results and maybe not have the anomalous result.
We could also have extended the preliminary work by trying out more catalysts to see if we could have got a better one than Manganese Oxide.
Bibliography
“Chemistry For You” by Lawrie Ryan
“The Essentials of GCSE Chemistry” by Mary James
www.gcsechemistry.com
