Introduction In this investigation I will firstly be finding out which catalyst out of the three available breaks down hydrogen peroxide the quickest. Then with the catalyst that works fastest I will use it to see if a change in volume of the catalyst which change the rate of hydrogen peroxide breaking down. The equation for the decomposition of hydrogen peroxide is: 2H2O2 (aq) 2H2O (l) + O2 (g) (It is worth noting that the compounds used are catalyst and therefore they are not used up, nor do they play and part in the equation. They simply act to speed up the natural process of hydrogen peroxide decomposition.)

The catalysts being used in this investigation are: Copper (II) Oxide CuO Iron (III) Oxide Fe2O3 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.

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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: 1. 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. 2. 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. 3. 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. 4. 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:  Give the molecules more energy i?? temperature  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.

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: Time (secs) Manganese Oxide (cm3) Iron Oxide (cm3) Copper Oxide (cm3)

Rate of Reaction 1. 7cm3/s 0. 6cm3/s 0. 25cm3/s 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.

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