-spatula
-glass beaker
-bung (with hole through it)
-rubber tubing (approximately 60cm)
-large, clear, flat bowl
-500cm measuring cylinder
(I intend to use a 500cm measuring cylinder as it is the largest available, therefore the margin for fault is less than if I used, for example, a 100cm measuring cylinder.)
-cling film
-bunsen burner
-tripod
-tap water
-1.5 grams of copper carbonate
Diagram of apparatus set up:
It is very important that there aren’t any chemicals left from other experiments on the equipment as they may affect the results, so all of the equipment must be washed thoroughly before use. For the volumes measured in this experiment to be accurate, it is vital that the experiment is carried out at constant room temperature of 25 and pressure of 1 atmosphere. Without these conditions the rule that 1 mole of gas occupies 24dm does not apply and it would be impossible to find out which equation is correct by measuring the volume of gas produced. The volume of gas measured may also be altered by the fact that the rubber tubing will contain air before the experiment has started, and this gas will add to the final volume collected , making the measurement less accurate. Another factor, which may make the results less precise, is the accuracy of the scales used to weigh the copper carbonate. It is important to use the most precise scales available.
Method:
-Firstly, place the glass beaker on a pair of scales and reset the weight to zero. Then, using a spatula, measure out as accurately as possible, 1.5 grams of the green copper carbonate powder. Place the glass beaker on the tripod, above the bunsen burner.
-Put one end of the rubber tubing through the bung, and tightly seal the glass beaker with it.
-Fill around half of the bowl with tap water, so that it allows for more water which will fill it during the experiment. Fill the measuring cylinder completely with tap water, then cover the top of it with cling film to prevent water from escaping.
-Carefully, holding the clingfilm down, turn the measuring cylinder upside down and place the top end into the bowl of water.
-Remove the clingfilm, and make sure that there are no bubbles of gas in the measuring cylinder. It is very important that there are not as the reliability of the results depends on it.
-Take the free end of the rubber tubing and put it into the bowl of water, so that the end is directly underneath the measuring cylinder. For this part of the experiment it is necessary to work in a group of at least two, so that one person can hold the measuring cylinder steady, whilst the other places the end of the rubber tube underneath it. It is important that this is done properly, so that none of the gas will escape during the reaction.
-Using a lighted splint carefully light the bunsen burner under the tripod. As the reaction takes place bubbles of gas will travel along the rubber tubing and into the measuring cylinder, causing the water level to decrease. When all of the copper carbonate has decomposed there will be no more bubbles of gas and a black powder (copper oxide) will be left in the glass beaker.
-At this point read off the volume of gas that has been produced from the measuring cylinder and record it. When doing this make sure your eyes are level with the water level, so that you can clearly see the volume markings, and make an accurate reading. The volume should indicate which equation is correct for the reaction that has taken place.
I have chosen to use this method, as I think that it is the most reliable way within my means to find out which equation is correct. If carried out carefully and safely, according the above method, all of the gas produced should be obtained. And, allowing for a small margin of error, due to air in the rubber tubing and a small amount of air which may get into the measuring cylinder when it is turned over, the volume of gas measured should be equal or close to one of the volumes I have predicted on page 1, therefore indicating which equation is correct for the reaction.
Results
To get the rate of reaction I will need to perform the following calculations:
20% = 48.7 – 46.6 = 2.1
49.3 – 47.2 = 2.1
49.1 – 47.2 = 1.9
= (2.1 + 2.1 + 1.9) ÷ 3
= 6.1 ÷ 3
=2.03 ÷ 10
= 0.203 cm³s¯¹
The same calculations needs to take place for the other concentrations (i.e. 40%, 60%, 80% and 100%) and this will give us the average rate of reaction for all the for all the concentrations. The average rates of reaction for the concentrations are:
Conclusion
As the concentration of hydrogen peroxide increases the rate of reaction also increases at more or less proportional values this is evident from the graph as we can clearly see that the graph is more or less increasing in a linear form. We cal also see this from the results that at a 20% concentration of hydrogen peroxide the average rate of reaction is 0.20cm³s¯¹ and at 100% concentration of hydrogen peroxide the average rate of reaction is 0.99cm³s¯¹.
From the above we can clearly see that generally the amount of gas being released is at a steady increase this is due to the fact that when the reaction occurs more substrate molecules are available to bind to the active site on the liver homogenate and as they are more substrate molecules available it is likely that collisions occurs which form the enzyme-substrate complex. As there is an enzyme present the activation energy of the reaction is lowered, the activation energy is the energy required to break the bonds so a reaction can occurs. For example sugar only reacts with oxygen when heat is applied, the energy from the heat gives the substrate molecules more kinetic energy and as a result the molecules of glucose vibrate more vigorously and this result in collisions with oxygen molecules and these bonds break. Once these bonds break new bonds can form.
Enzymes are very good at lowering the activation energy as the diagram of a graph below shows us:
Therefore a more concentrated substance will have more collisions resulting in a faster rate of reaction and a diluted concentration will result in a slower rate of reaction as collisions between substrate and active sites are less likely. And the reaction slows down due to the fact that the substrate molecules are being used up.
Limitations
I believe it was a number of factors such as the bung being placed onto the conical flask after the reaction had already started and the time not being measured accurately as well as not being familiar to the apparatus resulted in the anomalous result I had.
Overall I was able to determine that the different amount of gas produced and the concentration of the hydrogen peroxide was proportional.
Bibliography
The books that I used for some of my research for this project were:
- Revise AS Biology for AQA (B) by Graham Read and Ray Skwierczynski published by Heinemann.
- Biology AS second edition by Mike Bailey and Keith Hirst published by Collins.