For the particles to react, they must collide with a certain amount of energy, called the activation energy. Changing the concentration does not increase the proportion of collisions with enough energy, it increases the frequency of collisions.
I predict that the relationship between CO2 produced and concentration of acid is directly proportional. This means that if the concentration doubled, I would expect the rate of reaction to double. This is because there are twice as many hydrogen ions in the same volume, so there is twice as much chance of collision. This means that the frequency of collisions doubles, and so the speed of reaction doubles. I would expect my graph showing concentration of acid and rate of reaction to be a straight line graph through the origin.
Diagram
Apparatus
- Conical Flask
- Gas syringe
- Clamp
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75cm3 hydrochloric acid (concentration 2M)
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75cm3 water
- stopwatch
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4g Marble (CaCO3)
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25cm3 measuring cylinder
Method
- Set up apparatus as shown in diagram.
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Mix 25cm3 acid with the CaCO3 in the conical flask.
- Seal the conical flask using bung attached to gas syringe.
- Begin timing using stopwatch.
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Record the CO2 produced by the reaction every 10 seconds using the scale on the gas syringe, until 120 seconds, or the reaction has produced 100cm3 of CO2.
- Repeat steps 2-5 for the other 5 concentrations of acid.
Fair Test
To keep this experiment fair, and to ensure accurate results, I will keep all other factors that could affect the rate of reaction constant, with the exception of the variable I am investigating. I will keep temperature constant, because as the temperature of the acid increases, the speed of the reaction increases. This is because temperature affects the proportion of collisions which have enough energy to react (activation energy)
The graph above is showing how at a higher temperature, a higher proportion of collisions have enough energy to react, however the total number of collisions remains unchanged. It is especially important to control this factor, as a temperature change of just 10°c increases the average energy of the molecules by only 1%, but the reaction speed doubles. This is because there are a greater percentage of successful collisions at as the temperature increases.
I will keep the total volume of acid the same, changing the concentration by using different ratios of acid to water.
By using the same marble chips in each experiment, I can keep both volume and surface area of CaCO3 constant. It is important to keep the surface area constant because there are more collisions per second if there is a greater surface area. If I used chips with a large surface area, the speed of the reaction would increase.
By keeping temperature, volume of reactants, and surface area of marble constant, my test will be accurate.
Results
Accuracy – I have measured the volume of acid, water, and CO2 to the nearest cm3. I have measured time to the nearest second.
Instead of repeating the same values to ensure accuracy, I have decided to test a greater range, taking 6 readings instead of the specified 5. I think this will be more useful than repeating experiments, as it will give me a better idea of the general trend of results, and help me to identify anomalies.
Conclusion
As concentration of acid increases the speed of reaction increases. The speed of this reaction can be determined by the gradient of the graph showing time against CO2 produced. For each of the concentrations I tested I produced a graph. I calculated the gradient of each of these graphs using the formula ∆y /∆x. This showed me the speed the reaction was occurring for each concentration. I then plotted a graph of concentration against speed of reaction. This showed me that as concentration increased, speed of reaction increased.
This happened because for a reaction to occur, the particles must first collide. In this experiment the atoms which need to collide in order to react are hydrogen. In a higher concentration of acid, there are more hydrogen atoms. If there are more hydrogen ions within a given volume, these collisions occur more often. If these collisions react with enough energy (activation energy) they will react. If there are more collisions, a greater number will be successful than if there are only a few collisions. The greater the number of successful collisions, the faster the experiment occurs.
Results of speed of reaction using gradient of the graph
I calculated the concentration of acid in M, using the formula
Concentration = original concentration x total volume of acid
Total volume
As my final graph showing how concentration effects rate of reaction shows, after drawing on a straight line through the majority of points, the rate of reaction is directly proportional to the concentration of acid. I found that as the concentration doubles, the rate of reaction also doubles. This is because there are twice as many hydrogen ions in the same volume, so there is twice as much chance of collision. This means that the frequency of collisions doubles, and so the speed of reaction doubles.
Evaluation
I think the evidence I have obtained is reliable enough to support my prediction, and it shows the general trend I expected. The procedure was not very accurate, and there are several places where inaccuracies could have occurred. There is a significant time lag from when the reactants are mixed to when the first reading is taken. In the faster reactions, a significant part of the reaction could be over before readings begin to be taken. This could be avoided by finding a way to mix the two reactants whilst they were in the sealed equipment. If I were to do the experiment again, I would use the set-up shown in the diagram below. The acid and marble could be mixed by shaking the test tube, or turning it upside down.
When the bung is forced into the flask, it displaced air into the tube and syringe, which means gas is being measured which has not been produced by the reaction. Again, this could be prevented by using the set-up above. The bung would already be in the flask, and wouldn’t cause a displacement of air, and so an inaccurate reading.
Before any CO2 is seen in the gas syringe, it must saturate in the acid. This can make it appear as if the reaction begins very slowly, especially in the slow reactions.
Some of my graphs appear to jump between readings. This is because sometime the piston in the gas syringe gets stuck until enough pressure builds up behind it, forcing it to jump out quickly.
My final graph showed the results I expected, yet there were 2 anomalies. At a concentration of 0.8M the rate of reaction was much slower than I predicted. Looking at the graph showing just the rate of reaction at 0.8M, I can see that the reaction began very slowly. This is because it took a long time for the acid to become saturated with CO2. This makes the reaction appear much slower than it really was.
At a concentration of 1.6M the reaction appears to be much faster than I expected. This could be because of a rise in temperature of the acid. At a higher temperature, there are a higher proportion of successful collisions, and so the rate of reaction increases.
To further this investigation, I would extend my range of concentration of acid, to find out if my theory that if the concentration of acid doubles, the speed of reaction doubles is still correct at much stronger acids. I would also test the rate of reaction over a longer time, to see if the rate is constant throughout. I could also investigate another factor which would affect the rate of reaction, such as temperature of acid.
Nicola Dearnley 10a1
