4. Measuring the release of carbon dioxide.
As there is a limited amount of equipment we are unable to do numbers 1 to 3, however, there is sufficient equipment to measure the release of carbon dioxide.
Factors
There are many factors that could alter the accuracy of the results I get.
Pressure – If I was to increase the pressure in the conical flask the calcium carbonate and hydrochloric acid particles would be forced closer together and would be less dispersed, therefore are more likely to collide. This would speed up the rate of reaction.
Temperature – If the hydrochloric acid was heated the reaction would be much quicker because the heat would give the acid particles more energy which would make them faster and they would collide with the calcium carbonate more frequently and with a lot more force causing the rate of reaction to be accelerated.
Concentration of acid – If I was to use a 0.5 molar concentration of acid it would take much longer to react with a 5mg piece of calcium carbonate than 2.0 molar acid would take to react with the same mass of calcium carbonate this is because the percentage of the solution that is acid is much higher in a 2.0 molar concentration , therefore there are more acid particles to collide with the calcium carbonate particles so a higher acid concentration would speed up the reaction time.
Surface area – If I used a piece of calcium carbonate with a large surface area and then a piece with a smaller surface area, the piece with the larger surface area would react more quickly with the hydrochloric acid than the piece with the smaller surface area. If I was to use powdered calcium carbonate rather than a block the powder would react more quickly this is because there is a larger area of calcium carbonate for the hydrochloric acid particles to collide with, as shown in the diagram below.
Powdered calcium carbonate
Hydrochloric acid molecules are able to get to
the calcium carbonate more easily and can collide
with each molecule far more quickly,
this speeds up the rate of reaction.
Lump of calcium Carbonate
Hydrochloric acid molecules have to break
down the lump of calcium carbonate layer
by layer which takes much longer and
slows down the reaction speed.
Method
You will need…
- 5 concentrations of hydrochloric acid …
* 0.10 molar
* 0.50 molar
* 1.00 molar
* 1.75 molar
* 2.00 molar
- 75mg of medium sized calcium carbonate (5mg per experiment)
- 1 gas syringe with rubber bung
- 1 conical flask
- measuring cylinder
- electronic scales (to measure the mass of calcium carbonate)
- Stopwatch
- Clamp Stand
Step 1- Set up the equipment in the list above as shown in the diagram on the following page.
Step 2- Measure 20ml of hydrochloric acid and pour into conical flask.
Step 3- Weigh 5mg of calcium carbonate and add it to the hydrochloric acid in the conical flask and quickly put the rubber bung in the top of the conical flask and start stopwatch.
Step 4 - Time the reaction for 1 minute and read the measure on the gas syringe after 1 minute. (If the gas syringe reaches 100ml before 1 minute stop the timer and record the result for the length of time that it took to get to 100ml)
Step 5- Clean out conical flask and repeat the experiment for the desired amount of times.
Fair testing
During this experiment I will only change the concentration of acid.
- The temperature was kept constant because the heat in laboratory does not vary enough to affect this experiment.
- The pressure was kept constant because again the laboratory pressure does not alter enough to affect this experiment.
- I used an electronic balance to weigh the mass of the limestone each time; therefore there would not be enough difference in mass to affect the investigation.
- The same equipment was used each time in order to keep my results fair.
- There was only one size of limestone used, so as not to differentiate the surface area.
- I calculated the rate so as the results are over an equal time.
- To ensure that the volume of acid remained the same throughout the experiment I used a measuring cylinder each time.
Time
From a previous investigation I found that I got suitable results using 20ml of acid and 5mg of limestone over a 1 minute period.
The results of this experiment are included.
Analysis
My graph indicates that the largest increase in reaction rate occurs between 0.1 and 0.5 acid concentrations. The gradient of this increase is 16.67. However the actual largest gradient shown on my graph is between 1.75 and 2.0 acid concentrations the gradient is 22. I would have expected the largest increase in the rate of reaction to have been between 1.0 and 1.75 acid concentrations as there is a difference of 0.75 moles which is the largest gap between concentrations but the gradient at this point on the graph shows a gradient of only 9.
The line of best fit shows that as the acid concentration increases so does the rate of reaction this is because the ratio of acid to water is increasing, therefore increasing the number of acid particles in the solution. The graph also shows a positive correlation.
Evaluation
As I mentioned in my evaluation I had one anomalous result, this was the rate of reaction between 1.0 and 1.75 molar acid concentrations, in comparison to my other results the reaction rate should have shown a larger difference than what it did as there is a greater difference between the concentrations of acids used than the rest. For better accuracy there should be a larger range of acids used between 1.0 and 1.75, for example: 1.0, 1.1, 1.2, 1.3, 1.4 and so on. This could be done throughout in order to give far more accurate results. I think that there are many possible reasons for the anomalous result that I got, such as, tampered with acids, larger surface area of limestone (the different sizes of limestone may have become muddled up), there may have been a slight leak in the rubber bung and gas could possibly have dispersed into the surrounding atmosphere.
The experiment could have been more accurate by using an electronic balance to measure the loss of mass rather than using a gas syringe to measure the gas given off by the hydrochloric acid and limestone.
