I think if the amount of HCA particles are increased the will be more chance of them colliding with the SSC, and by doing so cause will cause more reactions.
I assume If I make the concentration twice as strong the reaction will be twice as fast this is because there is twice as many particles of HCA that could collide with the SCC.
I can assume this because as I said before a reaction will occur and when that happens a product will formed, this means there must be an equation. By researching I found that the equation for this reaction is
CaCO3(s) + 2HCl (aq) CaCl2(aq) + CO2(g) + H2O(l)
If I double one side the other side must double too.
2CaCO3(s) + 4HCl (aq) 2CaCl2(aq) + 2CO2(g) + 2H2O(l)
I’ve doubled what went in to the equation and in doing so doubled the reaction.
So my results should gradually rise as I increase the concentration. If I put my results in a graph it should look like this.
The amount of carbon dioxide collected should rise as the concentration thickens, though at a certain point I think the reaction rate will level because the two substances can only react so fast.
Fair test:
To keep the experiment a fair test, the factors that have to be kept constant are:
The amount of water in the burette- if this were to change the measurements would be different each time.
The amount of hydrochloric acid used- if this were to change the rate of reaction would be affected because there would be more acid particles in a large amount of acid than in a smaller amount of acid of the same molar.
The amount of lime stone used- if this were to change, the rate of reaction would be affected because more limestone would take longer to fully react with the acid.
Temperature- if this were to change it would affect the results because it would speed up or slow down the particles, causing either more or less collisions between particles.
The input variable in this experiment is the acid concentration and the output variable is the rate of reaction, measured by how much carbon dioxide gas is produced.
Apparatus: Burette, beaker of water, flask with tube, hydrochloric acid, limestone.
Method:
Limestone is put into the flask filled with hydrochloric acid.
The bung connected to the tube is put on the flask.
The gas produced travels along the tube and is released up the burette filled with water.
The carbon dioxide produced displaces some of the water in the burette.
The gas produced is measured using the burette, to show the rate of the reaction.
The experiment is repeated for different measurements of concentration.
For each measurement of concentration the experiment will be repeated a minimum of three times to ensure accuracy and to delete any anomalous results.
Pre-test:
___ = Anomalous result, which is therefore not included in the average result.
I carried out a pre-test to find out what range was suitable to use. I started with increasing by 0.2 M of acid concentration, but found that this gave to wide a gap between each result. Therefore I decided to go up in 0.1M of acid concentration. I decided to stop at 0.9 M of acid concentration because I tried a few times with 10M but the burette was not long enough to hold all of the gas produced, so I wouldn’t have been able to get an accurate reading.
Results:
Analysis:
The results show that as the concentration of the acid increases, the rate of reaction will increase, which means my prediction was correct. When the acid concentration was 0.1M, the rate of reaction, shown by the amount of carbon dioxide produced, was 6 cm²/min. When the acid concentration was 0.9M, the amount of carbon dioxide released was 45.2 cm²/min. The graph backs this up, showing a strong correlation in results.
This is because the more particles of acid there are in the solution, the more collisions between acid and limestone particles will occur. Chemical reactions only occur when reacting particles collide with each other, so the higher the concentration of the acid, the more acid particles there are to collide with the limestone particles.
Evaluation:
There were a few anomalies in my results. This could have been down to taking too long to put the bung on the flask, which would have meant some gas would’ve been lost, or not all the gas produced going up the burette. The results could have also been affected by the surface area of the limestone. Although it is easy to keep the same mass of limestone used in the experiment, each piece would have had a different surface area which would have affected the results because the larger the surface area, the faster the rate of reaction because more limestone particles would have been exposed to the hydrochloric acid in pieces of a larger surface area.
To extend this investigation I could use higher concentrations of acid to see if the results level out as the concentration gets higher, or if the rate of reaction continues to go up. I could also experiment with other factors that would affect the rate of reaction, like temperature and surface area of the limestone.
