We did a lab where we reacted antacid with HCL and measured the pressure increase after the reaction. Then we found the percent concentrations of calcium carbonate of the antacid tablet. The question is: say the percent of calcium carbonate from trials 1 & 2 are consistent while trial 3 is much higher. How could a wet flask be used to explain this result? A wet flask would allow water to react with the calcium carbonate before we measure it, but i'm not sure what do say next. I think I need to use a gas law?
I can't tell from your description but it appears that you used the HCl in a flask containing the antacid tablet and the gas generated flowed into another flask. If that is the case, then the wet "second flask" could absorb some of the CO2 gas generated. Was the pressure lower than the others? That would explain it. Higher than the others, there must be another explanation. If the set up is not as I have it pictured, please repost and describe the set up for us. Otherwise, we're operating in the dark.
To explain the result of trial 3 being higher when using a wet flask, we can consider the reaction of water with calcium carbonate and how it affects the overall pressure increase observed.
When water reacts with calcium carbonate, it forms carbon dioxide gas, which is responsible for the pressure increase in the experiment. The reaction can be represented by the following equation:
H2O + CaCO3 -> CO2↑ + H2O + Ca^(2+)
In trials 1 and 2, where a dry flask was used, the calcium carbonate in the antacid tablet reacts directly with hydrochloric acid (HCl) to produce carbon dioxide gas. The pressure increase observed in these trials reflects the amount of calcium carbonate present in the tablet.
However, in trial 3 where a wet flask was used, water reacts with calcium carbonate before the addition of HCl. This means that some of the calcium carbonate may have already reacted to form carbon dioxide gas before the actual experiment started. As a result, the observed pressure increase in trial 3 would be higher than expected.
To determine the difference in pressure increase due to water reacting with calcium carbonate, we can apply the ideal gas law. The ideal gas law states:
PV = nRT
Where:
P = pressure
V = volume
n = number of moles of gas
R = gas constant
T = temperature
In this case, the volume and temperature are constant throughout all the trials, so we can simplify the equation to:
P ∝ n
This means that the pressure is directly proportional to the number of moles of gas. If water reacts with calcium carbonate to form carbon dioxide gas before the experiment, then the amount of calcium carbonate available to react with HCl will be reduced, resulting in a smaller number of moles of gas produced and a lower pressure increase.
In trial 3, where a wet flask was used, it is expected that the pressure increase will be higher since the calcium carbonate has already reacted partially with water, leading to less calcium carbonate available for reaction with HCl. This explains the higher observed pressure increase in trial 3 compared to trials 1 and 2.
In conclusion, the use of a wet flask in trial 3 allows water to react with calcium carbonate before the experiment, leading to a decrease in the available calcium carbonate for reaction with HCl. This results in a higher pressure increase than expected.