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?
The setup: put a vial of HCl inside a flask of CaCO3. Plug the top with a stopper and suck out air to make it airtight, and there are instruments in the stopper that measure the inside temperature and pressure. When we are ready to measure, we tip the vial so that the HCl reacts with the CaCO3 to produce CO2 and since the flask should be airtight, the pressure will go up.
Personally, I don't think a wet flask CAN account for high results (or low results, for that matter). The CaCO3 is essentially insoluble in water and won't react with any water there before the HCl vial is tipped. The CaCO3 can't tell the difference between water that might be there before the HCl vial is tipped and the water that's in the HCl solution. I think you need to look for other reasons that would account for the high results. Perhaps the tablet's mass was recorded or weighed incorrectly. Perhaps the flask was not cleaned properly between trials 2 and 3.
Perhaps another tutor will give his/her opinion.
Nevermind I think i figured it out. But CaCO3 is supposed to react with water in the presence of CO2
CO2(g) + CaCO3(s)==> Ca(HCO3)2(aq)
action with acid
properties of hydrogen
can you please send your answer here ?
To explain the difference in results between trials 1 & 2 and trial 3 using a wet flask, let's consider the reaction that takes place between the HCl and calcium carbonate (CaCO3):
CaCO3 + 2HCl -> CaCl2 + H2O + CO2
In trials 1 & 2, where the percent concentrations of calcium carbonate were consistent, the reaction proceeded as expected. However, in trial 3, where a wet flask was used, water had the opportunity to react with the calcium carbonate before the actual measurement.
When water interacts with calcium carbonate, it can form a compound called calcium hydroxide (Ca(OH)2). This reaction is:
CaCO3 + H2O -> Ca(OH)2 + CO2
The formation of calcium hydroxide consumes a certain amount of calcium carbonate. Therefore, if the flask was wet in trial 3, some of the calcium carbonate would already be converted to calcium hydroxide before the HCl is introduced.
Now, you're correct in thinking that a gas law could help explain the results. In this case, we can use the ideal gas law, which states:
PV = nRT
Where:
P = pressure
V = volume
n = moles of gas
R = ideal gas constant
T = temperature
Since we are measuring the pressure increase, let's rewrite the equation as:
ΔP * V = nRT
Assuming the volume and initial temperature are constant for all three trials, we can focus on the number of moles of gas (n).
In trials 1 & 2, where the percent concentrations of calcium carbonate were consistent, we can assume that the number of moles of CO2 produced is proportional to the percent of calcium carbonate present. Therefore, the calculated moles of CO2 can be directly related to the concentration of calcium carbonate in the antacid tablet.
In trial 3, where the wet flask was used, it is likely that some of the calcium carbonate had reacted with water to form calcium hydroxide. As a result, the moles of CO2 measured will be lower than expected, leading to a higher concentration of calcium carbonate calculated.
So to summarize, the use of a wet flask allowed water to react with some of the calcium carbonate before measurement, resulting in a decreased amount of CO2 produced and an erroneously higher concentration of calcium carbonate calculated.
In conclusion, the wet flask condition in trial 3 affected the accuracy of the experiment because it caused an additional reaction that depleted the original amount of calcium carbonate, consequently yielding a higher measured concentration.