# Question #f31ca

Mar 12, 2017

We are flying a bit blind here.......you have take this question from a laboratory manual WITHOUT the context of the experiment.

#### Explanation:

Given your question, the most likely answer is the first option, i.e. to create a solution where the impurities can be filtered (i.e. this is NOT the reason why we add acid to the pill).

When we take a pill (of any sort, aspirin, prescription drugs, the contraceptive pill), MOST of the mass of the pill is calcium carbonate. Why? Because calcium carbonate is reasonably inert (none too basic), fairly insoluble, and during manufacture of the pill the dry ingredients can be mixed together in the appropriate mass ratios, and then this mixture can be machined into pills that contain the appropriate mass of drug. Drug companies would be very careful in this respect. They don't want to give some poor punter a double dose of Viagra, while the next user gets a dud dose.

Most of the drug (I assume) will be water soluble. So when we add acid to calcium carbonate, we get the following reaction:

$C a C {O}_{3} \left(s\right) + 2 H C l \left(a q\right) \rightarrow C a C {l}_{2} \left(a q\right) + {H}_{2} O \left(l\right) + C {O}_{2} \left(g\right) \uparrow$

And I assume then when you add the acid, part of your experimental observation will read: $\text{the crushed pill was treated with acid}$ $\text{to yield some effervescence and a clear solution}$. But you have done the experiment. I assume that when you added the acid you got a clear solution with no PRECIPITATE. If you did not filter this solution, then it is the first option.