Question #2a988

The idea here is that aspirin (acetylsalicyclic acid) is a weak acid that does not dissociate completely in aqueous solution to form its conjugate base and produce hydronium ions, #"H"_3"O"^(+)#.

Let's start with a solution that contains taspirin, not the sodium salt of its conjugate base

Aspirin contains a carboxylic acid group, #-"COOH"#, so I'll represent it as #"R"-"COOH"#.

In aqueous solution, aspirin will partially dissociate to produce

#"R"-"COOH"_text((aq]) + "H"_2"O"_text((l]) rightleftharpoons "H"_3"O"_text((aq])^(+) + "R"-"COO"_text((aq])^(-)#

Adding hydrochloric acid, #"HCl"#, a strong acid that dissociates completely in aqueous solution, will increase the concentration of hydronium ions.

In order to counteract this increase in the concentration of hydronium ions, the equilibrium will shift to the left, since the reverse reaction will consume hydronium ions and the conjugate base of aspirin and produce aspirin - think Le Chatelier's Principle.

#"R"-"COOH"_text((aq]) + "H"_2"O"_text((l]) rightleftharpoons "H"_3"O"_text((aq])^(+) + "R"-"COO"_text((aq])^(-)#

#color(white)(aaaaa)stackrel(color(green)(larr))(color(white)(aaaaaaa)color(red)("Shift to the left")color(white)(aaaaaaaaaa)#

This will result in fewer ionized molecules of aspirin in solution. In essence, the conjugate base will be protonated to reform the weak acid #-># the solubility of aspirin decreases.

Now, your solution contains the sodium salt of aspirin's conjugate base, so here's what happens when you add the hydrochloric acid.

The sodium salt of aspirin's conjugate base is soluble in aqueous solution

#"R"-"COONa"_text((aq]) -> "Na"_text((aq])^(+) + "R"-"COO"_text((aq])^(-)#

The added hydronium ions will react with the conjugate base of aspirin to form molecular aspirin, i.e. the weak acid

#"R"-"COO"_text((aq])^(-) + "H"_3"O"_text((aq])^(+) -> "R"-"COOH"_text((s]) darr + "H"_2"O"_text((l])#

Since aspirin is "not quite soluble'" in aqueous solution, it will actually precipitate out of solution.