Question #b41c8
1 Answer
Here's what I got.
Explanation:
You know that a weak monoprotic acid will only partially ionize to produce hydronium cations and the conjugate base of the acid in
#"HA"_ ((aq)) + "H"_ 2"O"_ ((l)) rightleftharpoons "H"_ 3"O"_ ((aq))^(+) + "A"_ ((aq))^(-)#
Now, you know that you start with
This implies that
#"175 moles " - " 1.15 moles" = "0.60 moles"#
of weak acid have ionized to produce hydronium cations and the conjugate base of the weak acid.
This implies that, at equilibrium, the solution contains
Use the volume of the solution to calculate the equilibrium concentrations of the species involved in the reaction.
#["HA"] = "1.15 moles"/"1.5 L" = "0.767 mol L"^(-1)#
#["H"_3"O"^(+)] = "0.60 moles"/"1.5 L" = "0.40 moles L"^(-1)#
#["A"^(-)] = "0.60 moles"/"1.5 L" = "0.40 moles L"^(-1)#
By definition, the acid dissociation constant,
#K_a = (["H"_3"O"^(+)] * ["A"^(-)])/(["HA"])#
Plug in your values to find--I'll leave the answer without added units
#K_a = (0.40 * 0.40)/(0.767) = color(darkgreen)(ul(color(black)(0.21)))#
The answer is rounded to two sig figs, the number of sig figs you have for the volume of the solution.
Finally, the
#color(blue)(ul(color(black)("pH" = - log(["H"_3"O"^(+)]))))#
In your case, the solution will have a
#"pH" = - log(0.40) = color(darkgreen)(ul(color(black)(0.40)))#
The answer is rounded to two decimal places because you have two sig figs for the volume of the solution.
Now, notice that the
In your case, the acid does not ionize completely, i.e.