On addition of #NaOH# to #CH_3COOH# solution, #60%# of the acid is neutralised. If #pK_a# of #CH_3COOH# is 4.7 then the pH of the resulting solution will be?

1 Answer
Truong-Son N.
Feb 3, 2018

Acetic acid is a weak enough acid that the Henderson-Hasselbalch equation applies...

#"pH" = "pK"_a + log((["base"]_(eq))/(["acid"]_(eq)))#

Since we know that #60%# of weak acid is neutralized, it must be that #60%# of it became conjugate weak base. Clearly it has one proton, so the mol ratio is #1:1#.

#"HA"(aq) + "OH"^(-)(aq) rightleftharpoons "A"^(-)(aq) + "H"_2"O"(l)#

Thus:

#color(blue)("pH") = 4.75 + log(((0.00+0.60)n_("A"^-))/((1.00-0.60)n_"HA"))#

#= 4.75 + log(((0.00+0.60)n_"HA")/((1.00-0.60)n_"HA"))#

#= 4.75 + log(0.60/0.40)#

#= color(blue)(4.93)#

Why does this make sense? You added base, so why did the pH rise?

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