Question

Which of the following result(s) is/are correct for the equilibrium state in a solution originally having $0.1M-C{H}_{3}COOH$ and $0.1M-HCl$? $K_a$ of $C{H}_{3}COOH=1.8\times {10}^{-5}$.

• A. $\left[H^{+}\right]=0.1M$
• B. $\left[C{H}_{3}CO{O}^{-}\right]=1.8\times {10}^{-5}M$
• C. Degree of dissociation of acetic acid $=1.8\times {10}^{-4}$
• D. $\left[H^{+}\right]$ from water ${10}^{-13}M$

Answer 1

The correct option is A $\left[H^{+}\right]=0.1M$

$\left[H^{+}\right]=0.1M\to H^{+}Concentration$
The introduction of HCl into an acetic acid solution would shift the equilibrium of the acetic acid forward forming less ions.
This is Le-Chatlier's principle, since acetic acid is weak acid.