Acetic acid dissociates to $1.20 \%$ in its $0.01 \ M$ solution. What is the value of its dissociation constant?
- A$2.20 \times 10^{-2}$
- B$1.60 \times 10^{-4}$
- C$1.44 \times 10^{-6}$
- D$2.40 \times 10^{-4}$
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Weak acid $HX$ has dissociation constant $1 \times 10^{-5}$. Calculate the percent dissociation in its $0.1 \ M$ solution. (in $\%$)
The degree of ionization of $0.4 \ M$ acetic acid will be $(K_{a} = 1.8 \times 10^{-5})$.
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View SolutionWhat is the percent dissociation of acetic acid in its $0.01 \ M$ solution if the dissociation constant of the acid is $1.34 \times 10^{-6}$ (in $\%$)?
For the ionization of a weak acid $HA$ as $HA \rightleftharpoons H^{+} + A^{-}$,the $pH$ of a $1.0 \ M$ solution is $5$. The dissociation constant $K_a$ is:
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