Question

Calculate the percentage ionization of 0.01 M acetic acid in 0.1 M HCl. \[{K_a}\] of acetic acid is \[1.8 \times {10^{ - 5}}\].
A.) 0.18%
B.) 0.018%
C.) 1.8%
D.) 18%

Answer 1

Hint: HCl is a strong acid. It dissociates completely into hydrogen ion and chlorine ion. However, acetic acid is a weak acid. It undergoes a common ion effect by the hydrogen ions of HCl and its ionization decreases.

Complete step by step answer:
A weak acid is an acid that ionizes only slightly in an aqueous solution. An example of a weak acid is acetic acid. The extent of ionization of weak acids varies but is generally less than 10%. Weak acids, like strong acids, ionize to yield the \[{H^ + }\] ion and a conjugate base. Also, the \[{H^ + }\] ions of acetic acid undergo the common ion effect by the \[{H^ + }\] ions of HCl. The common-ion effect is the decrease in solubility of an ionic precipitate by the addition to the solution of a soluble compound with an ion in common with the precipitate. It is a consequence of Le Chatelier’s principle for the equilibrium reaction of the ionic association/dissociation.
In the given question, acetic acid decomposes according to the reaction:
\[\mathop {C{H_3}COOH}\limits_{0.01 - 0.01\alpha } \to \mathop {C{H_3}CO{O^ - }}\limits_{0.01\alpha } + \mathop {{H^ + }}\limits_{0.01\alpha + 0.1} \]
The equilibrium equation is:
\[{K_a} = \dfrac{{[C{H_3}CO{O^ - }][{H^ + }]}}{{[C{H_3}COOH]}}\]
\[1.8 \times {10^{ - 5}} = \dfrac{{(0.01\alpha )\left( {0.01\alpha + 0.1} \right)}}{{0.01 - 0.01\alpha }}\]
As <<< 1, 0.01 <<< 0.01
\[\therefore 1.8 \times {10^{ - 5}} = \dfrac{{(0.01\alpha )(0.1)}}{{0.01}}\]
\[\therefore \alpha = 1.8 \times {10^{ - 4}}\]
Percentage dissociation of acetic acid = 100%
                         = 0.018%

Hence, the correct answer is (B) 0.018%.

Note: Remember that the common ion effect is by the strong acid on the weak acid. It hinders its dissociation if they have any ion in common.