Given
$(i)$ $\begin{gathered}
HCN\left( {aq} \right) + {H_2}O\left( l \right) \rightleftharpoons {H_3}{O^ + }\left( {aq} \right) + C{N^ - }\left( {aq} \right) \hfill \\
{K_a} = 6.2 \times {10^{ - 10}} \hfill \\
\end{gathered} $
$(ii)$ $\begin{gathered}
C{N^ - }\left( {aq} \right) + {H_2}O\left( l \right) \rightleftharpoons HCN\left( {aq} \right) + O{H^ - }\left( {aq} \right) \hfill \\
{K_b} = 1.6 \times {10^{ - 5}} \hfill \\
\end{gathered} $
These equilibria show the following order of the relative base strength
$O{H^ - } > {H_2}O > C{N^ - }$
$O{H^ - } > C{N^ - } > {H_2}O$
${H_2}O > C{N^ - } > O{H^ - }$
$C{N^ - } > {H_2}O > O{H^ - }$
A $0.1\, M$ solution of $HF$ is $1\%$ ionized. What is the $K_a$
The $pH$ of $0.1\, M$ monobasic acid is $4.50$ Calculate the concentration of species $H ^{+},$ $A^{-}$ and $HA$ at equilibrium. Also, determine the value of $K_{a}$ and $pK _{a}$ of the monobasic acid.
${K_b}$ of $N{H_4}OH = 1.8 \times {10^{ - 5}}$ calculate $pH$ of $0.15$ $mol$ $N{H_4}OH$ and $0.25$ $mol$ $N{H_4}OH$ containing solution.
Heat of neutralisation of weak acid and strong base is less than the heat of neutralisation of strong acid and strong base due to
The ionisation constant of acetic acid is $1.8 \times 10^{-5}$. The concentration at which it will be dissociated to $2\%$, is