The dissociation constant of ${\rm{N}}{{\rm{H}}_4}{\rm{OH}}$ is $1.8 \times {10^{ - 5}}.$ The hydrolysis constant of ${\rm{N}}{{\rm{H}}_4}{\rm{Cl}}$ would be:

JEE CHEMISTRY MCQ: Equilibrium

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The dissociation constant of

\mathrm{N}{\mathrm{H}_4}\mathrm{OH}

1.8 \times {10^{ - 5}}.

The hydrolysis constant of

\mathrm{N}{\mathrm{H}_4}\mathrm{Cl}

would be:

View answer and explanation

Correct option: D

5.55 \times {10^{ - 10}}

Option D (

5.55 imes {10^{ - 10}}

) is correct because it is calculated using the formula

{K_H} = \frac{{{K_w}}}{{{K_a}}}

, where

K_w

is the ionic product of water and

K_a

is the dissociation constant of the weak acid.

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\mathrm{K}_1

\mathrm{K}_2

\mathrm{N}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}(\mathrm{g})\quad\ldots

\mathrm{NO}(\mathrm{g}) \rightleftharpoons \frac{1}{2} \mathrm{~N}_2(\mathrm{~g})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g})\quad\ldots

K{a_1},\,\,\,K{a_2}\,\,\,\mathrm{and}\,\,\,K{a_3}

\begin{array}{l} {H_2}S\leftrightharpoons{H^ + } + H{S^ - }\\ H{S^ - }\leftrightharpoons{H^ + } + {S^{2 - }}\\ {H_2}S\leftrightharpoons2{H^ + } + {S^{2 - }} \end{array}

K{a_1},\,\,K{a_2}\,,K{a_3}

\mathrm{pH}

0.1 \mathrm{~M}-\mathrm{N}_2 \mathrm{H}_4

\mathrm{N}_2 \mathrm{H}_4,

K_{{b}_1}=3.6 \times 10^{-6}, K_{{b}_2}=6.4 \times 10^{-12}, \log 2=0.3

\log 3=0.48

{\text{Ba}}{{\text{O}}_2}(s) \rightleftharpoons {\text{BaO}}(s) + {{\text{O}}_2}(g);\Delta H = - {\text{ve}}.

\text{O}_2