$\mathrm{I}_2+\mathrm{I}^{-} \rightleftharpoons \mathrm{I}_3^{-}$This reaction is set-up in aqueous medium. We start with 1 mole of $\mathrm{I}_2$ and 0.5 mole of $\mathrm{I}^{-}$in $1 \mathrm{~L}$ flask. After equilibrium is reached, excess of $\mathrm{AgNO}_3$ gave 0.25 mole of yellow precipitate. equilibrium constant is

JEE CHEMISTRY MCQ: Equilibrium

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\mathrm{I}_2+\mathrm{I}^{-} \rightleftharpoons \mathrm{I}_3^{-}

This reaction is set-up in aqueous medium. We start with 1 mole of

\mathrm{I}_2

and 0.5 mole of

\mathrm{I}^{-}

1 \mathrm{~L}

flask. After equilibrium is reached, excess of

\mathrm{AgNO}_3

gave 0.25 mole of yellow precipitate. equilibrium constant is

View answer and explanation

Correct option: A

1.33

Option A (1.33) is correct because the equilibrium constant is calculated based on the concentrations of the reactants and products at equilibrium.

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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