If the equilibrium constant for $\mathrm{A}<mspace\; width="1em"></mspace>\rightleftharpoons <mspace\; width="1em"></mspace>\mathrm{B}<mspace\; width="1em"></mspace>+\hspace{0.17em}\mathrm{C}<mspace\; width="1em"></mspace>\mathrm{is}<mspace\; width="1em"></mspace>{\mathrm{K}}_{\mathrm{eq}}^{\left(1\right)}$ and that of $\mathrm{B}\hspace{0.17em}+\hspace{0.17em}\mathrm{C}<mspace\; width="1em"></mspace><mspace\; width="1em"></mspace>\rightleftharpoons <mspace\; width="1em"></mspace>\mathrm{P}<mspace\; width="1em"></mspace>\mathrm{is}<mspace\; width="1em"></mspace>{\mathrm{K}}_{\mathrm{eq}}^{\left(2\right)}$, the equilibrium constant for $A<mspace\; width="1em"></mspace>\leftrightharpoons <mspace\; width="1em"></mspace>P$ is: 

1. ${\mathrm{K}}_{\mathrm{eq}}^{\left(2\right)}<mspace\; width="1em"></mspace>-<mspace\; width="1em"></mspace>{\mathrm{K}}_{\mathrm{eq}}^{\left(1\right)}$

2. ${\mathrm{K}}_{\mathrm{eq}}^{(1)\hspace{0.17em}}\times <mspace\; width="1em"></mspace>{\mathrm{K}}_{\mathrm{eq}}^{\left(2\right)}$

3. ${\mathrm{K}}_{\mathrm{eq}}^{\left(1\right)}<mspace\; width="1em"></mspace>+<mspace\; width="1em"></mspace>{\mathrm{K}}_{\mathrm{eq}}^{\left(2\right)}$

4. ${\mathrm{K}}_{\mathrm{eq}}^{\left(2\right)}<mspace\; width="1em"></mspace>/<mspace\; width="1em"></mspace>{\mathrm{K}}_{\mathrm{eq}}^{\left(1\right)}$

The correct expression for the following reaction is:

Fe₂N_(s) + \(\frac{3}{2}\)H_2(g) \(\leftrightharpoons \) 2Fe_(s) + NH_3(g)

For the reaction, 2SO_2(g) + O_2(g) = 2SO_3(g), $∆$H = –57.2 kJ mol^–1 and K_C = 1.7 × 10¹⁶ .
Which of the following statements is incorrect?

In which one of the following equilibria, K_p \(\ne\) K_C?

1. $2NO\left(g\right)\rightleftharpoons N_2\left(g\right)+O_2\left(g\right)$

2. $2C\left(s\right)+O_2\left(g\right)\rightleftharpoons 2CO\left(g\right)$

3. $2HI\left(g\right)\rightleftharpoons H_2\left(g\right)+I_2\left(g\right)$

4. $N{O}_2\left(g\right)+S{O}_2\left(g\right)\rightleftharpoons NO\left(g\right)+S{O}_3\left(g\right)$

The standard Gibbs energy change at 300 K for the reaction $2\mathrm{A}\rightleftharpoons \mathrm{B}+\mathrm{C}$ is 2494.2 J. At a given time, the composition of the reaction mixture is $\left[\mathrm{A}\right]=\frac{1}{2},<mspace\; width="1em"></mspace>\left[B\right]=2<mspace\; width="1em"></mspace>and<mspace\; width="1em"></mspace>\left[C\right]=\frac{1}{2}$. The reaction proceeds in the:[assume R=8.314 J/K/mol; e^-1 = 0.37]

If the equilibrium constant \(\left(K_c\right)\) for the reaction \(\mathrm{N}_2(\mathrm{g})+\mathrm{O}_2(\mathrm{g}) \rightarrow 2 \mathrm{NO}(\mathrm{g})\) at temperature \(T\) is \(4 \times10^{-4}\) then what will be the value of \(K_c\) for the reaction, \(\mathrm{NO}(\mathrm{g}) \rightarrow \frac{1}{2} \mathrm{N}_2(\mathrm{g})+\frac{1}{2} \mathrm{O}_2(\mathrm{g})\) ?

1. \(2.5 \times10^{2}\)
2. \(4 \times10^{-4}\)
3. \(50.0\)
4. \(0.02\)

The first and second dissociation constants of an acid H₂A are 1.0 × 10^–5 & 5.0 × 10^–10 respectively. The overall dissociation constant of the acid will be :

The equilibrium constant for the reaction
\(\mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \rightleftharpoons 2 \mathrm{NO}(g)\)
at temperature T is \(4 \times 10^{-4}\) . The value of K_c for the reaction,
\(\mathrm{{NO}}(g) \rightleftharpoons \frac{1}{2} \mathrm{{~N}_{2}}(g)+\frac{1}{2} \mathrm{O}_{2}(g)\)at the same temperature) is: