At $380 \ K$,$NH_4HS$ decomposes and the total pressure is $1.12 \ bar$. Find $K_p$ for the reaction: $NH_4HS_{(s)} \rightleftharpoons NH_{3(g)} + H_2S_{(g)}$ (in $bar^2$)

At $717 \ K$,$1.50 \ mol$ each of hydrogen and iodine are placed in a $10 \ L$ closed vessel. At equilibrium,$1.25 \ mol$ of each hydrogen and iodine remain. The value of the equilibrium constant $K_c$ for the reaction $H_2(g) + I_2(g) \rightleftharpoons 2HI(g)$ at $717 \ K$ is .......

If for a reaction $\Delta G^{o} > 0$,then:

For the reversible reaction $A + B \rightleftharpoons C + D$,the equilibrium concentrations of $C$ and $D$ are $0.8 \ mol/L$ each. If the initial concentrations of $A$ and $B$ were $1 \ mol/L$ each,calculate the equilibrium constant $K_c$.

If $K_C$ for the equilibrium reaction $2 ABC_{(g)} \rightleftharpoons 2 AB_{(g)} + C_{2(g)}$ is $X$ at $T \ K$,its $K_P$ at the same temperature is:

At equilibrium,the concentrations are $[N_2] = 3.0 \times 10^{-3} \ M$,$[O_2] = 4.2 \times 10^{-3} \ M$,and $[NO] = 2.8 \times 10^{-3} \ M$ in a sealed vessel at $800 \ K$ and $1 \ atm$ pressure. What will be $K_p$ for the given reaction?
$N_{2(g)} + O_{2(g)} \rightleftharpoons 2NO_{(g)}$