Thermodynamics of Equilibrium

The Haber’s process for the formation of ${\mathrm{NH}}_3$ at $298 \mathrm{K}$ is ${\mathrm{N}}_2+3{\mathrm{H}}_2\rightleftharpoons 2{\mathrm{NH}}_3;\mathrm{\hspace{0.17em}}\mathrm{\Delta H}=-46.0\mathrm{KJ}$. Which of the following is the correct statement?

The Haber’s process for the formation of ${\mathrm{NH}}_3$ at $298 \mathrm{K}$ is ${\mathrm{N}}_2+3{\mathrm{H}}_2\rightleftharpoons 2{\mathrm{NH}}_3;\mathrm{\hspace{0.17em}}\mathrm{\Delta H}=-46.0\mathrm{KJ}$. Which of the following is the correct statement?

Which of the following is correct at equilibrium?

Select the graph which can be used to determine standard free energy of the reaction given below:

${\mathrm{CaCO}}_3\left(\mathrm{s}\right)\rightleftharpoons \mathrm{CaO}\left(\mathrm{s}\right)+{\mathrm{CO}}_2\left(\mathrm{g}\right)$
Assume that the reaction is at equilibrium.

Among the following set of conditions, which condition necessarily holds true for a non-feasible process?
(Where, $\mathrm{K}=$equilibrium constant)

Calculate the temperature given that the value of ${\mathrm{K}}_{\mathrm{c}} = {10}^3$ and ${\mathrm{\Delta G}}_{\mathrm{f}}°$of $\mathrm{A}, \mathrm{B} \mathrm{and} \mathrm{C}$ are $-100, -200 \mathrm{and} -500 \mathrm{kJ}/\mathrm{mol}$and the reactants $\mathrm{A}$ and $\mathrm{B}$ is in equilibrium with $\mathrm{C}$.

The value of ${\mathrm{K}}_{\mathrm{c}}$ for the following reaction at $298 \mathrm{K}$ is:

$2{\mathrm{NH}}_3\left(\mathrm{g}\right)+{\mathrm{CO}}_2\left(\mathrm{g}\right)\rightleftharpoons {\mathrm{NH}}_2{\mathrm{CONH}}_2\left(\mathrm{aq}\right)+2{\mathrm{H}}_2\mathrm{O}\left(\mathrm{\ell }\right)$  

Given, standard Gibbs energy change $\left(\mathrm{\Delta G}°\right)$ at the given temperature is $-13.6\mathrm{kJ}/{\mathrm{mol}}^{-1}$

For the following reaction,

${\mathrm{C}}_2{\mathrm{H}}_6\left(\mathrm{g}\right)\rightleftharpoons {\mathrm{C}}_2{\mathrm{H}}_4\left(\mathrm{g}\right)+{\mathrm{H}}_2\left(\mathrm{g}\right)$

the ${\mathrm{k}}_{\mathrm{p}}=0.05 \mathrm{atm}$. What is the value of $\mathrm{\Delta G}°$ of the reaction at ${627}^{\circ }\mathrm{C}$?

Consider that for a hypothetical reaction, $\mathrm{\Delta H}°$ and $\mathrm{\Delta S}°$ are $-30 \mathrm{kJ} {\mathrm{mol}}^{-1}$ and $-100 {\mathrm{JK}}^{-1}{\mathrm{mol}}^{-1}$ at $300\mathrm{K}.$ What is the equilibrium constant for the reaction at $298 \mathrm{K}?$

Consider the curve between $\ln \mathrm{K}$ and $1/\mathrm{T}$. Which plot will be correct for exothermic reaction?

The plot of $\log \mathrm{K}$ against $\frac{1}{\mathrm{T}}$ is a straight line with a positive slope ($\mathrm{K}$ being the equilibrium constant of a reaction). The correct statement among the following is:

What is the value of free energy change for a reversible reaction at equilibrium?

The change in standard Gibbs energy for a reaction at equilibrium, $eg . \mathrm{PCl}_{5}(g) \rightleftharpoons \mathrm{PCl}_{3}(g)+\mathrm{Cl}_{2}(g),$ on addition of an inert gas at constant pressure and then at constant volume respectively are :

$5$ moles of $N_2{O}_4$ taken in a one litre flask dissociates to$20\%$ extent to reach equilibrium. What is $\Delta G°$ of the reaction ?

$\underset{\left(g\right)}{N_2{O}_4}\stackrel{}{\rightleftharpoons }\underset{\left(g\right)}{2N{O}_2}$

Solid $CaC{O}_3$ is heated in a closed container to $900 K.$ When equilibrium is reached, the pressure becomes$500 Torr.$Similarly, the equilibrium pressure at $1000 K$ was found to be $2000 Torr$. Hence $∆H$ for the given reaction.

$CaC{O}_3\left(s\right)\rightleftharpoons CaO\left(s\right)+C{O}_2\left(g\right)$

At this temperature range is:

Consider the reaction :

$A\rightleftharpoons B \Delta H_f^{\circ } kJ S^{\circ }J{K}^{-1}$

$A -24.27 246.4$

$B -25.44 252.3$

Then the composition of the equilibrium mixture at ${25}^{o}C$ is

Cotyledons are also called-

For the following reaction

$\mathrm{A}{\mathrm{g}}_{\left(\mathrm{a}\mathrm{q}\right)}^{+}+\mathrm{C}{\mathrm{l}}_{\left(\mathrm{a}\mathrm{q}\right)}^{-}\to \mathrm{A}\mathrm{g}\mathrm{C}{\mathrm{l}}_{\left(\mathrm{s}\right)}$

Given: $\mathrm{\Delta }{\mathrm{G}}_{\mathrm{f}}^{\mathrm{o}}$ , AgCl = - 112.44 kJ/mol; $\mathrm{\Delta }{\mathrm{G}}_{\mathrm{f}}^{\mathrm{o}}$ ${\mathrm{C}\mathrm{l}}^{–}$ = - 130 kJ/mol ; $\mathrm{\Delta }{\mathrm{G}}_{\mathrm{f}}^{\mathrm{o}}\mathrm{A}{\mathrm{g}}^{+}$ = 75 kJ/mol
Report your answer by rounding it upto nearest whole number. The ${\mathrm{K}}_{\mathrm{s}\mathrm{p}}$ of AgCl is $\mathrm{n}\times 10^{-10}$ .The value of 'n' is.

The standard enthalpy of neutralization of strong acid and strong base is $-57.3$ $\mathrm{k}\mathrm{J}$ ${\text{equiv}}^{-1}.$ If the enthalpy of neutralization of the first proton of aqueous ${\mathrm{H}}_2\mathrm{S}$ is $-33.7$ $\mathrm{k}\mathrm{J}\mathrm{}\mathrm{m}\mathrm{o}{\mathrm{l}}^{-1}$ then the ${\left({\text{pK}}_{\text{a}}\right)}_1$ of ${\mathrm{H}}_2\mathrm{S}$ is

At $25°\mathrm{C}$ the enthalpy change, for the ionization of trichloroacetic acid is $+6.3 \mathrm{k}\mathrm{J} \mathrm{m}\mathrm{o}{\mathrm{l}}^{-1}$ and the entropy change, is $+0.0084 \mathrm{k}\mathrm{J} \mathrm{m}\mathrm{o}{\mathrm{l}}^{-1} {\mathrm{K}}^{-1}$. Then $\mathrm{p}\mathrm{K}\mathrm{a}$ of tricholoro acetic acid is