For the reaction $A \rightleftharpoons B$,find the value of $log_{10}K$. Given: $\Delta_rH^o_{298\,K} = -54.07\, kJ\, mol^{-1}$,$\Delta_rS^o_{298\,K} = 10\, J\, K^{-1}\, mol^{-1}$,$R = 8.314\, J\, K^{-1}\, mol^{-1}$,$2.303 \times 8.314 \times 298 = 5705$.

Assertion: For every chemical reaction at equilibrium,the standard Gibbs energy change is zero.
Reason: At constant temperature and pressure,a chemical reaction is spontaneous in the direction of decreasing Gibbs energy.

Of the following reactions:
$(i) \, A \rightleftharpoons B, \Delta G^{\circ} = 250 \, kJ \, mol^{-1}$
$(ii) \, D \rightleftharpoons E, \Delta G^{\circ} = -100 \, kJ \, mol^{-1}$
$(iii) \, F \rightleftharpoons G, \Delta G^{\circ} = -150 \, kJ \, mol^{-1}$
$(iv) \, M \rightleftharpoons N, \Delta G^{\circ} = 150 \, kJ \, mol^{-1}$
The reaction with the largest equilibrium constant is:

Find the value of the equilibrium constant $(K)$ of a reaction at $300 \ K$, when standard Gibbs free energy change is $-25 \ kJ \ mol^{-1}$? (Consider $R = 8.33 \ J \ mol^{-1} \ K^{-1}$)

The equilibrium concentrations of the species in the reaction $A + B \rightleftharpoons C + D$ are $2, 3, 10$ and $6 \, mol \, L^{-1}$,respectively at $300 \, K$. $\Delta G^{\circ}$ for the reaction is $(R = 2 \, cal \, mol^{-1} \, K^{-1})$ (in $, cal$)

When a reaction is carried out at standard states,then at equilibrium: