Gibbs Energy Change and Equilibrium

Match list-I (Equations) with List-II (Type of processes) and select the correct option.
 

For the reaction $2{\mathrm{O}}_3\to 3{\mathrm{O}}_2, \mathrm{\Delta H}, \mathrm{\Delta S}$ and $∆\mathrm{G}$ are respectively:

Products are favoured in a chemical reaction taking place at a constant temperature and pressure. Consider the following statements:
(i) The change in Gibbs energy for the reaction is negative.
(ii) The total change in Gibbs energy for the reaction and the surroundings is negative.
(iii) The change in entropy for the reaction is positive.
(iv) The total change in entropy for the reaction and the surroundings is positive.

The statements which are ALWAYS true are:-

For the reaction $2{\mathrm{O}}_3\to 3{\mathrm{O}}_2, \mathrm{\Delta H}, \mathrm{\Delta S}$ and $∆\mathrm{G}$ are respectively:

The value of ${\log }_{10}\mathrm{K}$ for a reaction $\mathrm{A}\rightleftharpoons \mathrm{B}$ is

(Given ${\mathrm{\Delta }}_{\mathrm{r}}{{\mathrm{H}}^{\mathrm{o}}}_{298\mathrm{K}}=-54.07 {\mathrm{kJmol}}^{-1}, {\mathrm{\Delta }}_{\mathrm{r}}{{\mathrm{S}}^{\mathrm{o}}}_{298\mathrm{K}}=10 {\mathrm{JK}}^{-1}{\mathrm{mol}}^{-1}$ and $\mathrm{R}=8.314 {\mathrm{JK}}^{-1}{\mathrm{mol}}^{-1}; 2.303\times 8.314\times 298=5705$)

Data is given for some reaction: $∆\mathrm{H} = -110\mathrm{KJ}$ and $∆\mathrm{S} = 40 {\mathrm{JK}}^{-1}$ at $400 \mathrm{K}$.

Choose the correct option for the reaction.

For the reaction:

${\mathrm{N}}_2{\mathrm{O}}_4\left(\mathrm{g}\right)\to 2{\mathrm{NO}}_2\left(\mathrm{g}\right)$

$(∆{\mathrm{H}}^{\circ }=+57.24 \mathrm{kJ}, ∆{\mathrm{S}}^{\circ }= 175.8 {\mathrm{Jk}}^{–1})$

At what temperature the reaction will be spontaneous?

Find the equilibrium constant for the formation of ${\mathrm{NO}}_2$ from $\mathrm{NO}$ and ${\mathrm{O}}_2$ at $298 \mathrm{K}$ ,as per the reaction given below. Given ${\mathrm{\Delta }}_{\mathrm{f}}{\mathrm{G}}^0\left({\mathrm{NO}}_2\right)=52 \mathrm{kJ}.{\mathrm{mol}}^{-1}, {\mathrm{\Delta }}_{\mathrm{f}}{\mathrm{G}}^0\left(\mathrm{NO}\right)=87 \mathrm{kJ}.{\mathrm{mol}}^{-1}$
${\mathrm{\Delta }}_{\mathrm{f}}{\mathrm{G}}^0\left({\mathrm{O}}_2\right)=0 \mathrm{kJ}.{\mathrm{mol}}^{-1}$ (Antilog of $0.122=1.3210$)

A container holds $2$ moles of gas $\mathrm{A}$ and $3$ moles of gas $\mathrm{B}$ in it at a constant pressure of $2 \mathrm{bar}$. The gases are separated by a wall and the whole container is held in an ice bath at a constant temperature of $0°\mathrm{C}$. The wall is removed and the gases mix and react to form $\mathrm{C}$ by the reaction given below. After the reaction comes to equilibrium how many grams of the ice have melted ? Consider that all heat produced in reaction has been utilized in melting of ice.

$\text{A}\left(\text{g}\right)+\text{B}\left(\text{g}\right)\rightleftharpoons \text{C}\left(\text{g}\right)$

All data above is at $298 \mathrm{K}$. $\Delta {}_{\text{FUS}}{\text{H}}^{\circ }\left(\text{water, 0}\mathrm{℃}\right)={\text{6 kJ mol}}^{-1}$.

Read the passage and answer the following questions as single choice correct .$\left(\text{Given 273}\times \text{8.314}\times \text{2.303}={\text{5227 J m}}^{-1}\right)$

The value of log ${\mathrm{K}}_{\mathrm{p}}$for the reaction may be

The efficiency of the fuel cell will be

A reaction occurs spontaneously if

If $\mathrm{K}<1.0$, what will be the value of $∆\mathrm{G}°$ out of the following ?