For a reaction,$\Delta H = -40 \, kJ$ and $\Delta S = -50 \, J/K$. At what temperature range will it change from spontaneous to non-spontaneous?

For a chemical reaction $A^{+} + B \rightleftharpoons C^{+} + D$ $(\Delta_{r} H^{0} = 80 \, kJ \, mol^{-1})$,the entropy change $\Delta_{r} S^{0}$ depends on the temperature $T$ (in $K$) as $\Delta_{r} S^{0} = 2T \, J \, K^{-1} \, mol^{-1}$.
Minimum temperature at which it will become spontaneous is ..... $K$. (Integer)

For the change $H_2O_{(l)} \to H_2O_{(g)}$ at $P = 1 \ atm$ and $T = 373 \ K$,the free energy change $\Delta G = 0$. This indicates that:

At the transition temperature $T$, $\Delta G^0 = 0$ and $\Delta G^0 = 105 - 35 \log T$, where $A$ and $B$ are two states of substance $X$. The transition temperature in $^\circ\text{C}$ when pressure is $1 \text{ atm}$ is . . . . . . .

Calculate the value of $\Delta G$ for the following reaction: $N_2O_{4(g)} \longrightarrow 2NO_{2(g)}$ if $\Delta H = 57.44 \ kJ$ and $\Delta S = 176 \ J \ K^{-1} \ mol^{-1}$ at $300 \ K$. (in $kJ$)

Which of the following formulas is correct for Gibbs free energy $(G)$?