For the decomposition reaction of $N_2O_4$,$\Delta H = 58.04 \, kJ$ and $\Delta S = 176.7 \, J/K$. Calculate $\Delta G$ in $kJ$ at $T = 298 \, K$.

For the given reaction: $H_{2(g)} + S_{(s)} \to H_2S_{(g)}$; $\Delta H_r = 100\, kJ/mol$ and $\Delta S_r = 400\, J/mol\cdot K$. The temperature at which the above reaction occurs reversibly is (Assuming $\Delta H_r$ and $\Delta S_r$ are independent of temperature) ..... $K$.

$A$ process will be spontaneous at all temperatures if

For the vaporization of water at $1 \ atm$ pressure,the values of $\Delta H$ and $\Delta S$ are $40.63 \ kJ/mol$ and $108.8 \ J/K \cdot mol$ respectively. At what temperature $(K)$ will the change in Gibbs energy $(\Delta G)$ be zero (in $.4$)?

The value of ${\left[ {\frac{dG}{dT}} \right]_P}$ is ....

For the reaction $x + y \rightarrow z$,the process is spontaneous at room temperature,and the reverse reaction is spontaneous at high temperature. Which of the following is true?