For the reaction A + B → C + D + q (kJ/mol), entropy change is positive. The reaction will be

1. Possible only at high temperature

2. Possible only at low temperature

3. Not possible at any temperature

4. Possible at any temperature

Consider the given reaction:
\(2 \mathrm{Cl}(\mathrm{~g}) \rightarrow \mathrm{Cl}_2(\mathrm{~g})\)

What are the values of \(∆\mathrm{H}\) and \(∆\mathrm{S}\), respectively?
1. \(\Delta \mathrm{H}=0, \Delta \mathrm{~S}=-\mathrm{ve}\)
2. \(\Delta \mathrm{H}=0, \Delta \mathrm{~S}=0\)
3. \(\Delta \mathrm{H}=-\mathrm{ve}, \Delta \mathrm{~S}=-\mathrm{ve}\)
4. \(\Delta \mathrm{H}=+\mathrm{ve}, \Delta \mathrm{~S}=+\mathrm{ve}\)

For the reaction:

The enthalpy change (ΔH) is 400 kJ mol⁻¹ and entropy change (ΔS) is 0.2 kJ K⁻¹ mol⁻¹.

Determine the temperature at which the reaction becomes spontaneous:

The entropy change in the surroundings when 1.00 mol of H₂O_(l) is formed under standard conditions is:

∆_fH^θ = –286 kJ mol^–1 

1. 952.5 J mol^-1

2. $979.7$ $\mathrm{J}$ ${\mathrm{mol}}^{-1}$

3. $949.7$ $\mathrm{J}$ ${\mathrm{mol}}^{-1}$

4. $959.7$ $\mathrm{J}$ ${\mathrm{mol}}^{-1}$