The entropy change in the surroundings when 1.00 mol of H2O(l) is formed under standard conditions is-
∆fHθ = –286 kJ mol–1
1. 952.5 J mol-1
2.
3.
4.
The thermodynamic stability of NO(g) based on the above data is:
1. Less than NO2(g)
2. More than NO2(g)
3. Equal to NO2(g)
4. Insufficient data
The equilibrium constant for a reaction is 10. The value of will be:
( )
For the reaction 2A(g) + B(g) → 2D(g) ; ∆U° = - 10.5 kJ and ∆S° = - 44.1 J K-1, the value of ∆G° for the given reaction would be-
1. 1.6 J
2. -0.16 kJ
3. 0.16 kJ
4. 1.6 kJ
For the reaction at 298 K,
2A + B → C
ΔH = 400 kJ mol−1 and ΔS = 0.2 kJ K−1 mol−1. The reaction will become spontaneous at-
1. 1500 K
2. 2000 K
3. 100 K
4. 1900K
For an isolated system with ∆U = 0, the ∆S value will be-
1. | Positive | 2. | Negative |
3. | Zero | 4. | Not possible to define |
The standard enthalpy of the formation of CH3OH(l) from the following data is:
\(\small{\mathrm{CH}_3 \mathrm{OH}_{(l)}+\frac{3}{2} \mathrm{O}_2(\mathrm{g}) \rightarrow \mathrm{CO}_2(\mathrm{g})+2 \mathrm{H}_2 \mathrm{O}_{(l)} \text {; }}\) \( \Delta_{\mathrm{r}} \mathrm{H}^{\circ}=-726 \mathrm{~kJ} \mathrm{~mol}{ }^{-1}\) |
\(\small{\mathrm{C}(\mathrm{s})+\mathrm{O}_2(\mathrm{g}) \rightarrow \mathrm{CO}_2(\mathrm{g}) \text {; } }\) \(\Delta_{\mathrm{c}} \mathrm{H}^{\circ}=-393 \mathrm{~kJ} \mathrm{~mol}{ }^{-1}\) |
\(\small{\mathrm{H}_{2(\mathrm{g})}+\frac{1}{2} \mathrm{O}_{2(\mathrm{g})} \rightarrow \mathrm{H}_2 \mathrm{O}_{(l)} \text {; } } \) \(\Delta_{\mathrm{f}} \mathrm{H}^{\circ}=-286 \mathrm{~kJ} \mathrm{~mol}^{-1}\) |
1. | −239 kJ mol−1 | 2. | +239 kJ mol−1 |
3. | −47 kJ mol−1 | 4. | +47 kJ mol−1 |
. The standard enthalpy of formation of gas in the above reaction would be-
1. | -92.4 J (mol)-1 | 2. | -46.2 kJ (mol)-1 |
3. | +46.2 J (mol)-1 | 4. | +92.4 kJ (mol)-1 |
The enthalpy of formation of are
–110 kJ , – 393 kJ , 81 kJ and 9.7 kJ respectively.
The value of for the reaction would be-