The correct thermodynamic conditions for a spontaneous reaction at all temperatures is:
1. $∆$H > 0 and $∆$S< 0
2. $∆$H < 0 and $∆$S> 0
3. $∆$H < 0 and $∆$S< 0
4. $∆$H > 0 and $∆$S = 0

The heat of combustion of carbon to CO₂ is –393.5 KJ/mol. The heat released upon the formation of 35.2 g of CO₂ from carbon and oxygen gas is:
1. –315 KJ
2. +315 KJ
3. –630 KJ
4. +630 KJ

The correct statement for a reversible process in a state of equilibrium is:
1. $∆$G = – 2.30RT log K
2. $∆$G = 2.30RT log K
3. $∆$G^o = – 2.30RT log K
4. $∆$G^o = 2.30RT log K

For the reaction:
\(\mathrm{X}_2 \mathrm{O}_4(l) \rightarrow 2 \mathrm{XO}_2(g)\)
with the given values \(\Delta U = 2.1 \, \text{kcal}\) and \(\Delta S = 20 \, \text{cal K}^{-1}\) at \(300 \, \text{K}\), what is the value of \(\Delta G\)?

1. +2.7 kcal
2. –2.7 kcal
3. +9.3 kcal
4. –9.3 kcal

In which of the following reactions, the standard reaction entropy change
$(∆S^0)$ is positive, and standard Gibb's energy change
$(∆G^0)$ decreases sharply with increasing temperature?

The enthalpy of fusion of water is 1.435 kcal/mol. The molar entropy change for the melting of ice at 0 ^oC is:

1. 10.52 cal/(mol K)

2. 21.04 cal/(mol K)

3. 5.260 cal/(mol K)

4. 0.526 cal/(mol K)

The standard enthalpy of vaporization ${∆}_{\mathrm{vap}}{\mathrm{H}}^{\mathrm{o}}$ for water at 100 ^oC is 40.66 kJ mol^–1. The internal energy of vaporization of water at 100 ^oC (in kJ mol^–1) is: 
(Assume water vapour behaves like an ideal gas.)

1. +37.56

2. –43.76

3. +43.76

4. +40.66

Given the following reaction:
\(4H(g)\)→  \(2 H_{2}\)\((g)\)
The enthalpy change for the reaction is -869.6 kJ. The dissociation energy of the H-H bond is:
1. -869.6 kJ
2. +434.8kJ
3. +217.4kJ
4. -434.8 kJ