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

Identify which of the following is the correct option for free expansion of an ideal gas under adiabatic condition.

1.  $\mathrm{ \Delta q}\ne 0,$ $∆\mathrm{T}=0,$ $\mathrm{W}=0$
2.  $\mathrm{ \Delta q}=0,$ $∆\mathrm{T}=0,$ $\mathrm{W}=0$
3.  $\mathrm{ \Delta q}=0,$ $∆\mathrm{T}<0,$ $\mathrm{W}\ne 0$
4.  $\mathrm{ \Delta q}=0,$ $∆\mathrm{T}\ne 0,$ $\mathrm{W}=0$

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

Standard entropies of X₂, Y₂ and XY₃ are 60, 40 and 50JK^-1mol^-1 respectively. For the reaction 
$\frac{1}{2}{\mathrm{X}}_2$ $+$ $\frac{3}{2}{\mathrm{Y}}_2$ $$ $\leftrightharpoons$ ${\mathrm{XY}}_3$ $;$ $∆\mathrm{H}$ $=$ $-30$ $\mathrm{kJ}$ to be at equilibrium, the temperature should be:

1. 750 K

2. 1000 K

3. 1250 K

4. 500 K

From the following bond energies:
H—H bond energy: 431.37 kJ mol^-1 
C=C bond energy: 606.10 kJ mol^-1 
C—C bond energy: 336.49 kJ mol^-1 
C—H bond energy: 410.50 kJ mol^-1 
Enthalpy for the reaction, 

will be:

The values of ΔH and ΔS for the given reaction are 170 kJ and 170 JK^-1, respectively.
C_(graphite) + CO₂(g)→2CO(g) 
This reaction will be spontaneous at:

1. 710 K
2. 910 K
3. 1110 K
4. 510 K