Which of the following options correctly describes the free expansion of an ideal gas under adiabatic conditions?
1. \(\mathrm{q} \neq 0, \quad \Delta \mathrm{T}=0, \quad \mathrm{~W}=0\)
2. \(\mathrm{q}=0, \quad \Delta \mathrm{T}=0, \quad \mathrm{~W}=0\)
3. \(\mathrm{q}=0, \quad \Delta \mathrm{T}<0, \quad \mathrm{~W} \neq 0\)
4. \(\mathrm{q}=0, \quad \Delta \mathrm{T} \neq 0, \quad \mathrm{~W}=0\)

Enthalpy change for the reaction,

 4H(g) $\to$2H₂(g) is -869.6 kJ

The dissociation energy of H-H bond is

1. -869.6 kJ                 

2. + 434.8 kJ

3. +217.4 kJ                 

4. -434.8 kJ

Given the following bond energies:

Based on the data given above, enthalpy change for the following reaction will be:
 

1. 1523.6 kJ mol^-1

2. -243.6 kJ mol^-1

3. -120.0 kJ mol^-1

4. 553.0 kJ mol^-1

The values of $∆$H and $∆$S for the given reaction are 170 kJ and 170 JK^–1, respectively.
\(\mathrm{C} \text { (graphite) }+\mathrm{CO}_2(\mathrm{~g}) \rightarrow 2 \mathrm{CO}(\mathrm{g})\)
This reaction will be spontaneous at:

1. 710 K

2. 910 K

3. 1110 K

4. 510 K

Which of the following are not state functions?

(I) q + W                        (II) q

(III) W                           (IV) H-TS

1. (I) and (IV)                     

2. (II), (III) and (IV)

3. (I) , (II) and (III)               

4.  (II) and (III)

Consider the following reactions :

(i) ${\mathrm{H}}^{+}\left(\mathrm{aq}\right) + {\mathrm{OH}}^{-}\left(\mathrm{aq}\right) = {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) ∆\mathrm{H} = -{\mathrm{x}}_1 \mathrm{kJ} {\mathrm{mol}}^{-1}$

(ii) ${\mathrm{H}}_2\left(\mathrm{g}\right) + \frac{1}{2}{\mathrm{O}}_2\left(\mathrm{g}\right) = {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) ∆\mathrm{H} = -{\mathrm{x}}_2 \mathrm{kJ} {\mathrm{mol}}^{-1}$

(iii) ${\mathrm{CO}}_2\left(\mathrm{g}\right) + {\mathrm{H}}_2\left(\mathrm{g}\right) = \mathrm{CO}\left(\mathrm{g}\right) + {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) ∆\mathrm{H} = -{\mathrm{x}}_{3 }\mathrm{kJ} {\mathrm{mol}}^{-1}$

(iv) ${\mathrm{C}}_2{\mathrm{H}}_5\left(\mathrm{g}\right) + \frac{5}{2}{\mathrm{O}}_2\left(\mathrm{g}\right) = 2{\mathrm{CO}}_2\left(\mathrm{g}\right) + {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) ∆\mathrm{H} = -{\mathrm{x}}_4 \mathrm{kJ} {\mathrm{mol}}^{-1}$

Enthalpy of formation of H₂O(l) is:

1. -x₂ kJ mol^-1 

2. +x₃ kJ mol^-1

3. -x₄ kJ mol^-1

4. -x₁ kJ mol^-1

Identify the correct statement for change of Gibbs energy for a system ($∆$G_system) at constant temperature and pressure:

(1) If $∆$G_system > 0, the process is spontaneous

(2) If $∆$G_system = 0, the system has attained equilibrium

(3) If $∆$G_system = 0, the system is still moving in a particular direction

(4) If $∆$G_system < 0, the process is not spontaneous

The enthalpy and entropy change for the reaction :

Br₂ (l) + Cl₂ (g)$\to$ 2BrCl (g)

are 30 kJ mol^-1 and 105 J K^-1 mol^-1 respectively.

The temperature at which the reaction will be in equilibrium is :

The enthalpy of combustion of H₂, cyclohexene (C₆H₁₀) and cyclohexane (C₆H₁₂) are -241, -3800 and -3920 kJ per mol respectively. Heat of hydrogenation of cyclohexene is:

(a) -121 kJ per mol

(b) +121 kJ per mol

(c) +242 kJ per mol

(d) -242 kJ per mol

Consider the reactionat 300K

H₂₍₉₎ + Cl₂₍₉₎ →2HCI_(g), ΔH° = — 185 KJ

If 3 mole of H₂ completely react with 3 mol of Cl₂ to form Cl, $∆$U° of the reaction will be

(1) Zero

(2) –185 KJ

(3) -555 KJ

(4) None