Given the following reaction:

$N_2{H}_4\left(g\right)\underset{<mspace\; width="1em"></mspace><mspace\; width="1em"></mspace><mspace\; width="1em"></mspace><mspace\; width="1em"></mspace><mspace\; width="1em"></mspace><mspace\; width="1em"></mspace><mspace\; width="1em"></mspace><mspace\; width="1em"></mspace><mspace\; width="1em"></mspace><mspace\; width="1em"></mspace>}{\to }{N}_2{H}_2\left(g\right)+H_2\left(g\right);$

[Given:  ∆ H^∘ = 109 kJ/mol, B.E. of (N-N) = 163 kJ/mol, B.E. of (N-H) = 391 kJ/mol and B.E. of (H-H) = 436 kJ/mol]

Calculate the bond enthalpy of N = N:

1. 182 kJ/mol

2. 400 kJ/mol

3. 300 kJ/mol

4. 218 kJ/mol

The pair of isochoric among the transformation of state is:

1. K to L and L to M

2. L to M and N to K

3. L to M and M to N

4. M to N and N to K

Reversible expansion of an ideal gas under isothermal and adiabatic conditions are shown in the figure:

AB$\to$Isothermal expansion

AC$\to$Adiabatic expansion

Which of the following options is not correct?

What is the amount of work done by an ideal gas, if the gas expands isothermally from \(10^{-3}~m^3\) to \(10^{-2}~m^3\) at \(300~K\)against a constant pressure of \(10^{5}~Nm^{-2}\)?

Thermodynamics is not concerned about:

Consider the following reactions: 

Enthalpy of formation of H₂O(l) is :

1. $-x_3$ $kJ$ $mo{l}^{-1}$

2. $-x_4$ $kJ$ $mo{l}^{-1}$

3. $-x_1$ $kJ$ $mo{l}^{-1}$

4. $-x_2$ $kJ$ $mo{l}^{-1}$

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:

If for a certain reaction ${∆}_{r}H$ is 30 kJ mol^–1 at 450 K, the value of ${∆}_{r}S$ (in JK^–1 mol^–1) for which the same reaction will be spontaneous at the same temperature is:

Identify the correct statement regarding entropy:

Following reaction occurs at ${25}^{\circ }C$ :

2NO_(g) , (1 × 10^-5   atm) + Cl₂(g), ( 1×10^-2 atm) ⇌ 2NOCl_(g), ( 1×10^{- 2} atm) ∆G^o   is-

1. -45.65 kJ

2. -28.53 kJ

3. -22.82 kJ

4. -57.06 kJ