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:

1.	70	2.	–33
3.	33	4.	–70

At standard conditions, if the change in the enthalpy for the following reaction is –109 kJ mol^–1 
H_2(g)+Br_2(g)$\to$2HBr_(g) and the bond energy of H₂ and Br₂ is 435 kJ mol^–1 and 192 kJ mol^–1 respectively, what is the bond energy (in kJ mol^–1) of HBr?

Which of the following options correctly represents the relationship between \(C_p \text { and } C_V\) for one mole of an ideal gas?

For irreversible expansion of an ideal gas under isothermal conditions, the correct option is :

1. $∆\mathrm{U}=0,<mspace\; width="1em"></mspace>∆{\mathrm{S}}_{\mathrm{total}}\ne 0$

2. $∆\mathrm{U}\ne 0,<mspace\; width="1em"></mspace>∆{\mathrm{S}}_{\mathrm{total}}=0$

3. $∆\mathrm{U}=0,<mspace\; width="1em"></mspace>∆{\mathrm{S}}_{\mathrm{total}}=0$

4. $∆\mathrm{U}\ne 0,<mspace\; width="1em"></mspace>∆{\mathrm{S}}_{\mathrm{total}}\ne 0$

The correct option for free expansion of an ideal gas under adiabatic condition is:

1.  $q=0,$ $∆T<0$ and $w>0$

2.  $q<0,$ $∆T=0$ and $w=0$

3.  $q>0,$ $∆T>0$ and $w>0$

4.  $q=0,$ $∆T=0$ and $w=0$

Hydrolysis of sucrose is given by the following reaction

Sucrose + H₂O $\rightleftharpoons$ Glucose + Fructose

If the equilibrium constant (K_c) is 2$\times$10¹³ at 300 K, the value of ${∆}_r{G}^{⊝}$ at the same temperature will be:

1. 8.314 J mol^–1 K^–1$\times$300 K$\times$ln (2$\times$10¹³)

2. 8.314 J mol^–1 K^–1$\times$300 K$\times$ln (3$\times$10¹³)

3. –8.314 J mol^–1 K^–1$\times$300 K$\times$ln (4$\times$10¹³)

4. –8.314 J mol^–1 K^–1$\times$300 K$\times$ln (2$\times$10¹³)

The work done when 1 mole of gas expands reversibly and isothermally from a pressure of 5 atm to 1 atm at 300 K is:
[Given: log 5 = 0.6989 and R = 8.314 J K^-1 mol^-1​]