The standard change is Gibbs energy for the reaction,

H₂O$\rightleftarrows$H⁺ + OH^- at 25$°$C is:

1. 100 kJ                         

2. -90 kJ

3. 90 kJ                           

4. -100 kJ

The entropy change in the fusion of one mole of a solid melting at 27$°$C  is  -       

(the latent heat of fusion is 2930 J mol^-1)                                                                                   

1. 9.77 JK^-1mol^-1               

2. 19.73 JK^-1mol^-1               

3. 2930 JK^-1mol^-1               

4. 108.5 JK^-1mol^-1          

The maximum work done in expanding 16 g oxygen at 300 K and occupying a volume of 5 dm³ isothermally until the volume becomes 25 dm³ is:

1. 2.01 x 10³ J                       

2. +2.81 x 10³ J

3. 2.01 x 10^-3 J                     

4. +2.01 x 10^-6 J

1 mole of an ideal gas at 25$°\mathrm{C}$ is subjected to expand reversibly ten times of its initial volume. The change in entropy of expansion is:

1. 19.15 JK^-1mol^-1                         

2. 16.15 JK^-1mol^-1

3. 22.15 JK^-1mol^-1                         

4. None of the above

For the process

H₂O(l) $\to$H₂O(g)

 at T=100$°$C and 1 atmosphere pressure, the correct choice is:

1. $∆{\mathrm{S}}_{\mathrm{system}}>0 \mathrm{and} ∆{\mathrm{S}}_{\mathrm{surrounding}}>0$

2. $∆{\mathrm{S}}_{\mathrm{system}}>0 \mathrm{and} ∆{\mathrm{S}}_{\mathrm{surrounding}}<0$

3. $∆{\mathrm{S}}_{\mathrm{system}}<0 \mathrm{and} ∆{\mathrm{S}}_{\mathrm{surrounding}}>0$

4. $∆{\mathrm{S}}_{\mathrm{system}}<0 \mathrm{and} ∆{\mathrm{S}}_{\mathrm{surrounding}}<0$

 During an adiabatic process:

1. pressure is maintained constant

2. gas is isothermally expanded

3. there is perfect heat insulation

4. the system changes heat with surroundings

Heat of combustion $∆\mathrm{H}°$ for C(s), H₂(g) and CH₄(g) are -94, -68 and -213 kcal/mol. Then, $∆\mathrm{H}°$ for 

C(s) + 2H₂(g) $\to$CH₄(g) is                                                                 

1. -17 kcal/mol                         

2. -111 kcal/mol

3. -170 kcal/mol                       

4. -85 kcal/mol

If 50 calorie are added to a system and system does work of 30 calorie on surroundings, the change in internal energy of system is:

1. 20 cal                             

2. 50 cal

3. 40 cal                             

4. 30 cal

For the process:

H₂O(l)[1 bar, 373 K] $\rightleftharpoons$H₂O(g)[1 bar, 373 K] the correct set of thermodynamic parameters are:

1. $∆\mathrm{G}=0; ∆\mathrm{S}=+\mathrm{ve}$

2. $∆\mathrm{G}=0; ∆\mathrm{S}=-\mathrm{ve}$

3. $∆\mathrm{G}=+\mathrm{ve}; ∆\mathrm{S}=0$

4. $∆\mathrm{G}=-\mathrm{ve}; ∆\mathrm{S}=+\mathrm{ve}$

The internal energy change when a system goes from state A to B is 40 kJ/mol. If the system goes from A to B by a reversible path and returns to state A by an irreversible path. What would be the change in internal energy?

1. 40 kJ                                   

2. >40 kJ

3. <40 kJ                                 

4. Zero