What is the resulting change in the internal energy of the system when 50 calories are added to
the system and the system does work of 30 calories on the surroundings?

1. 20 cal                             
2. 50 cal
3. 40 cal                             
4. 30 cal

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

 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

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

Change in entropy is negative for:

1. Bromine (l)$\to$Bromine(g)

2. C(s) + H₂O(g) $\to$CO(g) + H₂(g)

3. N₂(g,10 atm)$\to$N₂(g,1 atm) 

4. Fe ( 1mol, 400 K) $\to$ Fe( 1mol, 300 K)

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

The maximum work done in expanding 16 g of 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

For the process

H₂O(l) $\to$H₂O(g)

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

1. $∆{\mathrm{S}}_{\mathrm{system}}>0<mspace\; width="1em"></mspace>\mathrm{and}<mspace\; width="1em"></mspace>∆{\mathrm{S}}_{\mathrm{surrounding}}>0$

2. $∆{\mathrm{S}}_{\mathrm{system}}>0<mspace\; width="1em"></mspace>\mathrm{and}<mspace\; width="1em"></mspace>∆{\mathrm{S}}_{\mathrm{surrounding}}<0$

3. $∆{\mathrm{S}}_{\mathrm{system}}<0<mspace\; width="1em"></mspace>\mathrm{and}<mspace\; width="1em"></mspace>∆{\mathrm{S}}_{\mathrm{surrounding}}>0$

4. $∆{\mathrm{S}}_{\mathrm{system}}<0<mspace\; width="1em"></mspace>\mathrm{and}<mspace\; width="1em"></mspace>∆{\mathrm{S}}_{\mathrm{surrounding}}<0$

For the process:

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

1. $∆\mathrm{G}=0;<mspace\; width="1em"></mspace>∆\mathrm{S}=+\mathrm{ve}$

2. $∆\mathrm{G}=0;<mspace\; width="1em"></mspace>∆\mathrm{S}=-\mathrm{ve}$

3. $∆\mathrm{G}=+\mathrm{ve};<mspace\; width="1em"></mspace>∆\mathrm{S}=0$

4. $∆\mathrm{G}=-\mathrm{ve};<mspace\; width="1em"></mspace>∆\mathrm{S}=+\mathrm{ve}$