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

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

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

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 mathematical form of the first law of thermodynamics when heat (q) is supplied and W is work done by the system (-ve) is:

1. $∆$U=q+W                       

2. $∆$U=q-W

3. $∆$U=-q+W                       

4. $∆$U= -q-W

Which statements are correct?

1. $2.303 \log \frac{{\mathrm{P}}_2}{{\mathrm{P}}_1}=\frac{∆{\mathrm{H}}_{\mathrm{vap}.}}{\mathrm{R}}\frac{\left[{\mathrm{T}}_2-{\mathrm{T}}_1\right]}{{\mathrm{T}}_1{\mathrm{T}}_2}$ is called Clausius-Clapeyron equation

2. $\frac{∆{\mathrm{H}}_{\mathrm{vap}.}}{\mathrm{Boiling} \mathrm{Point}}=88 J mo{l}^{-1}{K}^{-1}$ is called Trouton's rule

3. Entropy is a measure of unavailable energy, i.e.,

     unavailable energy = entropy x temperature

4. All of the above

$∆\mathrm{S}°$ will be highest for the reaction:

1. Ca(s) + 1/2 O₂(g) $\to$CaO(s)

2. CaCO₃(s)$\to$CaO(s) + CO₂(g)

3. C(s) + O₂(g)$\to$CO₂(g)

4. N₂(g) + O₂(g)$\to$2NO(g)