As an isolated box, equally partitioned, contains two ideal gasses A and B as shown:

 

When the partition is removed, the gases mix. The changes in enthalpy $(∆\mathrm{H})$ and entropy $(∆\mathrm{S})$ in the process, respectively, are

1. Zero, positive

2. Zero, negative

3. Positive, zero

4. Negative, zero

The entropy change in the isothermal reversible expansion of 2 moles of an ideal gas from 10 to 100 L at 300 K is:
1. \(42.3 \mathrm{~J} \mathrm{~K}^{-1}\)
2. \(35.8 \mathrm{~J} \mathrm{~K}^{-1}\)
3. \(38.3 \mathrm{~J} \mathrm{~K}^{-1}\)
4. \(32.3 \mathrm{~J} \mathrm{~K}^{-1}\)

$∆{\mathrm{H}}_{\mathrm{f}}^0$ (298K) of methanol is given by the chemical equation:

1. \(\mathrm{C}(\text { diamond })+\frac{1}{2} \mathrm{O}_{2(\mathrm{~g})}+2 \mathrm{H}_{2(\mathrm{~g})} \rightarrow \mathrm{CH}_3 \mathrm{OH}_{(\mathrm{l})}\)

2. \(\mathrm{CH}_{4(\mathrm{~g})}+\frac{1}{2} \mathrm{O}_{2(\mathrm{~g})} \rightarrow \mathrm{CH}_3 \mathrm{OH}_{(\mathrm{g})}\)

3. \(\mathrm{CO}_{(\mathrm{g})}+2 \mathrm{H}_{2(\mathrm{~g})} \rightarrow \mathrm{CH}_3 \mathrm{OH}_{(\mathrm{l})}\)

4. \(\mathrm{C}(\text { graphite })+\frac{1}{2} \mathrm{O}_{2(\mathrm{~g})}+2 \mathrm{H}_{2(\mathrm{~g})} \rightarrow \mathrm{CH}_3 \mathrm{OH}_{(\mathrm{l})}\)

A piston filled with 0.04 mol of an ideal gas expands reversibly from 50.0 mL to 375 mL at a constant temperature of 37.0ºC. As it does so, it absorbs 208 J of heat. The values of q and w for the process will be-
(R = 8.314 J/mol K) (ln 7.5 = 2.01)

Entropy decreases during:

1. Crystallization of sucrose from solution

2. Rusting of iron

3. Melting of ice

4. Vaporization of camphor

Find the process showing negative entropy change:

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)

For the following given equations and $∆\mathrm{H}°$ values, determine the enthalpy of reaction at 298 K for the reaction:

C₂H₄(g) + 6F₂(g) $\to$ 2CF₄(g) + 4HF(g)

H₂(g) + F₂(g) $\to$ 2HF(g)       $∆{\mathrm{H}}_1^{°}$= -537 kJ

C(s) + 2F₂(g) $\to$CF₄(g)         $∆{\mathrm{H}}_2^{°}$=-680 kJ

2C(s) + 2H₂(g) $\to$C₂H₄(g)    $∆{\mathrm{H}}_3^{°}$= 52 kJ

1. –1165 kJ

2. –2486 kJ

3. +1165 kJ

4. +2486 kJ

When 4 g of iron is burnt in oxygen to form ferric oxide at constant pressure, 29.28 kJ of heat is evolved. Calculate the enthalpy of formation of ferric oxide.

(At. mass of Fe = 56) ?

1. $-$81.98 kJ

2. $-$ 819.8 kJ

3. $-$ 40.99 kJ

4.  +819.8 kJ

The values of ΔH and ΔS for the given reaction are 170 kJ and 170 JK^-1, respectively.
C_(graphite) + CO₂(g)→2CO(g) 
This reaction will be spontaneous at:

1. 710 K
2. 910 K
3. 1110 K
4. 510 K