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)

1.	q = +208 J, w = -208 J	2.	q = -208 J, w = -208 J
3.	q = -208 J, w = + 208 J	4.	q = +208 J, w = + 208 J

Entropy decreases during:

1. Crystallization of sucrose from solution

2. Rusting of iron

3. Melting of ice

4. Vaporization of camphor

At a temperature of 300K, what is the entropy change for the reaction given below?

2H₂ (g) + O₂ (g) $\to$2H₂O(l)

Standard entropies of H₂ (g), O₂(g) and H₂O(l) are 126.6, 201.20 and 68.0 JK^-1mol^-1 respectively.

1. -318.4 JK^-1mol^-1                           

2. 318.4 JK^-1mol^-1 

3. 31.84 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)

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 

2. -2486

3. +1165

4. +2486

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

1. C(diamond)+$\frac{1}{2}{O}_2$+2H₂$\to$CH₃OH

2. CH₄+$\frac{1}{2}{O}_2$$\to$CH₃OH

3. CO+2H₂$\to$CH₃OH

4. C(graphite)+$\frac{1}{2}{O}_2$+2H₂$\to$CH₃OH

An ideal gas expands isothermally from ${10}^{-3}{m}^3 to {10}^{-2} m^3$ at 300 K against a constant pressure of ${10}^5 N{m}^{-2}$. The work done by  the gas is:

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. $32.3 \mathrm{J} {\mathrm{K}}^{-1}$

The $∆\mathrm{H}$ for vaporization of a liquid is \(20 \mathrm{~kJ} / \mathrm{mol}.\) Assuming ideal behaviour, the change in internal energy for the vaporization of \(1 \mathrm{~mol}\) of the liquid at \(60^{\circ} \mathrm{C}\) and 1 bar is close to:

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