The maximum work done in expanding 16g ${\mathrm{O}}_2$ isothermally at 300 K and occupying a volume of 5 dm³ until the volume becomes 25 dm³ is-

1. –2.01$\times {10}^3$ J

2. 2.01×${10}^{-3}$ J

3. +2.81$\times {10}^3$ J

4. +2.01$\times {10}^{-6}$J

Which of the following conditions will always lead to a non-spontaneous change?

1. Positive $∆H$ and positive $∆S$

2. Negative$∆H$ and negative $∆S$

3. Positive $∆H$ and negative $∆S$

4. Negative $∆H$ and positive $∆P$

To calculate the amount of work done in joules during reversible isothermal expansion of an ideal gas, the volume must be expressed in:

For the conversion C_(graphite) $\to$C_(diamond) the $∆S$ is:

1. Zero

2. Positive

3. Negative

4. Unknown

The entropy change in the conversion of one mole of liquid water at 373 K to vapour at the same temperature would be: (Latent heat of vaporization of water, $∆H_{vap}=2.257$ $kJ/g$)

1. 105.9 JK^-1mol^-1

2. 107.9 JK^-1mol^-1

3. 108.9 JK^-1mol^-1

4. 109.9 JK^-1mol^-1

\(\underset {-10.7 \mathrm{JK}^{-1}} {\mathrm{H}_{(a q)}^{+}+\mathrm{OH^{-}}_{(a q)}} \rightarrow \underset {+70 \mathrm{JK}^{-1} \mathrm{~mol}^{-1} } {\mathrm{H}_{2} \mathrm{O}_{(l)}} ~~ \mathrm{S}^{\circ}(298 \mathrm{~K})\)

Standard entropy change for the above reaction is:

1. 60.3 JK^-1 mol^-1

2. 80.7 JK^-1 mol^-1

3. -70 JK^-1 mol^-1

4. +10.7 JK^-1 mol^-1

Hydrolysis of sucrose is given by the following reaction

Sucrose + H₂O $\rightleftharpoons$ Glucose + Fructose

If the equilibrium constant (K_c) is 2$\times$10¹³ at 300 K, the value of ${∆}_r{G}^{⊝}$ at the same temperature will be:

1. 8.314 J mol^–1 K^–1$\times$300 K$\times$ln (2$\times$10¹³)

2. 8.314 J mol^–1 K^–1$\times$300 K$\times$ln (3$\times$10¹³)

3. –8.314 J mol^–1 K^–1$\times$300 K$\times$ln (4$\times$10¹³)

4. –8.314 J mol^–1 K^–1$\times$300 K$\times$ln (2$\times$10¹³)

For the reaction, 2Cl(g) $\to$ Cl₂(g), the correct option is:

1. ${∆}_{\mathrm{r}}\mathrm{H}>0 \mathrm{and} {∆}_{\mathrm{r}}\mathrm{S}<0$

2. ${∆}_{\mathrm{r}}\mathrm{H}<0 \mathrm{and} {∆}_{\mathrm{r}}\mathrm{S}>0$

3. ${∆}_{\mathrm{r}}\mathrm{H}<0 \mathrm{and} {∆}_{\mathrm{r}}\mathrm{S}<0$

4. ${∆}_{\mathrm{r}}\mathrm{H}>0 \mathrm{and} {∆}_{\mathrm{r}}\mathrm{S}>0$

When 4 g of iron is burnt to ferric oxide at a constant pressure, 29.28 kJ of heat is evolved.

The enthalpy of formation of ferric oxide will be-

(At. mass of Fe = 56) ?

1. $-$81.98 kJ

2. $-$ 819.8 kJ

3. $-$ 40.99 kJ

4.  +819.8 kJ

The difference in $∆H$ $and$ $∆U$ for combustion of benzoic acid at 300 K is equal to-

1. -1.247 kJ

2. +1.247 kJ

3. -1.247 J

4. +1.247 J