Thermodynamics Part 2: Hess's Law, 2nd & 3rd Law of Thermodynamics and Gibb's Free Energy

1.

The entropy change in the fusion of one mole of a solid melting at 27 $°$ C is:
(the latent heat of fusion is 2930 J mol^-1)
1. 9.77 JK^-1mol^-1
2. 19.73 JK^-1mol^-1
3. 2930 JK^-1mol^-1
4. 108.5 JK^-1mol^-1

2.

For exothermic reaction to be spontaneous ( $∆ S$ =negative) temperature must be -

1.	High	2.	Zero
3.	Constant	4.	Low

3.

Which of the following is not a state function?

1.	ΔS	2.	ΔG
3.	ΔH	4.	ΔQ

4.

The unit of entropy is -
1. J mol^–1
2. JK mol^–1
3. J mol^–1 K^–1
4. J^–1 K^–1 mol^–1

5.

The enthalpy of hydration of Na⁺(g) and Cl^-(g) ions are -406 kJ mol^-1 and -364 kJ mol^-1 respectively.
The enthalpy of the solution of NaCl(s) is:
$\text{The lattice energy of NaCl is}~ 780~{ kJ mol}^{-1}. $
1. 23 kJ mol^-1
2. 10 kJ mol^-1
3. -10 kJ mol^-1
4. -82 kJ mol^-1

6.

Which of the following thermodynamic quantities is an outcome of the second law of thermodynamics?
1. Work
2. Enthalpy
3. Internal energy
4. Entropy

7.

The correct statement for a reversible process in a state of equilibrium is:
1. $∆$ G = – 2.30RT log K
2. $∆$ G = 2.30RT log K
3. $∆$ G^o = – 2.30RT log K
4. $∆$ G^o = 2.30RT log K

8.

For the reaction:
$\mathrm{X}_2 \mathrm{O}_4(l) \rightarrow 2 \mathrm{XO}_2(g)$
with the given values $\Delta U = 2.1 \, \text{kcal}$ and $\Delta S = 20 \, \text{cal K}^{-1}$ at $300 \, \text{K}$, what is the value of $\Delta G$?
1. +2.7 kcal
2. –2.7 kcal
3. +9.3 kcal
4. –9.3 kcal

9.

Give the following bond energies:
H—H bond energy: 431.37 kJ mol^-1
C=C bond energy: 606.10 kJ mol^-1
C—C bond energy: 336.49 kJ mol^-1
C—H bond energy: 410.50 kJ mol^-1
Enthalpy for the following reaction will be:

1. 1523.6 kJ mol^-1
2. -243.6 kJ mol^-1
3. -120.0 kJ mol^-1
4. 553.0 kJ mol^-1

10.

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

11.

The enthalpy of combustion of H₂, cyclohexene (C₆H₁₀) and cyclohexane (C₆H₁₂) are -241, -3800 and -3920 kJ per mol respectively. Heat of hydrogenation of cyclohexene is:
1. -121 kJ per mol
2. +121 kJ per mol
3. +242 kJ per mol
4. -242 kJ per mol

12.

Entropy decreases during:
1. Crystallization of sucrose from solution
2. Rusting of iron
3. Melting of ice
4. Vaporization of camphor

13.

For the following given equations and $∆ H °$ 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) $∆ H_{1}^{°}$ = -537 kJ
C(s) + 2F₂(g) $\to$ CF₄(g) $∆ H_{2}^{°}$ =-680 kJ
2C(s) + 2H₂(g) $\to$ C₂H₄(g) $∆ H_{3}^{°}$ = 52 kJ
1. –1165 kJ
2. –2486 kJ
3. +1165 kJ
4. +2486 kJ

14.

The enthalpy and entropy change for the reaction :
Br₂(l)+Cl₂(g) $\to$ 2BrCl(g) are 30 kJ mol^-1 and 105 JK^-1 mol^-1 respectively.
The temperature at which the reaction will be in equilibrium is :
1. 285.7 K
2. 273 K
3. 450 K
4. 300 K

15. For the reaction, 2Cl(g)

\to

Cl₂(g), the correct option is:

1.	$\Delta_{\mathrm{r}} \mathrm{H}>0$ and $\Delta_{\mathrm{r}} \mathrm{S}<0 $	2.	$\Delta_{\mathrm{r}} \mathrm{H}<0$ and $ \Delta_{\mathrm{r}} \mathrm{S}>0 $
3.	$\Delta_{\mathrm{r}} \mathrm{H}<0 $ and $\Delta_{\mathrm{r}} \mathrm{S}<0 $	4.	$\Delta_{\mathrm{r}} \mathrm{H}>0$ and $ \Delta_{\mathrm{r}} \mathrm{S}>0$

16.

The free energy change $(∆ G °)$ is negative when -
1. The surroundings do no electrical work on the system.
2. The surroundings do electrical work on the system.
3. The system does electrical work on the surroundings.
4. The system does no electrical work on the surroundings.

17.

What happens to the specific heat of a gas when its volume is reduced to half of its initial value?

1.	Reduce to half	2.	Be Doubled
3.	Remain constant	4.	Increase four times

18.

The entropy change can be calculated by using the expression $∆ S = \frac{q_{rev}}{T}$ . When water freezes in a glass beaker, the correct statement among the following is:

1.	∆ S (system) decreases but ∆ S (surroundings) remains the same.
2.	∆ S (system) increases but ∆ S (surroundings) decreases.
3.	∆ S (system) decreases but ∆ S (surroundings) increases.
4.	∆ S (system) decreases but ∆ S (surroundings) also decreases.

19.

Which of the following algebraic relationships is correct based on the given thermochemical equations?

(i)	C(graphite) + O₂(g) → CO₂(g) ; ∆_rH = x kJmol^-1
(ii)	C(graphite) + $\frac{1}{2}$O₂(g) → CO(g) ; ∆_rH = y kJ mol^-1
(iii)	CO(g) + $\frac{1}{2}$O₂(g) → CO₂(g); ∆_rH = z kJ mol^-1

1. z=x+y
2. x=y-z
3. x=y+z
4. y=2z-x

20.

Which of the following is not correct?
1. $∆ G$ is zero for a reversible reaction.
2. $∆ G$ is positive for a spontaneous reaction.
3. $∆ G$ is negative for a spontaneous reaction.
4. $∆ G$ is positive for a non-spontaneous reaction.

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1.	\(\Delta_{\mathrm{r}} \mathrm{H}>0\) and \(\Delta_{\mathrm{r}} \mathrm{S}<0 \)	2.	\(\Delta_{\mathrm{r}} \mathrm{H}<0\) and \( \Delta_{\mathrm{r}} \mathrm{S}>0 \)
3.	\(\Delta_{\mathrm{r}} \mathrm{H}<0 \) and \(\Delta_{\mathrm{r}} \mathrm{S}<0 \)	4.	\(\Delta_{\mathrm{r}} \mathrm{H}>0\) and \( \Delta_{\mathrm{r}} \mathrm{S}>0\)

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