The $Δ H$ and $Δ S$ for a reaction at one atmospheric pressure are +30.558 kJ and 0.066 k J $k^{- 1}$ respectively. The temperature at which the free energy change will be zero and below this temperature the nature of reaction would be:

483 K, spontaneous
443 K, non-spontaneous
443 K, spontaneous
463 K, non-spontaneous

View Solution

Q. For the reaction

C O (g) + \frac{1}{2} O_{2} (g) \to C O_{2} (g), Δ H

and

Δ S

are

- 283 kJ

and

- 87 {J K}^{- 1} {mol}^{- 1}

respectively. It was intended to carry out this reaction at

1000, 1500, 3000

and

3500 K

. At which of these temperatures would this reaction be thermodynamically spontaneous?

$1500 a n d 3500 K$
$3000 a n d 3500 K$
$1000, 1500 a n d 3000 K$
$1500, 3000 a n d 3500 K$

View Solution

Q. For the process

H_{2} O (l) ⇌ H_{2} O (g) (1 b a r, 373 K),

the correct set of thermodynamic parameters is

$△ G = 0, △ S = + v e$
$△ G = 0, △ S = - v e$
$△ G = + v e, △ S = 0$
$△ G = - v e, △ S = + v e$

View Solution

Q. What can be concluded about the values of

△ H

and

△ S

from this graph?

$△ H > 0, △ S > 0$
$△ H > 0, △ S < 0$
$△ H < 0, △ S > 0$
$△ H < 0, △ S < 0$

View Solution

Q. Silane,

S i H_{4}

like methane, burns in air. The product silica (silicon dioxide) is solid quite unlike carbondioxide.

S i H_{4 (g)} + 2 O_{2 (g)} \to S i O_{2} (s) + 2 H_{2} O_{(g)}

standard Gibbs energy of formation of

S i O_{2} (s)

H_{2} O

and

S i H_{4}

are 805, – 228.6 and + 52.3 (all is KJ

m o l^{- 1}

) respectively. Calculate the value of

Δ_{r} G

$- 1314.5 K J m o l^{- 1}$
$- 1413 K J m o l^{- 1}$
$1134.5 K J m o l^{- 1}$
$3411.5 K J m o l^{- 1}$

View Solution

Q. Calculate free energy change for the following reaction at

298 K

2 N O (g) + B r_{2} (l) \to 2 N O B r (g)

Given the partial pressure of

N O

0.1 atm

and the partial pressure of

N O B r

2.0 a t m

and

Δ G_{f N O B r}^{0} = 82.4 {kJ mol}^{- 1}, Δ G_{N O}^{0} = 86.55 {kJ mol}^{- 1}

$6.5 kJ$
$- 8.3 kJ$
$- 1.4 kJ$
$+ 8.3 kJ$

View Solution

Given: For ethylene oxide

Δ H_{f}^{\circ}

and

Δ S^{\circ}

are

- 51 {kJ mol}^{- 1}

and

243 {J mol}^{- 1} K^{- 1}

respectively and acetaldehyde,

Δ H_{f}^{\circ}

and

Δ S^{\circ}

are

- 166 {kJ mol}^{- 1}

and

266 {J mol}^{- 1} K^{- 1}

respectively, then select the correct statements :

$Δ H^{\circ}$ per the reaction is $- 115 {kJ mol}^{- 1}$
$Δ S^{\circ} = 0.023 {kJ mol}^{- 1} K^{- 1}$
Reaction is thermodynamically favourable
Low temperature can favour formation of more product

View Solution

Q. For the process

H_{2} O (l) ⇌ H_{2} O (g) (1 b a r, 373 K),

the correct set of thermodynamic parameters is

$△ G = 0, △ S = + v e$
$△ G = 0, △ S = - v e$
$△ G = + v e, △ S = 0$
$△ G = - v e, △ S = + v e$

View Solution

Q. The enthalpy change for a given reaction at

298 K

- x J

m o l^{- 1}

(x being positive). If the reaction occurs spontaneously at

298 K

, the entropy change at that temperautre

can be negative but numerically larger than x/298
can be negative but numerically smaller than x/298
cannot be negative
cannot be positive

View Solution

Q. Zinc reacts with dilute hydrochloric acid to give hydrogen gas at

17^{o} C .

The enthalpy change of the reaction is

- 12.55 k J m o l^{- 1}

and entropy change equals to

5.0 J K^{- 1} m o l^{- 1}

for the reaction. Calculate the free energy change and predict whether the reaction is spontaneous or not.

$- 14.00 k J m o l^{- 1}$ , spontaneous
$- 18.00 k J m o l^{- 1}$ , spontaneous
$14.00 k J m o l^{- 1}$ , non-spontaneous
$18.00 k J m o l^{- 1}$ , non-spontaneous

View Solution

Q. For a given reaction:

x A + y B \to m C + n D

△ H = - 30 k J m o l^{- 1}, △ S = - 100 J K^{- 1} m o l^{- 1}

At what temperature the reaction is at equilibrium?

$50^{o} C$
$250^{o} C$
$100^{o} C$
$27^{o} C$

View Solution

△ H a n d △ S

for the reaction,

A g_{2} O (s) \to 2 A g (s) + \frac{1}{2} O_{2} (g)

are

30.56 k J m o l^{- 1}

and

66.0 J m o l^{- 1} K^{- 1}

respectively. Calculate the temperature at which free energy change for the reaction will be zero. Predict whether the reaction will be sponteneous above or below this temperature (in K)

$563 K$ , reaction will be sponteneous above this temperature.
$463 K$ , reaction will be sponteneous above this temperature.
$563 K$ , reaction will be sponteneous below this temperature.
$463 K$ , reaction will be sponteneous below this temperature.

View Solution

Q. The standard reaction Gibbs energy for a chemical reaction at an absolute temperature T is given

△_{r} G^{\circ} = A - B T,

where A and B are non-zero constants. Which of the following is true about this reaction?

Exothermic if A > 0 and B < 0
Endothermic if A < 0 and B > 0
Endothermic if A > 0
Exothermic if B < 0

View Solution

Q. Predict whether the reaction

C O (g) + (\frac{1}{2}) O_{2} (g) \to C O_{2} (g)

at 300 K is spontaneous & exothermic or non-spontaneous & endothermic, when the standard entropy change is

- 0.094 k J m o l^{- 1} K^{- 1} .

The standard Gibb's free energies of formation for

C O_{2}

and

C O

are

- 394.4 a n d - 137.2 k J m o l^{- 1}

respectively.

value of $△ H^{o}$ is negative so the given reaction is non-spontaneous and endothermic.
value of $△ H^{o}$ is positive so the given reaction is non-spontaneous and endothermic.
value of $△ H^{o}$ is negative so the given reaction is spontaneous and exothermic.
value of $△ H^{o}$ is positive so the given reaction is spontaneous and exothermic.

View Solution

Q. The standard enthalpies of formation of

C O_{2} (g), H_{2} O (l)

and glucose(s) at

25^{\circ} C

are – 400 kJ/mol, - 300 kJ/mol and – 1300 kJ/mol, respectively. The standard enthalpy of combusion per gram of glucose at

25^{\circ} C

+ 2900kJ
– 2900 kJ
– 16.11 kJ
+ 16.11 kJ

View Solution

When is $Δ H$ negative and when is $Δ S$ negative?

View Solution

Q. Assume

△ H^{o} a n d △ S^{o}

to be independent of temperature, at what temperature will the reaction given below becomes spontaneous?

N_{2} (g) + O_{2} (g) \to 2 N O (g)

S_{N_{2} (g)}^{o} = 191.4 J K^{- 1} m o l^{- 1}

S_{O_{2} (g)}^{o} = 204.9 J K^{- 1} m o l^{- 1}

S_{N O (g)}^{o} = 210.5 J K^{- 1} m o l^{- 1}

△ H^{o} = 180.8 k J m o l^{- 1}

7320 K
3650 K
6595 K
8500 K

View Solution

For a reaction at $25^{\circ} C$ enthalpy change ( $△ H$ ) and entropy change ( $△ S$ ) are -11.7 K J $m o l^{- 1}$ and -105 J $m o l^{- 1}$ $K^{- 1}$ , respectively. It is a spontaneous reaction

True
False

View Solution

Q. The change in Gibb's free energy,

Δ G = - W_{n o n e x p a n s i o n}

True
False

View Solution

Q. For a reaction at equilibrium, the value of Gibbs free energy change is:

> 0
< 0
= 0
$\neq$ 0

View Solution

Q. The free energy change under standard conditions

(G)

is related to the equilibrium constant by the van't Hoff isotherm:

G = - R T l n K

. But what does the symbol

R

represent?

The entropy of the system
The gas constant
The Avogadro constant
The reaction quotient

View Solution

Q. Which is correct about

Δ G

$Δ G = Δ H - T Δ S$
At equilibrium, $Δ G^{o} = 0$
At equilibrium, $Δ G = - R T l o g K$
$Δ G = Δ G^{o} + R T l o g K$

View Solution

Q. What is the free energy change

(△ G)

when 1.0 mole of water at

100^{o} C

and 1 atmospheric pressure is converted into steam at

100^{o} C

and 1 atmospheric pressure?

80 cal
540 cal
620 cal
zero

View Solution

Q. The molar entropy content of 1 mole of oxygen

(O_{2})

gas at

300 K

and

1 atm

250 {J mol}^{- 1} K^{- 1}

. Calculate

Δ G

when 1 mole of oxygen is expanded reversibly and isothermally from

300 K

1 atm

to double its volume (Take

R = 8.314 {J mol}^{- 1} K^{- 1}, log e = 2.303

)

$1.728 {kJ mol}^{- 1} K^{- 1}$
$0$
$- 1.728 {kJ mol}^{- 1} K^{- 1}$
$0.75 {kJ mol}^{- 1} K^{- 1}$

View Solution

Q. For the equilibrium

N_{2} O_{4} ⇌ 2 N O_{2}

(G_{N_{2} O_{4}}^{o})_{298 K} = 100 k J m o l^{- 1} a n d (G_{N O_{2}}^{o})_{298 K} = 50 k J m o l^{- 1} .

When 5 mole/litre of each is taken, calculate the value of

△ G

k J m o l^{- 1}

for the reaction at 298 K.(nearest integer)

View Solution