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Equilibrium C...

Equilibrium Constant and Standard Free Energy Change

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Q.

Aluminum reacts with sulfuric acid to form aluminum sulfate and hydrogen. What is the volume of hydrogen gas in liters (L) produced at $300 K$ and $1.0 atm$ pressure, when $5.4 g$ of aluminum and $50.0 mL$ of $5.0 M$ sulfuric acid are combined for the reaction?

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Q. The standard electrode potential of a Daniell cell is

1.1 V

. Find the standard Gibbs energy for the reaction.

Z n (s) + C u^{2 +} (a q .) \to Z n^{2 +} (a q .) + C u (s)

$- 212.3 kJ/mol$
$- 2.123 kJ/mol$
$- 21.23 kJ/mol$
$- 21230 kJ/mol$

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Q.

Give some examples of second-order reactions.

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Q. Which of the following equilibriums degree of dissociation is independent of initial concentration?

$H I ⇌ \frac{1}{2} H_{2} + \frac{1}{2} I_{2}$
$P C l_{5} ⇌ P C l_{3} + C l_{2}$
$N_{2} O_{4} ⇌ 2 N O_{2}$
all the above

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Q.

Sucrose hydrolyses in acid solution into glucose and fructose following first order rate law with a half life of 3.33 h at 25°C. After 9 h, the fraction of sucrose remaining is f. The value of log10 1 f 1f is ......x _10^-2. (Rounded off to the nearest integer) [Assume : ln 10 = 2.303, ln 2 = 0.693]

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Q.

A B_{3} (g)

is dissociated as

A B_{3} (g) ⇌ A B_{2} (g) + \frac{1}{2} B_{2} (g)

.
When the initial pressure of

A B_{2}

is 800 torr and the total pressure developed at equlibrium is

900 t o r r

. What fraction of

A B_{3} (g)

is dissociated?

10%
25%
20%
30%

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Q. Calculate free energy change for the conversion of oxygen to ozone at 298 K, if

K_{p}

for the change

\frac{3}{2} O_{2} (g) \to O_{3} (g)

at 298 K is

2.47 \times 10^{- 29} (a t m)^{- \frac{1}{2}} .

log 2.47=0.3926

163.23 kJ
170.14 KJ
159.24 KJ
160.00 KJ

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Q. For a reaction

A ⇌ P

, the plots of [A] and [P] with time temperatrues

T_{1}

and

T_{2}

are given below.

If

T_{2} > T_{1}

, the correct statement(s) is (are) (Assume

Δ H^{θ}

and

Δ S^{θ}

are independent of temperature and ratio of ln K at

T_{1}

to lnK at

T_{2}

is greater than

\frac{T_{2}}{T_{1}}

. Here H, S, G and K are enthalpy, entropy, Gibbs energy and equilibrium constant, respectively.)

$Δ H^{θ} < 0, Δ S^{θ} < 0$
$Δ G^{θ} < 0, Δ H^{θ} > 0$
$Δ G^{θ} < 0, Δ S^{θ} < 0$
$Δ G^{θ} < 0, Δ S^{θ} > 0$

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Q.

Calculate

$Δ G^{0}$

the equilibrium constant for the formation of $N O_{2}$ from NO and $O_{2}$ at 298 K
$N O (g) + \frac{1}{2} O_{2} (g) \leftrightarrow N O_{2} (g)$

Where:
$Δ_{f} G^{\circ} (N O_{2}) = 52.0 k J / m o l Δ_{f} G^{\circ} (N O) = 87.0 k J / m o l Δ_{f} G^{\circ} (O_{2}) = 0 k J / m o l$

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Q. If

0.2 m o l

of

H_{2} (g)

and

2 m o l

of

S (s)

are mixed in a

1 d m^{3}

vessel at

90^{o} C,

the partial pressure of

H_{2} S (g)

formed according to the reaction would be :

H_{2} (g) + S (s) ⇌ H_{2} S; K_{P} = 6.8 \times 10^{- 2}

0.19 atm
0.38 atm
0.6 atm
0.072 atm

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Q. The standard reduction potentials for

Z n^{2 +} / Z n, N i^{2 +} / N i

and

F e^{2 +} / F e

are

- 0.74 V, - 0.22 V and - 0.44 V

respectively. The reaction

X + Y^{2 +} ⟶ X^{2 +} + Y

will be spontaneous when

$X = N i, Y = F e$
$X = N i, Y = Z n$
$X = F e, Y = Z n$
$X = Z n, Y = N i$

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Q. The following equilibrium constants are given:

N_{2} (g) + H_{2} (g) ⇌ 2 N H_{3} (g); K_{1}

N_{2} (g) + O_{2} (g) ⇌ 2 N O (g); K_{2}

H_{2} (g) + \frac{1}{2} O_{2} (g) ⇌ H_{2} O (g); K_{3}

The equilibrium constant for the oxidation of the

N H_{3}

by oxygen to give

N O

is:

2 N H_{3} (g) + \frac{5}{2} O_{2} (g) ⇌ 2 N O (g) + 3 H_{2} O (g)

$\frac{K_{1} K_{2}}{K_{3}}$
$\frac{K_{2} K_{3}^{3}}{K_{1}}$
$\frac{K_{2} K_{3}^{2}}{K_{1}}$
$\frac{K_{2}^{2} K_{3}}{K_{1}}$

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Q. For the melting of ice at

25^{\circ} C

, the enthalpy of fusion is

6.97 K J m o l^{- 1}

and entropy of fusion is

25.4 J K^{- 1} g^{- 1}

. The free energy change and nature of the reaction at this temperature is :

– 599.2 J , spontaneous
14.54 K J , non spontaneous
– 14.54 KJ , spontaneous
+ 599.2 J , non spontaneous

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Q. Thermal decomposition of gaseous

X_{2}

to gaseous X at 298 K takes place according to the following equation:

X_{2} (g) ⇌ 2 X (g)

The standard reaction Gibbs energy,

Δ_{r} G^{0},

of this reaction is positive. At the start of the reaction, there is one mole of

X_{2}

and no X. As the reaction proceeds, the number of moles of X formed is given by

β

. Thus,

β_{e q u i l i b r i u m}

is the number of moles of X formed at equilibrium. The reaction is carried out at a constant total pressure of 2 bar. Consider the gases to behave ideally. (Given: R = 0.083 L bar

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

).

The INCORRECT statement among the following, for this reaction is:

Decrease in the total pressure will result in formation of more moles of gaseous X
At the start of the reaction, dissociation of gaseous $X_{2}$ takes place spontaneously
$β_{e q u i l i b r i u m} = 0.7$
$K_{c} < 1$

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Q. For the reaction

2 S O_{2} (g) + O_{2} (g) ⇌ 2 S O_{3} (g)

The incorrect statement is :

Decreasing the pressure will shift the equilibrium to the left
Addition of inert gas at constant pressure will shift the equilibrium towards right
Value of $\frac{K_{P}}{K_{C}}$ is less than one
No change in state of equilibrium when inert gas is added at constant volume

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Q. Which of the following will favour the formation of

N H_{3}

by Haber's process?

Decrease in temperature
Increase in pressure
Addition of catalyst
All of the above

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Q.

For the reaction $N H_{4} H S_{(g)} ⇌ N H_{3 (g)} + H_{2} S_{(g)}$ in a closed flask, the equilibrium pressure is P atm. The standard free energy of the reaction would be :

-2 RT (In p - In 2)
-RT In p
-RT(In p - In 2)
-2 RT In p

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Q.

The reaction N $_{2}$ O $_{4}$ (g) $⇌$ 2NO $_{2}$ (g) is
carried out at 298 K and 20 bar. Five moles of each of N $_{2}$ O $_{4}$ and
NO $_{2}$ , are taken initially. Given: $Δ$ $_{f}$ G $^{\circ}$
(N $_{2}$ O $_{4}$ ) = 100 kJ mol $^{- 1}$ and $Δ$ $_{f}$ G $^{\circ}$
(NO $_{2}$ ) = 50 kJ mol $^{- 1}$ .Choose the appropriate option: The values of
$Δ$ G and K $^{\circ}$ $_{p}$ at 298 K are:

$Δ$ G = 0, $K_{p}^{\circ}$ = 10 bar; The reaction is in equilibrium
5.7058 kJ $m o l^{- 1}$ , $K_{p}^{\circ}$ = 1 bar; The reaction goes in the reverse direction
-5.7058 kJ $m o l^{- 1}$ , $K_{p}^{\circ}$ = 1 bar; The reaction goes in the forward direction
$Δ$ G = 0, $K_{p}^{\circ}$ = 1 bar; The reaction is in equilibrium

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Q. Consider the binding of oxygen to haemoglobin:

H b_{a q} + O_{2 (g)} ⇌ H b - O_{(a q)} .

If we decrease the pressure of the system, what is the effect on this equilibrium reaction?

The equilibrium shifts to the left
The equilibrium shifts to the right
There is no change in the system
None of the above

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Q.

Why there is an increase in mass when a substance is oxidized?

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Q. The standard free energy change of a reaction is

△ G^{o} = - 115 k J a t 298 K .

Calculate the equilibrium constant

K_{P}

in

l o g K_{P} . (R = 8.314 J K^{- 1} m o l^{- 1})

20.16
2.303
2.016
13.83

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Q.

At a given temperature, the equilibrium constant for reaction

$P C l_{5} (g) ⇌ P C l_{3} (g) + C l_{2} (g)$ is $2.4 \times 10^{- 3}$ . At the same temperature, the

equilibrium constant for reaction $P C l_{3} (g) + C l_{2} (g ⇌ (P C l_{5} (g)$ is

2.4 x 10^-3
-2.4 x 10^-3
4.2 x 10²
4.8 x 10^-2

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Q. The value of

Δ G^{\circ}

for a reaction having K=1 would be:

0
+RT
-RT
-1

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Q.

The half-life of $Radium - 226$ is $1590 years$ . If a sample contains $300 mg$ , how many $m g$ will remain after $3000 years$ ?

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Q. The equilibrium constant

K_{p}

for the reaction

A ⇌ 2 B

is related to degree of dissociation

α

of A and pressure P as:

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Q. For which of the following reactions at equilibrium at constant temperature doubling the volume will cause a shift to the right ( favour forward reaction)?

$N_{2} O_{4} (g) ⇌ 2 N O_{2} (g)$
$2 C O (g) + O_{2} ⇌ 2 C O_{2} (g)$
All of these
$N_{2} (g) + O_{2} (g) ⇌ 2 N O (g)$

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Q. Which of following graphs correctly depicts the variation of Gibbs Free Energy with extent of the Reaction ?

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Q. Calculate the equilibrium concentration ratio of

C t o A

, if

2

moles each of

A

and

B

are allowed to come to equilibrium at

300 K

.

A (g) + B (g) ⇌ C (g) + D (g)

Δ G^{\circ} = - 1380 c a l / m o l

$0.56$
$4.45$
$1.50$
$3.16$

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Q. What will be the

E^{0}

value for the given half cell reaction?

S n^{4 +} (a q) + 4 e^{-} \to S n (s)

Given:

E^{0}

value for the half cell reaction,

S n^{4 +} (a q) + 2 e^{-} \to S n^{2 +} (a q); E^{0} = 0.15 V

S n^{2 +} (a q) + 2 e^{-} \to S n (s); E^{0} = - 0.14 V

$+ 0.001 V$
$+ 0.005 V$
$- 0.001 V$
$- 0.005 V$

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Q. Thermal decomposition of gaseous

X_{2}

to gaseous X at 298 K takes place according to the following equation:

X_{2} (g) ⇌ 2 X (g)

The standard reaction Gibbs energy,

Δ G^{o},

of this reaction is positive. At the start of the reaction, there is one moe of

X_{2}

and no X. As the reaction proceeds, the number of moles of X formed is given by

β,

Thus,

β_{e q u i l i b r i u m}

is the number of moles of X formed at equilibrium. The reaction is carried out at a constant total pressure of 2 bar. Consider the gases to behave ideally.
(Given: R = 0.083 L bar

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

)

The Incorrect statement among the following, for this reaction is

$K_{c} < 1$
At the start of the reaction, dissociation of gaseous $X_{2}$ takes place spontaneously
Decrease in the total pressure will result in formation of more moles of gaseous X
$β_{e q u i l i b r i u m}$ = 0.7

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