Class 11
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Chemistry
Thermodynamics
NCERT Solutions
NCERT Solutions
Thermodynamics
22 Solutions
Q1
EXERCISES
Choose the correct answer. A thermodynamic state function is a quantity
(i)
used to determine heat changes
(ii)
whose value is independent of path
(iii)
used to determine pressure volume work
(iv)
whose value depends on temperature only.
Q2
EXERCISES
For the process to occur under adiabatic conditions, the correct condition is:
(i)
Δ
T
=
0
\Delta T=0
Δ
T
=
0
(ii)
Δ
p
=
0
\Delta p=0
Δ
p
=
0
(iii)
q
=
0
q=0
q
=
0
(iv)
w
=
0
\mathrm{w}=0
w
=
0
Q3
EXERCISES
The enthalpies of all elements in their standard states are:
(i)
unity
(ii)
zero
(iii)
< 0
(iv)
different for each element
Q4
EXERCISES
Δ
U
⊖
\Delta U^{\ominus}
Δ
U
⊖
of combustion of methane is
−
X
k
J
m
o
l
−
1
-\mathrm{X} \mathrm{kJ} \mathrm{mol}^{-1}
−
X
kJ
mol
−
1
. The value of
Δ
H
⊖
\Delta H^{\ominus}
Δ
H
⊖
is
(i)
=
Δ
U
⊖
=\Delta U^{\ominus}
=
Δ
U
⊖
(ii)
>
Δ
U
⊖
>\Delta U^{\ominus}
>
Δ
U
⊖
(iii)
<
Δ
U
⊖
<\Delta U^{\ominus}
<
Δ
U
⊖
(iv)
=
0
=0
=
0
Q5
EXERCISES
The enthalpy of combustion of methane, graphite and dihydrogen at 298 K are,
−
890.3
k
J
m
o
l
−
1
−
393.5
k
J
m
o
l
−
1
-890.3 \mathrm{~kJ} \mathrm{~mol}^{-1} -393.5 \mathrm{~kJ} \mathrm{~mol}^{-1}
−
890.3
kJ
mol
−
1
−
393.5
kJ
mol
−
1
, and
−
285.8
k
J
m
o
l
−
1
-285.8 \mathrm{~kJ} \mathrm{~mol}^{-1}
−
285.8
kJ
mol
−
1
respectively. Enthalpy of formation of
C
H
4
(
g
)
\mathrm{CH}_{4}(\mathrm{g})
CH
4
(
g
)
will be
(i)
−
74.8
k
J
m
o
l
−
1
\quad -74.8 \mathrm{~kJ} \mathrm{~mol}^{-1}
−
74.8
kJ
mol
−
1
(ii)
−
52.27
k
J
m
o
l
−
1
\quad -52.27 \mathrm{~kJ} \mathrm{~mol}^{-1}
−
52.27
kJ
mol
−
1
(iii)
+
74.8
k
J
m
o
l
−
1
\quad +74.8 \mathrm{~kJ} \mathrm{~mol}^{-1}
+
74.8
kJ
mol
−
1
(iv)
+
52.26
k
J
m
o
l
−
1
\quad +52.26 \mathrm{~kJ} \mathrm{~mol}^{-1}
+
52.26
kJ
mol
−
1
.
Q6
EXERCISES
A reaction,
A
+
B
→
C
+
D
+
q
\mathrm{A}+\mathrm{B} \rightarrow \mathrm{C}+\mathrm{D}+\mathrm{q}
A
+
B
→
C
+
D
+
q
is found to have a positive entropy change. The reaction will be
(i)
possible at high temperature
(ii)
possible only at low temperature
(iii)
not possible at any temperature
(iv)
possible at any temperature
Q7
EXERCISES
In a process, 701 J of heat is absorbed by a system and 394 J of work is done by the system. What is the change in internal energy for the process?
Q8
EXERCISES
The reaction of cyanamide,
N
H
2
C
N
(
s
)
\mathrm{NH}_{2} \mathrm{CN}(\mathrm{s})
NH
2
CN
(
s
)
, with dioxygen was carried out in a bomb calorimeter, and
Δ
U
\Delta U
Δ
U
was found to be
−
742.7
k
J
m
o
l
−
1
-742.7 \mathrm{~kJ} \mathrm{~mol}^{-1}
−
742.7
kJ
mol
−
1
at 298 K . Calculate enthalpy change for the reaction at 298 K .
N
H
2
C
N
(
s
)
+
3
2
O
2
(
g
)
→
N
2
(
g
)
+
C
O
2
(
g
)
+
H
2
O
(
l
)
\mathrm{NH}_{2} \mathrm{CN}(\mathrm{s})+\frac{3}{2} \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{N}_{2}(\mathrm{~g})+\mathrm{CO}_{2}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})
NH
2
CN
(
s
)
+
2
3
O
2
(
g
)
→
N
2
(
g
)
+
CO
2
(
g
)
+
H
2
O
(
l
)
Q9
EXERCISES
Calculate the number of kJ of heat necessary to raise the temperature of 60.0 g of aluminium from
35
∘
C
35^{\circ} \mathrm{C}
3
5
∘
C
to
55
∘
C
55^{\circ} \mathrm{C}
5
5
∘
C
. Molar heat capacity of Al is
24
J
m
o
l
−
1
K
−
1
24 \mathrm{~J} \mathrm{~mol}^{-1} \mathrm{~K}^{-1}
24
J
mol
−
1
K
−
1
.
Q10
EXERCISES
Calculate the enthalpy change on freezing of 1.0 mol of water at
10.0
∘
C
10.0^{\circ} \mathrm{C}
10.
0
∘
C
to ice at
−
10.0
∘
C
-10.0^{\circ} \mathrm{C}
−
10.
0
∘
C
.
Δ
fus
H
=
6.03
k
J
m
o
l
−
1
\Delta_{\text{fus}} H=6.03 \mathrm{~kJ} \mathrm{~mol}^{-1}
Δ
fus
H
=
6.03
kJ
mol
−
1
at
0
∘
C
0^{\circ} \mathrm{C}
0
∘
C
.
C
p
[
H
2
O
(
l
)
]
=
75.3
J
m
o
l
−
1
K
−
1
C_{p}[\mathrm{H}_{2} \mathrm{O}(\mathrm{l})]=75.3 \mathrm{~J} \mathrm{~mol}^{-1} \mathrm{~K}^{-1}
C
p
[
H
2
O
(
l
)]
=
75.3
J
mol
−
1
K
−
1
C
p
[
H
2
O
(
s
)
]
=
36.8
J
m
o
l
−
1
K
−
1
C_{p}[\mathrm{H}_{2} \mathrm{O}(\mathrm{s})]=36.8 \mathrm{~J} \mathrm{~mol}^{-1} \mathrm{~K}^{-1}
C
p
[
H
2
O
(
s
)]
=
36.8
J
mol
−
1
K
−
1
Q11
EXERCISES
Enthalpy of combustion of carbon to
C
O
2
\mathrm{CO}_{2}
CO
2
is
−
393.5
k
J
m
o
l
−
1
-393.5 \mathrm{~kJ} \mathrm{~mol}^{-1}
−
393.5
kJ
mol
−
1
. Calculate the heat released upon formation of 35.2 g of
C
O
2
\mathrm{CO}_{2}
CO
2
from carbon and dioxygen gas.
Q12
EXERCISES
Enthalpies of formation of
C
O
(
g
)
,
C
O
2
(
g
)
,
N
2
O
(
g
)
\mathrm{CO}(\mathrm{g}), \mathrm{CO}_{2}(\mathrm{g}), \mathrm{N}_{2} \mathrm{O}(\mathrm{g})
CO
(
g
)
,
CO
2
(
g
)
,
N
2
O
(
g
)
and
N
2
O
4
(
g
)
\mathrm{N}_{2} \mathrm{O}_{4}(\mathrm{g})
N
2
O
4
(
g
)
are
−
110
,
−
393
,
81
-110, -393, 81
−
110
,
−
393
,
81
and
9.7
k
J
m
o
l
−
1
9.7 \mathrm{~kJ} \mathrm{~mol}^{-1}
9.7
kJ
mol
−
1
respectively. Find the value of
Δ
r
H
\Delta_{r} H
Δ
r
H
for the reaction:
N
2
O
4
(
g
)
+
3
C
O
(
g
)
→
N
2
O
(
g
)
+
3
C
O
2
(
g
)
\mathrm{N}_{2} \mathrm{O}_{4}(\mathrm{~g})+3 \mathrm{CO}(\mathrm{~g}) \rightarrow \mathrm{N}_{2} \mathrm{O}(\mathrm{~g})+3 \mathrm{CO}_{2}(\mathrm{~g})
N
2
O
4
(
g
)
+
3
CO
(
g
)
→
N
2
O
(
g
)
+
3
CO
2
(
g
)
Q13
EXERCISES
Given
N
2
(
g
)
+
3
H
2
(
g
)
→
2
N
H
3
(
g
)
;
Δ
r
H
⊖
=
−
92.4
k
J
m
o
l
−
1
\mathrm{N}_{2}(\mathrm{~g})+3 \mathrm{H}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{NH}_{3}(\mathrm{~g}) ; \Delta_{r} H^{\ominus}=-92.4 \mathrm{~kJ} \mathrm{~mol}^{-1}
N
2
(
g
)
+
3
H
2
(
g
)
→
2
NH
3
(
g
)
;
Δ
r
H
⊖
=
−
92.4
kJ
mol
−
1
What is the standard enthalpy of formation of
N
H
3
\mathrm{NH}_{3}
NH
3
gas?
Q14
EXERCISES
Calculate the standard enthalpy of formation of
C
H
3
O
H
(
l
)
\mathrm{CH}_{3} \mathrm{OH}(\mathrm{l})
CH
3
OH
(
l
)
from the following data:
C
H
3
O
H
(
l
)
+
3
2
O
2
(
g
)
→
C
O
2
(
g
)
+
2
H
2
O
(
l
)
;
Δ
r
H
⊖
=
−
726
k
J
m
o
l
−
1
\mathrm{CH}_{3} \mathrm{OH}(\mathrm{l})+\frac{3}{2} \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{CO}_{2}(\mathrm{~g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) ; \Delta_{r} H^{\ominus}=-726 \mathrm{~kJ} \mathrm{~mol}^{-1}
CH
3
OH
(
l
)
+
2
3
O
2
(
g
)
→
CO
2
(
g
)
+
2
H
2
O
(
l
)
;
Δ
r
H
⊖
=
−
726
kJ
mol
−
1
C (graphite)
+
O
2
(
g
)
→
C
O
2
(
g
)
;
Δ
c
H
⊖
=
−
393
k
J
m
o
l
−
1
+\mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{CO}_{2}(\mathrm{~g}) ; \Delta_{c} H^{\ominus}=-393 \mathrm{~kJ} \mathrm{~mol}^{-1}
+
O
2
(
g
)
→
CO
2
(
g
)
;
Δ
c
H
⊖
=
−
393
kJ
mol
−
1
H
2
(
g
)
+
1
2
O
2
(
g
)
→
H
2
O
(
l
)
;
Δ
f
H
⊖
=
−
286
k
J
m
o
l
−
1
\mathrm{H}_{2}(\mathrm{~g})+\frac{1}{2} \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) ; \Delta_{f} H^{\ominus}=-286 \mathrm{~kJ} \mathrm{~mol}^{-1}
H
2
(
g
)
+
2
1
O
2
(
g
)
→
H
2
O
(
l
)
;
Δ
f
H
⊖
=
−
286
kJ
mol
−
1
.
Q15
EXERCISES
Calculate the enthalpy change for the process
C
C
l
4
(
g
)
→
C
(
g
)
+
4
C
l
(
g
)
\mathrm{CCl}_{4}(\mathrm{~g}) \rightarrow \mathrm{C}(\mathrm{g})+4 \mathrm{Cl}(\mathrm{g})
CCl
4
(
g
)
→
C
(
g
)
+
4
Cl
(
g
)
and calculate bond enthalpy of
C
−
C
l
\mathrm{C}-\mathrm{Cl}
C
−
Cl
in
C
C
l
4
(
g
)
\mathrm{CCl}_{4}(\mathrm{~g})
CCl
4
(
g
)
.
Δ
vap
H
⊖
(
C
C
l
4
)
=
30.5
k
J
m
o
l
−
1
\Delta_{\text{vap}} H^{\ominus}(\mathrm{CCl}_{4})=30.5 \mathrm{~kJ} \mathrm{~mol}^{-1}
Δ
vap
H
⊖
(
CCl
4
)
=
30.5
kJ
mol
−
1
.
Δ
f
H
⊖
(
C
C
l
4
)
=
−
135.5
k
J
m
o
l
−
1
\Delta_{f} H^{\ominus}(\mathrm{CCl}_{4})=-135.5 \mathrm{~kJ} \mathrm{~mol}^{-1}
Δ
f
H
⊖
(
CCl
4
)
=
−
135.5
kJ
mol
−
1
.
Δ
a
H
⊖
(
C
)
=
715.0
k
J
m
o
l
−
1
\Delta_{a} H^{\ominus}(\mathrm{C})=715.0 \mathrm{~kJ} \mathrm{~mol}^{-1}
Δ
a
H
⊖
(
C
)
=
715.0
kJ
mol
−
1
, where
Δ
a
H
⊖
\Delta_{a} H^{\ominus}
Δ
a
H
⊖
is enthalpy of atomisation
Δ
a
H
⊖
(
C
l
2
)
=
242
k
J
m
o
l
−
1
\Delta_{a} H^{\ominus}(\mathrm{Cl}_{2})=242 \mathrm{~kJ} \mathrm{~mol}^{-1}
Δ
a
H
⊖
(
Cl
2
)
=
242
kJ
mol
−
1
Q16
EXERCISES
For an isolated system,
Δ
U
=
0
\Delta U=0
Δ
U
=
0
, what will be
Δ
S
\Delta S
Δ
S
?
Q17
EXERCISES
For the reaction at 298 K ,
2
A
+
B
→
C
2 \mathrm{~A}+\mathrm{B} \rightarrow \mathrm{C}
2
A
+
B
→
C
Δ
H
=
400
k
J
m
o
l
−
1
\Delta H=400 \mathrm{~kJ} \mathrm{~mol}^{-1}
Δ
H
=
400
kJ
mol
−
1
and
Δ
S
=
0.2
k
J
K
−
1
m
o
l
−
1
\Delta S=0.2 \mathrm{~kJ} \mathrm{~K}^{-1} \mathrm{~mol}^{-1}
Δ
S
=
0.2
kJ
K
−
1
mol
−
1
At what temperature will the reaction become spontaneous considering
Δ
H
\Delta H
Δ
H
and
Δ
S
\Delta S
Δ
S
to be constant over the temperature range.
Q18
EXERCISES
For the reaction,
2
C
l
(
g
)
→
C
l
2
(
g
)
2 \mathrm{Cl}(\mathrm{g}) \rightarrow \mathrm{Cl}_{2}(\mathrm{g})
2
Cl
(
g
)
→
Cl
2
(
g
)
, what are the signs of
Δ
H
\Delta H
Δ
H
and
Δ
S
\Delta S
Δ
S
?
Q19
EXERCISES
For the reaction
2
A
(
g
)
+
B
(
g
)
→
2
D
(
g
)
2 \mathrm{~A}(\mathrm{~g})+\mathrm{B}(\mathrm{~g}) \rightarrow 2 \mathrm{D}(\mathrm{~g})
2
A
(
g
)
+
B
(
g
)
→
2
D
(
g
)
Δ
U
⊖
=
−
10.5
k
J
\Delta U^{\ominus}=-10.5 \mathrm{~kJ}
Δ
U
⊖
=
−
10.5
kJ
and
Δ
S
⊖
=
−
44.1
J
K
−
1
\Delta S^{\ominus}=-44.1 \mathrm{JK}^{-1}
Δ
S
⊖
=
−
44.1
JK
−
1
. Calculate
Δ
G
⊖
\Delta G^{\ominus}
Δ
G
⊖
for the reaction, and predict whether the reaction may occur spontaneously.
Q20
EXERCISES
The equilibrium constant for a reaction is 10. What will be the value of
Δ
G
⊖
\Delta G^{\ominus}
Δ
G
⊖
?
R
=
8.314
J
K
−
1
m
o
l
−
1
,
T
=
300
K
\mathrm{R}=8.314 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}, \mathrm{~T}=300 \mathrm{~K}
R
=
8.314
JK
−
1
mol
−
1
,
T
=
300
K
.
Q21
EXERCISES
Comment on the thermodynamic stability of
N
O
(
g
)
\mathrm{NO}(\mathrm{g})
NO
(
g
)
, given
1
2
N
2
(
g
)
+
1
2
O
2
(
g
)
→
N
O
(
g
)
;
Δ
r
H
⊖
=
90
k
J
m
o
l
−
1
\frac{1}{2} \mathrm{~N}_{2}(\mathrm{~g})+\frac{1}{2} \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{NO}(\mathrm{g}) ; \quad \Delta_{r} H^{\ominus}=90 \mathrm{~kJ} \mathrm{~mol}^{-1}
2
1
N
2
(
g
)
+
2
1
O
2
(
g
)
→
NO
(
g
)
;
Δ
r
H
⊖
=
90
kJ
mol
−
1
N
O
(
g
)
+
1
2
O
2
(
g
)
→
N
O
2
(
g
)
:
Δ
r
H
⊖
=
−
74
k
J
m
o
l
−
1
\mathrm{NO}(\mathrm{g})+\frac{1}{2} \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{NO}_{2}(\mathrm{~g}): \quad \Delta_{r} H^{\ominus}=-74 \mathrm{~kJ} \mathrm{~mol}^{-1}
NO
(
g
)
+
2
1
O
2
(
g
)
→
NO
2
(
g
)
:
Δ
r
H
⊖
=
−
74
kJ
mol
−
1
Q22
EXERCISES
Calculate the entropy change in surroundings when 1.00 mol of
H
2
O
(
l
)
\mathrm{H}_{2} \mathrm{O}(\mathrm{l})
H
2
O
(
l
)
is formed under standard conditions.
Δ
f
H
⊖
=
−
286
k
J
m
o
l
−
1
\Delta_{f} H^{\ominus}=-286 \mathrm{~kJ} \mathrm{~mol}^{-1}
Δ
f
H
⊖
=
−
286
kJ
mol
−
1
.
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