Class 11
Dashboard
Chemistry
Redox Reactions
NCERT Solutions
NCERT Solutions
Redox Reactions
40 Solutions
Exercise:
All Exercises
EXERCISES
In-text Problems
Q1
EXERCISES
Assign oxidation number to the underlined elements in each of the following species:
(a)
N
a
H
2
P
‾
O
4
\mathrm{NaH}_{2} \underline{\mathrm{P}} \mathrm{O}_{4}
NaH
2
P
O
4
(b)
N
a
H
S
‾
O
4
\mathrm{NaH} \underline{\mathrm{S}} \mathrm{O}_{4}
NaH
S
O
4
(c)
H
4
P
‾
2
O
7
\mathrm{H}_{4} \underline{\mathrm{P}}_{2} \mathrm{O}_{7}
H
4
P
2
O
7
(d)
K
2
M
n
‾
O
4
\mathrm{K}_{2} \underline{\mathrm{Mn}} \mathrm{O}_{4}
K
2
Mn
O
4
(e)
C
a
O
‾
2
\mathrm{Ca} \underline{\mathrm{O}}_{2}
Ca
O
2
(f)
N
a
B
‾
H
4
\mathrm{Na} \underline{\mathrm{B}} \mathrm{H}_{4}
Na
B
H
4
(g)
H
2
S
‾
2
O
7
\mathrm{H}_{2} \underline{\mathrm{S}}_{2} \mathrm{O}_{7}
H
2
S
2
O
7
(h)
K
A
l
(
S
‾
O
4
)
2
⋅
12
H
2
O
\mathrm{KAl}(\underline{\mathrm{S}} \mathrm{O}_{4})_{2} \cdot 12 \mathrm{H}_{2} \mathrm{O}
KAl
(
S
O
4
)
2
⋅
12
H
2
O
Q2
EXERCISES
What are the oxidation number of the underlined elements in each of the following and how do you rationalise your results ?
(a)
K
I
‾
3
\mathrm{K}\underline{\mathrm{I}}_{3}
K
I
3
(b)
H
2
S
‾
4
O
6
\mathrm{H}_{2}\underline{\mathrm{S}}_{4}\mathrm{O}_{6}
H
2
S
4
O
6
(c)
F
e
‾
3
O
4
\underline{\mathrm{Fe}}_{3}\mathrm{O}_{4}
Fe
3
O
4
(d)
C
‾
H
3
C
‾
H
2
O
H
\underline{\mathrm{C}}\mathrm{H}_{3}\underline{\mathrm{C}}\mathrm{H}_{2}\mathrm{OH}
C
H
3
C
H
2
OH
(e)
C
‾
H
3
C
‾
O
O
H
\underline{\mathrm{C}}\mathrm{H}_{3}\underline{\mathrm{C}}\mathrm{OOH}
C
H
3
C
OOH
Q3
EXERCISES
Justify that the following reactions are redox reactions:
(a)
C
u
O
(
s
)
+
H
2
(
g
)
→
C
u
(
s
)
+
H
2
O
(
g
)
\mathrm{CuO}(\mathrm{s})+\mathrm{H}_{2}(\mathrm{~g}) \rightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{~g})
CuO
(
s
)
+
H
2
(
g
)
→
Cu
(
s
)
+
H
2
O
(
g
)
(b)
F
e
2
O
3
(
s
)
+
3
C
O
(
g
)
→
2
F
e
(
s
)
+
3
C
O
2
(
g
)
\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+3 \mathrm{CO}(\mathrm{~g}) \rightarrow 2 \mathrm{Fe}(\mathrm{s})+3 \mathrm{CO}_{2}(\mathrm{~g})
Fe
2
O
3
(
s
)
+
3
CO
(
g
)
→
2
Fe
(
s
)
+
3
CO
2
(
g
)
(c)
4
B
C
l
3
(
g
)
+
3
L
i
A
l
H
4
(
s
)
→
2
B
2
H
6
(
g
)
+
3
L
i
C
l
(
s
)
+
3
A
l
C
l
3
(
s
)
4 \mathrm{BCl}_{3}(\mathrm{~g})+3 \mathrm{LiAlH}_{4}(\mathrm{~s}) \rightarrow 2 \mathrm{B}_{2} \mathrm{H}_{6}(\mathrm{~g})+3 \mathrm{LiCl}(\mathrm{s})+3 \mathrm{AlCl}_{3}(\mathrm{~s})
4
BCl
3
(
g
)
+
3
LiAlH
4
(
s
)
→
2
B
2
H
6
(
g
)
+
3
LiCl
(
s
)
+
3
AlCl
3
(
s
)
(d)
2
K
(
s
)
+
F
2
(
g
)
→
2
K
+
F
−
(
s
)
2 \mathrm{~K}(\mathrm{~s})+\mathrm{F}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{~K}^{+} \mathrm{F}^{-}(\mathrm{s})
2
K
(
s
)
+
F
2
(
g
)
→
2
K
+
F
−
(
s
)
(e)
4
N
H
3
(
g
)
+
5
O
2
(
g
)
→
4
N
O
(
g
)
+
6
H
2
O
(
g
)
4 \mathrm{NH}_{3}(\mathrm{~g})+5 \mathrm{O}_{2}(\mathrm{~g}) \rightarrow 4 \mathrm{NO}(\mathrm{~g})+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{~g})
4
NH
3
(
g
)
+
5
O
2
(
g
)
→
4
NO
(
g
)
+
6
H
2
O
(
g
)
Q4
EXERCISES
Fluorine reacts with ice and results in the change:
H
2
O
(
s
)
+
F
2
(
g
)
→
H
F
(
g
)
+
H
O
F
(
g
)
\mathrm{H}_{2} \mathrm{O}(\mathrm{s})+\mathrm{F}_{2}(\mathrm{~g}) \rightarrow \mathrm{HF}(\mathrm{~g})+\mathrm{HOF}(\mathrm{~g})
H
2
O
(
s
)
+
F
2
(
g
)
→
HF
(
g
)
+
HOF
(
g
)
Justify that this reaction is a redox reaction.
Q5
EXERCISES
Calculate the oxidation number of sulphur, chromium and nitrogen in
H
2
S
O
5
,
C
r
2
O
7
2
−
\mathrm{H}_{2} \mathrm{SO}_{5}, \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}
H
2
SO
5
,
Cr
2
O
7
2
−
and
N
O
3
−
\mathrm{NO}_{3}^{-}
NO
3
−
. Suggest structure of these compounds. Count for the fallacy.
Q6
EXERCISES
Write formulas for the following compounds:
(a)
Mercury(II) chloride
(b)
Nickel(II) sulphate
(c)
Tin(IV) oxide
(d)
Thallium(I) sulphate
(e) Iron(III) sulphate
(f) Chromium(III) oxide
Q7
EXERCISES
Suggest a list of the substances where carbon can exhibit oxidation states from -4 to +4 and nitrogen from -3 to +5.
Q8
EXERCISES
While sulphur dioxide and hydrogen peroxide can act as oxidising as well as reducing agents in their reactions, ozone and nitric acid act only as oxidants. Why ?
Q9
EXERCISES
Consider the reactions:
(a)
6
C
O
2
(
g
)
+
6
H
2
O
(
l
)
→
C
6
H
12
O
6
(
a
q
)
+
6
O
2
(
g
)
6 \mathrm{CO}_{2}(\mathrm{~g})+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}(\mathrm{aq})+6 \mathrm{O}_{2}(\mathrm{~g})
6
CO
2
(
g
)
+
6
H
2
O
(
l
)
→
C
6
H
12
O
6
(
aq
)
+
6
O
2
(
g
)
(b)
O
3
(
g
)
+
H
2
O
2
(
l
)
→
H
2
O
(
l
)
+
2
O
2
(
g
)
\mathrm{O}_{3}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{l}) \rightarrow \mathrm{H}_{2} \mathrm{O}(\mathrm{l})+2 \mathrm{O}_{2}(\mathrm{~g})
O
3
(
g
)
+
H
2
O
2
(
l
)
→
H
2
O
(
l
)
+
2
O
2
(
g
)
Why it is more appropriate to write these reactions as :
(a)
6
C
O
2
(
g
)
+
12
H
2
O
(
l
)
→
C
6
H
12
O
6
(
a
q
)
+
6
H
2
O
(
l
)
+
6
O
2
(
g
)
6 \mathrm{CO}_{2}(\mathrm{~g})+12 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}(\mathrm{aq})+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})+6 \mathrm{O}_{2}(\mathrm{g})
6
CO
2
(
g
)
+
12
H
2
O
(
l
)
→
C
6
H
12
O
6
(
aq
)
+
6
H
2
O
(
l
)
+
6
O
2
(
g
)
(b)
O
3
(
g
)
+
H
2
O
2
(
l
)
→
H
2
O
(
l
)
+
O
2
(
g
)
+
O
2
(
g
)
\mathrm{O}_{3}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{l}) \rightarrow \mathrm{H}_{2} \mathrm{O}(\mathrm{l})+\mathrm{O}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g})
O
3
(
g
)
+
H
2
O
2
(
l
)
→
H
2
O
(
l
)
+
O
2
(
g
)
+
O
2
(
g
)
Also suggest a technique to investigate the path of the above (a) and (b) redox reactions.
Q10
EXERCISES
The compound
A
g
F
2
\mathrm{AgF}_{2}
AgF
2
is unstable compound. However, if formed, the compound acts as a very strong oxidising agent. Why ?
Q11
EXERCISES
Whenever a reaction between an oxidising agent and a reducing agent is carried out, a compound of lower oxidation state is formed if the reducing agent is in excess and a compound of higher oxidation state is formed if the oxidising agent is in excess. Justify this statement giving three illustrations.
Q12
EXERCISES
How do you count for the following observations ?
(a)
Though alkaline potassium permanganate and acidic potassium permanganate both are used as oxidants, yet in the manufacture of benzoic acid from toluene we use alcoholic potassium permanganate as an oxidant. Why ? Write a balanced redox equation for the reaction.
(b)
When concentrated sulphuric acid is added to an inorganic mixture containing chloride, we get colourless pungent smelling gas HCl , but if the mixture contains bromide then we get red vapour of bromine. Why ?
Q13
EXERCISES
Identify the substance oxidised, reduced, oxidising agent and reducing agent for each of the following reactions:
(a)
2
A
g
B
r
(
s
)
+
C
6
H
6
O
2
(
a
q
)
→
2
A
g
(
s
)
+
2
H
B
r
(
a
q
)
+
C
6
H
4
O
2
(
a
q
)
2 \mathrm{AgBr}(\mathrm{s})+\mathrm{C}_{6} \mathrm{H}_{6} \mathrm{O}_{2}(\mathrm{aq}) \rightarrow 2 \mathrm{Ag}(\mathrm{s})+2 \mathrm{HBr}(\mathrm{aq})+\mathrm{C}_{6} \mathrm{H}_{4} \mathrm{O}_{2}(\mathrm{aq})
2
AgBr
(
s
)
+
C
6
H
6
O
2
(
aq
)
→
2
Ag
(
s
)
+
2
HBr
(
aq
)
+
C
6
H
4
O
2
(
aq
)
(b)
H
C
H
O
(
l
)
+
2
[
A
g
(
N
H
3
)
2
]
+
(
a
q
)
+
3
O
H
−
(
a
q
)
→
2
A
g
(
s
)
+
H
C
O
O
−
(
a
q
)
+
4
N
H
3
(
a
q
)
+
2
H
2
O
(
l
)
\mathrm{HCHO}(\mathrm{l})+2[\mathrm{Ag}(\mathrm{NH}_{3})_{2}]^{+}(\mathrm{aq})+3\mathrm{OH}^{-}(\mathrm{aq}) \rightarrow 2 \mathrm{Ag}(\mathrm{s})+\mathrm{HCOO}^{-}(\mathrm{aq})+4 \mathrm{NH}_{3}(\mathrm{aq}) +2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})
HCHO
(
l
)
+
2
[
Ag
(
NH
3
)
2
]
+
(
aq
)
+
3
OH
−
(
aq
)
→
2
Ag
(
s
)
+
HCOO
−
(
aq
)
+
4
NH
3
(
aq
)
+
2
H
2
O
(
l
)
(c)
H
C
H
O
(
l
)
+
2
C
u
2
+
(
a
q
)
+
5
O
H
−
(
a
q
)
→
C
u
2
O
(
s
)
+
H
C
O
O
−
(
a
q
)
+
3
H
2
O
(
l
)
\mathrm{HCHO}(\mathrm{l})+2 \mathrm{Cu}^{2+}(\mathrm{aq})+5\mathrm{OH}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cu}_{2}O(\mathrm{s})+\mathrm{HCOO}^{-}(\mathrm{aq})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})
HCHO
(
l
)
+
2
Cu
2
+
(
aq
)
+
5
OH
−
(
aq
)
→
Cu
2
O
(
s
)
+
HCOO
−
(
aq
)
+
3
H
2
O
(
l
)
(d)
N
2
H
4
(
l
)
+
2
H
2
O
2
(
l
)
→
N
2
(
g
)
+
4
H
2
O
(
l
)
\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{l})+2 \mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{l}) \rightarrow \mathrm{N}_{2}(\mathrm{~g})+4 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})
N
2
H
4
(
l
)
+
2
H
2
O
2
(
l
)
→
N
2
(
g
)
+
4
H
2
O
(
l
)
(e)
P
b
(
s
)
+
P
b
O
2
(
s
)
+
2
H
2
S
O
4
(
a
q
)
→
2
P
b
S
O
4
(
s
)
+
2
H
2
O
(
l
)
\mathrm{Pb}(\mathrm{s})+\mathrm{PbO}_{2}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \rightarrow 2 \mathrm{PbSO}_{4}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})
Pb
(
s
)
+
PbO
2
(
s
)
+
2
H
2
SO
4
(
aq
)
→
2
PbSO
4
(
s
)
+
2
H
2
O
(
l
)
Q14
EXERCISES
Consider the reactions :
2
S
2
O
3
2
−
(
a
q
)
+
I
2
(
s
)
→
S
4
O
6
2
−
(
a
q
)
+
2
I
−
(
a
q
)
2 \mathrm{S}_{2} \mathrm{O}_{3}^{2-}(\mathrm{aq})+\mathrm{I}_{2}(\mathrm{~s}) \rightarrow \mathrm{S}_{4} \mathrm{O}_{6}^{2-}(\mathrm{aq})+2 \mathrm{I}^{-}(\mathrm{aq})
2
S
2
O
3
2
−
(
aq
)
+
I
2
(
s
)
→
S
4
O
6
2
−
(
aq
)
+
2
I
−
(
aq
)
S
2
O
3
2
−
(
a
q
)
+
2
B
r
2
(
l
)
+
5
H
2
O
(
l
)
→
2
S
O
4
2
−
(
a
q
)
+
4
B
r
−
(
a
q
)
+
10
H
+
(
a
q
)
\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(\mathrm{aq})+2 \mathrm{Br}_{2}(\mathrm{l})+5 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{SO}_{4}^{2-}(\mathrm{aq})+4 \mathrm{Br}^{-}(\mathrm{aq})+10 \mathrm{H}^{+}(\mathrm{aq})
S
2
O
3
2
−
(
aq
)
+
2
Br
2
(
l
)
+
5
H
2
O
(
l
)
→
2
SO
4
2
−
(
aq
)
+
4
Br
−
(
aq
)
+
10
H
+
(
aq
)
Why does the same reductant, thiosulphate react differently with iodine and bromine?
Q15
EXERCISES
Justify giving reactions that among halogens, fluorine is the best oxidant and among hydrohalic compounds, hydroiodic acid is the best reductant.
Q16
EXERCISES
Why does the following reaction occur ?
X
e
O
6
4
−
(
a
q
)
+
2
F
−
(
a
q
)
+
6
H
+
(
a
q
)
→
X
e
O
3
(
g
)
+
F
2
(
g
)
+
3
H
2
O
(
l
)
\mathrm{XeO}_{6}{ }^{4-}(\mathrm{aq})+2 \mathrm{F}^{-}(\mathrm{aq})+6 \mathrm{H}^{+}(\mathrm{aq}) \rightarrow \mathrm{XeO}_{3}(\mathrm{~g})+\mathrm{F}_{2}(\mathrm{~g})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})
XeO
6
4
−
(
aq
)
+
2
F
−
(
aq
)
+
6
H
+
(
aq
)
→
XeO
3
(
g
)
+
F
2
(
g
)
+
3
H
2
O
(
l
)
What conclusion about the compound
N
a
4
X
e
O
6
\mathrm{Na}_{4} \mathrm{XeO}_{6}
Na
4
XeO
6
(of which
X
e
O
6
4
−
\mathrm{XeO}_{6}^{4-}
XeO
6
4
−
is a part) can be drawn from the reaction.
Q17
EXERCISES
Consider the reactions:
(a)
H
3
P
O
2
(
a
q
)
+
4
A
g
N
O
3
(
a
q
)
+
2
H
2
O
(
l
)
→
H
3
P
O
4
(
a
q
)
+
4
A
g
(
s
)
+
4
H
N
O
3
(
a
q
)
\mathrm{H}_{3} \mathrm{PO}_{2}(\mathrm{aq})+4 \mathrm{AgNO}_{3}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{H}_{3} \mathrm{PO}_{4}(\mathrm{aq})+4 \mathrm{Ag}(\mathrm{s})+4 \mathrm{HNO}_{3}(\mathrm{aq})
H
3
PO
2
(
aq
)
+
4
AgNO
3
(
aq
)
+
2
H
2
O
(
l
)
→
H
3
PO
4
(
aq
)
+
4
Ag
(
s
)
+
4
HNO
3
(
aq
)
(b)
H
3
P
O
2
(
a
q
)
+
2
C
u
S
O
4
(
a
q
)
+
2
H
2
O
(
l
)
→
H
3
P
O
4
(
a
q
)
+
2
C
u
(
s
)
+
H
2
S
O
4
(
a
q
)
\mathrm{H}_{3} \mathrm{PO}_{2}(\mathrm{aq})+2 \mathrm{CuSO}_{4}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{H}_{3} \mathrm{PO}_{4}(\mathrm{aq})+2 \mathrm{Cu}(\mathrm{s})+\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq})
H
3
PO
2
(
aq
)
+
2
CuSO
4
(
aq
)
+
2
H
2
O
(
l
)
→
H
3
PO
4
(
aq
)
+
2
Cu
(
s
)
+
H
2
SO
4
(
aq
)
(c)
C
6
H
5
C
H
O
(
l
)
+
2
[
A
g
(
N
H
3
)
2
]
+
(
a
q
)
+
3
O
H
−
(
a
q
)
→
C
6
H
5
C
O
O
−
(
a
q
)
+
2
A
g
(
s
)
+
4
N
H
3
(
a
q
)
+
2
H
2
O
(
l
)
\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CHO}(\mathrm{l})+2[\mathrm{Ag}(\mathrm{NH}_{3})_{2}]^{+}(\mathrm{aq})+3 \mathrm{OH}^{-}(\mathrm{aq}) \rightarrow \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COO}^{-}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s})+ 4 \mathrm{NH}_{3}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})
C
6
H
5
CHO
(
l
)
+
2
[
Ag
(
NH
3
)
2
]
+
(
aq
)
+
3
OH
−
(
aq
)
→
C
6
H
5
COO
−
(
aq
)
+
2
Ag
(
s
)
+
4
NH
3
(
aq
)
+
2
H
2
O
(
l
)
(d)
C
6
H
5
C
H
O
(
l
)
+
2
C
u
2
+
(
a
q
)
+
5
O
H
−
(
a
q
)
→
\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CHO}(\mathrm{l})+2 \mathrm{Cu}^{2+}(\mathrm{aq})+5 \mathrm{OH}^{-}(\mathrm{aq}) \rightarrow
C
6
H
5
CHO
(
l
)
+
2
Cu
2
+
(
aq
)
+
5
OH
−
(
aq
)
→
No change observed.
What inference do you draw about the behaviour of
A
g
+
\mathrm{Ag}^{+}
Ag
+
and
C
u
2
+
\mathrm{Cu}^{2+}
Cu
2
+
from these reactions?
Q18
EXERCISES
Balance the following redox reactions by ion - electron method :
(a)
M
n
O
4
−
(
a
q
)
+
I
−
(
a
q
)
→
M
n
O
2
(
s
)
+
I
2
(
s
)
\mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq}) \rightarrow \mathrm{MnO}_{2}(\mathrm{~s})+\mathrm{I}_{2}(\mathrm{~s})
MnO
4
−
(
aq
)
+
I
−
(
aq
)
→
MnO
2
(
s
)
+
I
2
(
s
)
(in basic medium)
(b)
M
n
O
4
−
(
a
q
)
+
S
O
2
(
g
)
→
M
n
2
+
(
a
q
)
+
H
S
O
4
−
(
a
q
)
\mathrm{MnO}_{4}^{-}(\mathrm{aq}) +\mathrm{SO}_{2}(\mathrm{~g}) \rightarrow \mathrm{Mn}^{2+}(\mathrm{aq}) +\mathrm{HSO}_{4}^{-}(\mathrm{aq})
MnO
4
−
(
aq
)
+
SO
2
(
g
)
→
Mn
2
+
(
aq
)
+
HSO
4
−
(
aq
)
(in acidic solution)
(c)
H
2
O
2
(
a
q
)
+
F
e
2
+
(
a
q
)
→
F
e
3
+
(
a
q
)
+
H
2
O
(
l
)
\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})+\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})
H
2
O
2
(
aq
)
+
Fe
2
+
(
aq
)
→
Fe
3
+
(
aq
)
+
H
2
O
(
l
)
(in acidic solution)
(d)
C
r
2
O
7
2
−
+
S
O
2
(
g
)
→
C
r
3
+
(
a
q
)
+
S
O
4
2
−
(
a
q
)
\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+\mathrm{SO}_{2}(\mathrm{~g}) \rightarrow \mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{SO}_{4}{ }^{2-}(\mathrm{aq})
Cr
2
O
7
2
−
+
SO
2
(
g
)
→
Cr
3
+
(
aq
)
+
SO
4
2
−
(
aq
)
(in acidic solution)
Q19
EXERCISES
Balance the following equations in basic medium by ion-electron method and oxidation number methods and identify the oxidising agent and the reducing agent.
(a)
P
4
(
s
)
+
O
H
−
(
a
q
)
→
P
H
3
(
g
)
+
H
P
O
2
−
(
a
q
)
\mathrm{P}_{4}(\mathrm{~s})+\mathrm{OH}^{-}(\mathrm{aq}) \rightarrow \mathrm{PH}_{3}(\mathrm{~g})+\mathrm{HPO}_{2}^{-}(\mathrm{aq})
P
4
(
s
)
+
OH
−
(
aq
)
→
PH
3
(
g
)
+
HPO
2
−
(
aq
)
(b)
N
2
H
4
(
l
)
+
C
l
O
3
−
(
a
q
)
→
N
O
(
g
)
+
C
l
−
(
g
)
\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{l})+\mathrm{ClO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{NO}(\mathrm{~g})+\mathrm{Cl}^{-}(\mathrm{g})
N
2
H
4
(
l
)
+
ClO
3
−
(
aq
)
→
NO
(
g
)
+
Cl
−
(
g
)
(c)
C
l
2
O
7
(
g
)
+
H
2
O
2
(
a
q
)
→
C
l
O
2
−
(
a
q
)
+
O
2
(
g
)
+
H
+
\mathrm{Cl}_{2} \mathrm{O}_{7}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}) \rightarrow \mathrm{ClO}_{2}^{-}(\mathrm{aq})+\mathrm{O}_{2}(\mathrm{~g})+\mathrm{H}^{+}
Cl
2
O
7
(
g
)
+
H
2
O
2
(
aq
)
→
ClO
2
−
(
aq
)
+
O
2
(
g
)
+
H
+
Q20
EXERCISES
What sorts of informations can you draw from the following reaction ?
(
C
N
)
2
(
g
)
+
2
O
H
−
(
a
q
)
→
C
N
−
(
a
q
)
+
C
N
O
−
(
a
q
)
+
H
2
O
(
l
)
(\mathrm{CN})_{2}(\mathrm{~g})+2 \mathrm{OH}^{-}(\mathrm{aq}) \rightarrow \mathrm{CN}^{-}(\mathrm{aq})+\mathrm{CNO}^{-}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})
(
CN
)
2
(
g
)
+
2
OH
−
(
aq
)
→
CN
−
(
aq
)
+
CNO
−
(
aq
)
+
H
2
O
(
l
)
Q21
EXERCISES
The
M
n
3
+
\mathrm{Mn}^{3+}
Mn
3
+
ion is unstable in solution and undergoes disproportionation to give
M
n
2
+
\mathrm{Mn}^{2+}
Mn
2
+
,
M
n
O
2
\mathrm{MnO}_{2}
MnO
2
, and
H
+
\mathrm{H}^{+}
H
+
ion. Write a balanced ionic equation for the reaction.
Q22
EXERCISES
Consider the elements:
Cs, Ne, I and F
(a)
Identify the element that exhibits only negative oxidation state.
(b)
Identify the element that exhibits only postive oxidation state.
(c)
Identify the element that exhibits both positive and negative oxidation states.
(d)
Identify the element which exhibits neither the negative nor does the positive oxidation state.
Q23
EXERCISES
Chlorine is used to purify drinking water. Excess of chlorine is harmful. The excess of chlorine is removed by treating with sulphur dioxide. Present a balanced equation for this redox change taking place in water.
Q24
EXERCISES
Refer to the periodic table given in your book and now answer the following questions:
(a)
Select the possible non metals that can show disproportionation reaction.
(b)
Select three metals that can show disproportionation reaction.
Q25
EXERCISES
In Ostwald's process for the manufacture of nitric acid, the first step involves the oxidation of ammonia gas by oxygen gas to give nitric oxide gas and steam. What is the maximum weight of nitric oxide that can be obtained starting only with 10.00 g. of ammonia and 20.00 g of oxygen ?
Q26
EXERCISES
Using the standard electrode potentials given in the Table 7.1, predict if the reaction between the following is feasible:
(a)
F
e
3
+
(
a
q
)
\mathrm{Fe}^{3+}(\mathrm{aq})
Fe
3
+
(
aq
)
and
I
−
(
a
q
)
\mathrm{I}^{-}(\mathrm{aq})
I
−
(
aq
)
(b)
A
g
+
(
a
q
)
\mathrm{Ag}^{+}(\mathrm{aq})
Ag
+
(
aq
)
and
C
u
(
s
)
\mathrm{Cu}(\mathrm{s})
Cu
(
s
)
(c)
F
e
3
+
(
a
q
)
\mathrm{Fe}^{3+}(\mathrm{aq})
Fe
3
+
(
aq
)
and
C
u
(
s
)
\mathrm{Cu}(\mathrm{s})
Cu
(
s
)
(d)
A
g
(
s
)
\mathrm{Ag}(\mathrm{s})
Ag
(
s
)
and
F
e
3
+
(
a
q
)
\mathrm{Fe}^{3+}(\mathrm{aq})
Fe
3
+
(
aq
)
(e)
B
r
2
(
a
q
)
\mathrm{Br}_{2}(\mathrm{aq})
Br
2
(
aq
)
and
F
e
2
+
(
a
q
)
\mathrm{Fe}^{2+}(\mathrm{aq})
Fe
2
+
(
aq
)
.
Q27
EXERCISES
Predict the products of electrolysis in each of the following:
(i)
An aqueous solution of
A
g
N
O
3
\mathrm{AgNO}_{3}
AgNO
3
with silver electrodes
(ii)
An aqueous solution
A
g
N
O
3
\mathrm{AgNO}_{3}
AgNO
3
with platinum electrodes
(iii)
A dilute solution of
H
2
S
O
4
\mathrm{H}_{2} \mathrm{SO}_{4}
H
2
SO
4
with platinum electrodes
(iv)
An aqueous solution of
C
u
C
l
2
\mathrm{CuCl}_{2}
CuCl
2
with platinum electrodes.
Q28
EXERCISES
Arrange the following metals in the order in which they displace each other from the solution of their salts. Al, Cu, Fe, Mg and Zn .
Q29
EXERCISES
Given the standard electrode potentials,
K
+
/
K
=
−
2.93
V
,
A
g
+
/
A
g
=
0.80
V
\mathrm{K}^{+} / \mathrm{K}=-2.93 \mathrm{~V}, \mathrm{Ag}^{+} / \mathrm{Ag}=0.80 \mathrm{~V}
K
+
/
K
=
−
2.93
V
,
Ag
+
/
Ag
=
0.80
V
,
H
g
2
+
/
H
g
=
0.79
V
\mathrm{Hg}^{2+} / \mathrm{Hg}=0.79 \mathrm{~V}
Hg
2
+
/
Hg
=
0.79
V
M
g
2
+
/
M
g
=
−
2.37
V
.
C
r
3
+
/
C
r
=
−
0.74
V
\mathrm{Mg}^{2+} / \mathrm{Mg}=-2.37 \mathrm{~V} . \mathrm{Cr}^{3+} / \mathrm{Cr}=-0.74 \mathrm{~V}
Mg
2
+
/
Mg
=
−
2.37
V
.
Cr
3
+
/
Cr
=
−
0.74
V
arrange these metals in their increasing order of reducing power.
Q30
EXERCISES
Depict the galvanic cell in which the reaction
Z
n
(
s
)
+
2
A
g
+
(
a
q
)
→
Z
n
2
+
(
a
q
)
+
2
A
g
(
s
)
\mathrm{Zn}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s})
Zn
(
s
)
+
2
Ag
+
(
aq
)
→
Zn
2
+
(
aq
)
+
2
Ag
(
s
)
takes place, Further show:
(i)
which of the electrode is negatively charged,
(ii)
the carriers of the current in the cell, and
(iii)
individual reaction at each electrode.
Q1
In-text Problems
In the reactions given below, identify the species undergoing oxidation and reduction:
(i)
H
2
S
(
g
)
+
C
l
2
(
g
)
→
2
H
C
l
(
g
)
+
S
(
s
)
\mathrm{H}_{2} \mathrm{~S}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{HCl}(\mathrm{~g})+\mathrm{S}(\mathrm{s})
H
2
S
(
g
)
+
Cl
2
(
g
)
→
2
HCl
(
g
)
+
S
(
s
)
(ii)
3
F
e
3
O
4
(
s
)
+
8
A
l
(
s
)
→
9
F
e
(
s
)
+
4
A
l
2
O
3
(
s
)
3 \mathrm{Fe}_{3} \mathrm{O}_{4}(\mathrm{~s})+8 \mathrm{Al}(\mathrm{s}) \rightarrow 9 \mathrm{Fe}(\mathrm{s}) +4 \mathrm{Al}_{2} \mathrm{O}_{3}(\mathrm{~s})
3
Fe
3
O
4
(
s
)
+
8
Al
(
s
)
→
9
Fe
(
s
)
+
4
Al
2
O
3
(
s
)
(iii)
2
N
a
(
s
)
+
H
2
(
g
)
→
2
N
a
H
(
s
)
2 \mathrm{Na}(\mathrm{s})+\mathrm{H}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{NaH}(\mathrm{s})
2
Na
(
s
)
+
H
2
(
g
)
→
2
NaH
(
s
)
Q2
In-text Problems
Justify that the reaction:
2
N
a
(
s
)
+
H
2
(
g
)
→
2
N
a
H
(
s
)
2 \mathrm{Na}(\mathrm{s})+\mathrm{H}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{NaH}(\mathrm{s})
2
Na
(
s
)
+
H
2
(
g
)
→
2
NaH
(
s
)
is a redox change.
Q3
In-text Problems
Using Stock notation, represent the following compounds:
H
A
u
C
l
4
,
T
l
2
O
,
F
e
O
,
F
e
2
O
3
,
C
u
I
,
C
u
O
,
M
n
O
\mathrm{HAuCl}_{4}, \mathrm{Tl}_{2} \mathrm{O}, \mathrm{FeO}, \mathrm{Fe}_{2} \mathrm{O}_{3}, \mathrm{CuI}, \mathrm{CuO}, \mathrm{MnO}
HAuCl
4
,
Tl
2
O
,
FeO
,
Fe
2
O
3
,
CuI
,
CuO
,
MnO
and
M
n
O
2
\mathrm{MnO}_{2}
MnO
2
.
Q4
In-text Problems
Justify that the reaction:
2
C
u
2
O
(
s
)
+
C
u
2
S
(
s
)
→
6
C
u
(
s
)
+
S
O
2
(
g
)
2 \mathrm{Cu}_{2} \mathrm{O}(\mathrm{s})+\mathrm{Cu}_{2} \mathrm{S}(\mathrm{s}) \rightarrow 6 \mathrm{Cu}(\mathrm{s})+\mathrm{SO}_{2}(\mathrm{~g})
2
Cu
2
O
(
s
)
+
Cu
2
S
(
s
)
→
6
Cu
(
s
)
+
SO
2
(
g
)
is a redox reaction. Identify the species oxidised/reduced, which acts as an oxidant and which acts as a reductant.
Q5
In-text Problems
Which of the following species, do not show disproportionation reaction and why?
C
l
O
−
,
C
l
O
2
−
,
C
l
O
3
−
\mathrm{ClO}^{-}, \mathrm{ClO}_{2}^{-}, \mathrm{ClO}_{3}^{-}
ClO
−
,
ClO
2
−
,
ClO
3
−
and
C
l
O
4
−
\mathrm{ClO}_{4}^{-}
ClO
4
−
Also write reaction for each of the species that disproportionates.
Q6
In-text Problems
Suggest a scheme of classification of the following redox reactions
(a)
N
2
(
g
)
+
O
2
(
g
)
→
2
N
O
(
g
)
\mathrm{N}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{NO}(\mathrm{~g})
N
2
(
g
)
+
O
2
(
g
)
→
2
NO
(
g
)
(b)
2
P
b
(
N
O
3
)
2
(
s
)
→
2
P
b
O
(
s
)
+
4
N
O
2
(
g
)
+
O
2
(
g
)
2 \mathrm{~Pb}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{~s}) \rightarrow 2 \mathrm{PbO}(\mathrm{s})+4 \mathrm{NO}_{2}(\mathrm{~g})+ \mathrm{O}_{2}(\mathrm{~g})
2
Pb
(
NO
3
)
2
(
s
)
→
2
PbO
(
s
)
+
4
NO
2
(
g
)
+
O
2
(
g
)
(c)
N
a
H
(
s
)
+
H
2
O
(
l
)
→
N
a
O
H
(
a
q
)
+
H
2
(
g
)
\mathrm{NaH}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{NaOH}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})
NaH
(
s
)
+
H
2
O
(
l
)
→
NaOH
(
aq
)
+
H
2
(
g
)
(d)
2
N
O
2
(
g
)
+
2
O
H
−
(
a
q
)
→
N
O
2
−
(
a
q
)
+
N
O
3
−
(
a
q
)
+
H
2
O
(
l
)
2 \mathrm{NO}_{2}(\mathrm{~g})+2 \mathrm{OH}^{-}(\mathrm{aq}) \rightarrow \mathrm{NO}_{2}^{-}(\mathrm{aq})+ \mathrm{NO}_{3}^{-}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})
2
NO
2
(
g
)
+
2
OH
−
(
aq
)
→
NO
2
−
(
aq
)
+
NO
3
−
(
aq
)
+
H
2
O
(
l
)
Q7
In-text Problems
Why do the following reactions proceed differently ?
P
b
3
O
4
+
8
H
C
l
→
3
P
b
C
l
2
+
C
l
2
+
4
H
2
O
\mathrm{Pb}_{3} \mathrm{O}_{4}+8 \mathrm{HCl} \rightarrow 3 \mathrm{PbCl}_{2}+\mathrm{Cl}_{2}+4 \mathrm{H}_{2} \mathrm{O}
Pb
3
O
4
+
8
HCl
→
3
PbCl
2
+
Cl
2
+
4
H
2
O
and
P
b
3
O
4
+
4
H
N
O
3
→
2
P
b
(
N
O
3
)
2
+
P
b
O
2
+
2
H
2
O
\mathrm{Pb}_{3} \mathrm{O}_{4}+4 \mathrm{HNO}_{3} \rightarrow 2 \mathrm{~Pb}\left(\mathrm{NO}_{3}\right)_{2}+\mathrm{PbO}_{2}+ 2 \mathrm{H}_{2} \mathrm{O}
Pb
3
O
4
+
4
HNO
3
→
2
Pb
(
NO
3
)
2
+
PbO
2
+
2
H
2
O
Q8
In-text Problems
Write the net ionic equation for the reaction of potassium dichromate(VI),
K
2
C
r
2
O
7
\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}
K
2
Cr
2
O
7
with sodium sulphite,
N
a
2
S
O
3
\mathrm{Na}_{2} \mathrm{SO}_{3}
Na
2
SO
3
, in an acid solution to give chromium(III) ion and the sulphate ion.
Q9
In-text Problems
Permanganate ion reacts with bromide ion in basic medium to give manganese dioxide and bromate ion. Write the balanced ionic equation for the reaction.
Q10
In-text Problems
Permanganate(VII) ion,
M
n
O
4
−
\mathrm{MnO}_{4}^{-}
MnO
4
−
in basic solution oxidises iodide ion,
I
−
\mathrm{I}^{-}
I
−
to produce molecular iodine (
I
2
\mathrm{I}_{2}
I
2
) and manganese (IV) oxide (
M
n
O
2
\mathrm{MnO}_{2}
MnO
2
). Write a balanced ionic equation to represent this redox reaction.
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