NCERT Textbook: Redox Reactions | Chemistry Class 11 - NEET PDF Download

 Page 1


235 redox reactions
Chemistry deals with varieties of matter and change 
of one kind of matter into the other. Transformation of 
matter from one kind into another occurs through the 
various types of reactions. One important category of such 
reactions is Redox Reactions. A number of phenomena, 
both physical as well as biological, are concerned with 
redox reactions. These reactions find extensive use in 
pharmaceutical, biological, industrial, metallurgical and 
agricultural areas. The importance of these reactions is 
apparent from the fact that burning of different types of 
fuels for obtaining energy for domestic, transport and 
other commercial purposes, electrochemical processes 
for extraction of highly reactive metals and non-metals, 
manufacturing of  chemical  compounds like  caustic 
soda, operation of dry and wet batteries and corrosion of 
metals fall within the purview of redox processes. Of late, 
environmental issues like Hydrogen Economy (use of 
liquid hydrogen as fuel) and development of ‘Ozone Hole’ 
have started figuring under redox phenomenon. 
7.1	 CLASSICAL 	 IDEA 	 OF 	 REDOX 	 REACTIONS 	 –	 	
OXIDATION AND REDUCTION REACTIONS
Originally, the term oxidation was used to describe the 
addition of oxygen to an element or a compound. Because 
of the presence of dioxygen in the atmosphere (~20%), 
many elements combine with it and this is the principal 
reason why they commonly occur on the earth in the  
form of their oxides. The following reactions represent 
oxidation processes according to the limited definition of 
oxidation: 
2 Mg (s)  +  O
2
 (g)  ?  2 MgO (s) (7.1)
S (s) + O
2
 (g)  ?  SO
2
 (g) (7.2)
After studying this unit you will be 
able to 
• identify redox reactions as a class 
of reactions in which oxidation 
and reduction reactions occur 
simultaneously; 
• define the terms oxidation, 
reduction, oxidant (oxidising 
agent) and reductant (reducing 
agent); 
• explain mechanism of redox 
reactions by electron transfer 
process;
• use  the concept of oxidation 
number to identify oxidant and 
reductant in a reaction;
• classify redox reaction into 
combination (synthesis), 
decomposition, displacement 
and disproportionation 
reactions; 
• suggest a comparative order 
among various reductants and 
oxidants; 
• balance chemical equations 
using (i) oxidation number  
(ii) half reaction method;
• learn the concept of redox 
reactions in terms of electrode 
processes.
UNIT 7
REDOX REACTIONS
Where there is oxidation, there is always reduction –  
Chemistry is essentially a study of redox systems.
Unit 7.indd   235 10/10/2022   10:37:02 AM
2024-25
Page 2


235 redox reactions
Chemistry deals with varieties of matter and change 
of one kind of matter into the other. Transformation of 
matter from one kind into another occurs through the 
various types of reactions. One important category of such 
reactions is Redox Reactions. A number of phenomena, 
both physical as well as biological, are concerned with 
redox reactions. These reactions find extensive use in 
pharmaceutical, biological, industrial, metallurgical and 
agricultural areas. The importance of these reactions is 
apparent from the fact that burning of different types of 
fuels for obtaining energy for domestic, transport and 
other commercial purposes, electrochemical processes 
for extraction of highly reactive metals and non-metals, 
manufacturing of  chemical  compounds like  caustic 
soda, operation of dry and wet batteries and corrosion of 
metals fall within the purview of redox processes. Of late, 
environmental issues like Hydrogen Economy (use of 
liquid hydrogen as fuel) and development of ‘Ozone Hole’ 
have started figuring under redox phenomenon. 
7.1	 CLASSICAL 	 IDEA 	 OF 	 REDOX 	 REACTIONS 	 –	 	
OXIDATION AND REDUCTION REACTIONS
Originally, the term oxidation was used to describe the 
addition of oxygen to an element or a compound. Because 
of the presence of dioxygen in the atmosphere (~20%), 
many elements combine with it and this is the principal 
reason why they commonly occur on the earth in the  
form of their oxides. The following reactions represent 
oxidation processes according to the limited definition of 
oxidation: 
2 Mg (s)  +  O
2
 (g)  ?  2 MgO (s) (7.1)
S (s) + O
2
 (g)  ?  SO
2
 (g) (7.2)
After studying this unit you will be 
able to 
• identify redox reactions as a class 
of reactions in which oxidation 
and reduction reactions occur 
simultaneously; 
• define the terms oxidation, 
reduction, oxidant (oxidising 
agent) and reductant (reducing 
agent); 
• explain mechanism of redox 
reactions by electron transfer 
process;
• use  the concept of oxidation 
number to identify oxidant and 
reductant in a reaction;
• classify redox reaction into 
combination (synthesis), 
decomposition, displacement 
and disproportionation 
reactions; 
• suggest a comparative order 
among various reductants and 
oxidants; 
• balance chemical equations 
using (i) oxidation number  
(ii) half reaction method;
• learn the concept of redox 
reactions in terms of electrode 
processes.
UNIT 7
REDOX REACTIONS
Where there is oxidation, there is always reduction –  
Chemistry is essentially a study of redox systems.
Unit 7.indd   235 10/10/2022   10:37:02 AM
2024-25
236 chemistry In reactions (7.1) and (7.2), the elements 
magnesium and sulphur are oxidised on 
account of addition of oxygen to them. 
Similarly,  methane is oxidised owing to the 
addition of oxygen to it.
CH
4
 (g) + 2O
2
 (g) ? CO
2
 (g) + 2H
2
O (l) (7.3)
A careful examination of reaction (7.3) in 
which hydrogen has been replaced by oxygen 
prompted chemists to reinterpret oxidation 
in terms of removal of hydrogen from it and, 
therefore, the scope of term oxidation was 
broadened to include the removal of hydrogen 
from a substance. The following illustration is 
another reaction where removal of hydrogen 
can also be cited as an oxidation reaction. 
2 H
2
S(g) + O
2
 (g) ? 2 S (s) + 2 H
2
O (l) (7.4)
As knowledge of chemists grew, it was 
natural to extend the term oxidation for 
reactions similar to (7.1 to 7.4), which do 
not involve oxygen but other electronegative 
elements. The oxidation of magnesium with 
fluorine, chlorine and sulphur etc. occurs 
according to the following reactions : 
Mg (s) + F
2
 (g) ? MgF
2
 (s) (7.5)
Mg (s) + Cl
2
 (g) ? MgCl
2
 (s) (7.6)
Mg (s) + S (s) ? MgS (s) (7.7)
Incorporating the reactions (7.5 to 
7.7) within the fold of oxidation reactions  
encouraged chemists to consider not only 
the removal of hydrogen as oxidation, but 
also the removal of electropositive elements 
as oxidation. Thus the reaction : 
2K
4
 [Fe(CN)
6
](aq) + H
2
O
2
 (aq) ?2K
3
[Fe(CN)
6
](aq) 
+ 2 KOH (aq)
is interpreted as oxidation due to the removal 
of electropositive element potassium from 
potassium ferrocyanide before it changes to 
potassium ferricyanide. To summarise, the 
term “oxidation” is defined as the addition 
of oxygen/electronegative element to 
a substance or removal of hydrogen/
electropositive element from a substance.
In the beginning, reduction was considered 
as removal of oxygen from a compound. 
However, the term reduction has been 
broadened these days to include removal 
of oxygen/electronegative element from 
a substance or addition of hydrogen/
electropositive element to a substance. 
According to the definition given above, 
the following are the examples of reduction 
processes:
2 HgO (s)   2 Hg (l) + O
2 
(g) (7.8)
(removal of oxygen from mercuric oxide )
2 FeCl
3
 (aq) + H
2
 (g) ?2 FeCl
2
 (aq) + 2 HCl(aq)
(7.9)
(removal of electronegative element, chlorine 
from ferric chloride)
CH
2 
= CH
2
 (g) + H
2
 (g) ? H
3
C – CH
3
 (g) (7.10)
(addition of hydrogen)
2HgCl
2
 (aq) + SnCl
2
 (aq) ? Hg
2
Cl
2 
(s)+SnCl
4
 (aq)
(7.11)
(addition of mercury to mercuric chloride)
In reaction (7.11) simultaneous oxidation 
of stannous chloride to stannic chloride is 
also occurring because of the addition of 
electronegative element chlorine to it. It was 
soon realised that oxidation and reduction 
always occur simultaneously (as will be 
apparent by re-examining all the equations 
given above), hence, the word “redox” was 
coined for this class of chemical reactions. 
Problem 7.1
In the reactions given below, identify 
the species undergoing oxidation and 
reduction: 
(i)  H
2
S (g) + Cl
2
 (g) ? 2 HCl (g) + S (s)
(ii) 3Fe
3
O
4
 (s) + 8 Al (s) ? 9 Fe (s) 
                                                 + 4Al
2
O
3
 (s)
(iii) 2 Na (s) + H
2
 (g) ? 2 NaH (s)
Solution  
(i) H
2
S is oxidised because a more 
electronegative element, chlorine is added 
to hydrogen (or a more electropositive 
element, hydrogen has been removed  
from S). Chlorine is reduced due to 
addition of hydrogen to it.
(ii)  Aluminium is oxidised because 
oxygen is added to it. Ferrous ferric oxide 
Unit 7.indd   236 10/10/2022   10:37:03 AM
2024-25
Page 3


235 redox reactions
Chemistry deals with varieties of matter and change 
of one kind of matter into the other. Transformation of 
matter from one kind into another occurs through the 
various types of reactions. One important category of such 
reactions is Redox Reactions. A number of phenomena, 
both physical as well as biological, are concerned with 
redox reactions. These reactions find extensive use in 
pharmaceutical, biological, industrial, metallurgical and 
agricultural areas. The importance of these reactions is 
apparent from the fact that burning of different types of 
fuels for obtaining energy for domestic, transport and 
other commercial purposes, electrochemical processes 
for extraction of highly reactive metals and non-metals, 
manufacturing of  chemical  compounds like  caustic 
soda, operation of dry and wet batteries and corrosion of 
metals fall within the purview of redox processes. Of late, 
environmental issues like Hydrogen Economy (use of 
liquid hydrogen as fuel) and development of ‘Ozone Hole’ 
have started figuring under redox phenomenon. 
7.1	 CLASSICAL 	 IDEA 	 OF 	 REDOX 	 REACTIONS 	 –	 	
OXIDATION AND REDUCTION REACTIONS
Originally, the term oxidation was used to describe the 
addition of oxygen to an element or a compound. Because 
of the presence of dioxygen in the atmosphere (~20%), 
many elements combine with it and this is the principal 
reason why they commonly occur on the earth in the  
form of their oxides. The following reactions represent 
oxidation processes according to the limited definition of 
oxidation: 
2 Mg (s)  +  O
2
 (g)  ?  2 MgO (s) (7.1)
S (s) + O
2
 (g)  ?  SO
2
 (g) (7.2)
After studying this unit you will be 
able to 
• identify redox reactions as a class 
of reactions in which oxidation 
and reduction reactions occur 
simultaneously; 
• define the terms oxidation, 
reduction, oxidant (oxidising 
agent) and reductant (reducing 
agent); 
• explain mechanism of redox 
reactions by electron transfer 
process;
• use  the concept of oxidation 
number to identify oxidant and 
reductant in a reaction;
• classify redox reaction into 
combination (synthesis), 
decomposition, displacement 
and disproportionation 
reactions; 
• suggest a comparative order 
among various reductants and 
oxidants; 
• balance chemical equations 
using (i) oxidation number  
(ii) half reaction method;
• learn the concept of redox 
reactions in terms of electrode 
processes.
UNIT 7
REDOX REACTIONS
Where there is oxidation, there is always reduction –  
Chemistry is essentially a study of redox systems.
Unit 7.indd   235 10/10/2022   10:37:02 AM
2024-25
236 chemistry In reactions (7.1) and (7.2), the elements 
magnesium and sulphur are oxidised on 
account of addition of oxygen to them. 
Similarly,  methane is oxidised owing to the 
addition of oxygen to it.
CH
4
 (g) + 2O
2
 (g) ? CO
2
 (g) + 2H
2
O (l) (7.3)
A careful examination of reaction (7.3) in 
which hydrogen has been replaced by oxygen 
prompted chemists to reinterpret oxidation 
in terms of removal of hydrogen from it and, 
therefore, the scope of term oxidation was 
broadened to include the removal of hydrogen 
from a substance. The following illustration is 
another reaction where removal of hydrogen 
can also be cited as an oxidation reaction. 
2 H
2
S(g) + O
2
 (g) ? 2 S (s) + 2 H
2
O (l) (7.4)
As knowledge of chemists grew, it was 
natural to extend the term oxidation for 
reactions similar to (7.1 to 7.4), which do 
not involve oxygen but other electronegative 
elements. The oxidation of magnesium with 
fluorine, chlorine and sulphur etc. occurs 
according to the following reactions : 
Mg (s) + F
2
 (g) ? MgF
2
 (s) (7.5)
Mg (s) + Cl
2
 (g) ? MgCl
2
 (s) (7.6)
Mg (s) + S (s) ? MgS (s) (7.7)
Incorporating the reactions (7.5 to 
7.7) within the fold of oxidation reactions  
encouraged chemists to consider not only 
the removal of hydrogen as oxidation, but 
also the removal of electropositive elements 
as oxidation. Thus the reaction : 
2K
4
 [Fe(CN)
6
](aq) + H
2
O
2
 (aq) ?2K
3
[Fe(CN)
6
](aq) 
+ 2 KOH (aq)
is interpreted as oxidation due to the removal 
of electropositive element potassium from 
potassium ferrocyanide before it changes to 
potassium ferricyanide. To summarise, the 
term “oxidation” is defined as the addition 
of oxygen/electronegative element to 
a substance or removal of hydrogen/
electropositive element from a substance.
In the beginning, reduction was considered 
as removal of oxygen from a compound. 
However, the term reduction has been 
broadened these days to include removal 
of oxygen/electronegative element from 
a substance or addition of hydrogen/
electropositive element to a substance. 
According to the definition given above, 
the following are the examples of reduction 
processes:
2 HgO (s)   2 Hg (l) + O
2 
(g) (7.8)
(removal of oxygen from mercuric oxide )
2 FeCl
3
 (aq) + H
2
 (g) ?2 FeCl
2
 (aq) + 2 HCl(aq)
(7.9)
(removal of electronegative element, chlorine 
from ferric chloride)
CH
2 
= CH
2
 (g) + H
2
 (g) ? H
3
C – CH
3
 (g) (7.10)
(addition of hydrogen)
2HgCl
2
 (aq) + SnCl
2
 (aq) ? Hg
2
Cl
2 
(s)+SnCl
4
 (aq)
(7.11)
(addition of mercury to mercuric chloride)
In reaction (7.11) simultaneous oxidation 
of stannous chloride to stannic chloride is 
also occurring because of the addition of 
electronegative element chlorine to it. It was 
soon realised that oxidation and reduction 
always occur simultaneously (as will be 
apparent by re-examining all the equations 
given above), hence, the word “redox” was 
coined for this class of chemical reactions. 
Problem 7.1
In the reactions given below, identify 
the species undergoing oxidation and 
reduction: 
(i)  H
2
S (g) + Cl
2
 (g) ? 2 HCl (g) + S (s)
(ii) 3Fe
3
O
4
 (s) + 8 Al (s) ? 9 Fe (s) 
                                                 + 4Al
2
O
3
 (s)
(iii) 2 Na (s) + H
2
 (g) ? 2 NaH (s)
Solution  
(i) H
2
S is oxidised because a more 
electronegative element, chlorine is added 
to hydrogen (or a more electropositive 
element, hydrogen has been removed  
from S). Chlorine is reduced due to 
addition of hydrogen to it.
(ii)  Aluminium is oxidised because 
oxygen is added to it. Ferrous ferric oxide 
Unit 7.indd   236 10/10/2022   10:37:03 AM
2024-25
237 redox reactions
(Fe
3
O
4
) is reduced because oxygen has 
been removed from it. 
(iii)  With the careful application of the 
concept of electronegativity only we 
may infer that sodium is oxidised and 
hydrogen is reduced. 
Reaction (iii) chosen here prompts us to 
think in terms of another way to define 
redox reactions. 
7.2 REDOX REACTIONS IN TERMS OF 
ELECTRON  TRANSFER REACTIONS
We have already learnt that the reactions 
2Na(s) + Cl
2
(g) ?  2NaCl (s) (7.12)
4Na(s) + O
2
(g) ?  2Na
2
O(s) (7.13)
2Na(s) + S(s) ?  Na
2
S(s) (7.14)
are redox reactions because in each of these 
reactions sodium is oxidised due to the addition 
of either oxygen or more electronegative 
element to sodium. Simultaneously, chlorine, 
oxygen and sulphur are reduced because to 
each of these, the electropositive element 
sodium has been added. From our knowledge 
of chemical bonding we also know that sodium 
chloride, sodium oxide and sodium sulphide 
are ionic compounds and perhaps better 
written as Na
+
Cl
–
 (s), (Na
+
)
2
O
2–
(s), and (Na
+
)
2
 
S
2–
(s). Development of charges on the species 
produced suggests us to rewrite the reactions 
(7.12 to 7.14) in the following manner : 
For convenience, each of the above 
processes can be considered as two separate 
steps, one involving the loss of electrons 
and the other the gain of electrons. As an 
illustration, we may further elaborate one of 
these, say, the formation of  sodium chloride. 
2 Na(s) ? 2 Na
+
(g)
 
  +   2e
–
Cl
2
(g) + 2e
–
 ? 2 Cl
–
(g)
Each of the above steps is called a half 
reaction, which explicitly shows involvement 
of electrons. Sum of the half reactions gives 
the overall reaction : 
2 Na(s) + Cl
2
 (g)  ? 2 Na
+
 Cl
–
 (s) or 2 NaCl (s)
Reactions 7.12 to 7.14 suggest that half 
reactions that involve loss of electrons are 
called oxidation reactions. Similarly, the 
half reactions that involve gain of electrons 
are called reduction reactions.  It may not 
be out of context to mention here that the 
new way of defining oxidation and reduction 
has been achieved only by establishing a 
correlation between the behaviour of species 
as per the classical idea and their interplay 
in electron-transfer change. In reactions (7.12 
to 7.14) sodium, which is oxidised, acts as  
a reducing agent because it donates electron 
to each of the elements interacting with it and 
thus helps in reducing them. Chlorine, oxygen 
and sulphur are reduced and act as oxidising 
agents because these accept electrons from 
sodium. To summarise, we may mention that 
Oxidation : Loss of electron(s) by any species.
Reduction : Gain of electron(s) by any species.
Oxidising agent : Acceptor of electron(s).
Reducing agent : Donor of electron(s).
Problem 7.2  Justify that the reaction:
2 Na(s) + H
2
(g) ?  2 NaH (s) is a redox 
change.
Solution
Since in the above reaction the compound 
formed is an ionic compound, which may 
also be represented as Na
+
H
–
 (s), this 
suggests that one half reaction in this 
process is : 
2 Na (s) ? 2 Na
+
(g)  +   2e
–
Unit 7.indd   237 10/10/2022   10:37:03 AM
2024-25
Page 4


235 redox reactions
Chemistry deals with varieties of matter and change 
of one kind of matter into the other. Transformation of 
matter from one kind into another occurs through the 
various types of reactions. One important category of such 
reactions is Redox Reactions. A number of phenomena, 
both physical as well as biological, are concerned with 
redox reactions. These reactions find extensive use in 
pharmaceutical, biological, industrial, metallurgical and 
agricultural areas. The importance of these reactions is 
apparent from the fact that burning of different types of 
fuels for obtaining energy for domestic, transport and 
other commercial purposes, electrochemical processes 
for extraction of highly reactive metals and non-metals, 
manufacturing of  chemical  compounds like  caustic 
soda, operation of dry and wet batteries and corrosion of 
metals fall within the purview of redox processes. Of late, 
environmental issues like Hydrogen Economy (use of 
liquid hydrogen as fuel) and development of ‘Ozone Hole’ 
have started figuring under redox phenomenon. 
7.1	 CLASSICAL 	 IDEA 	 OF 	 REDOX 	 REACTIONS 	 –	 	
OXIDATION AND REDUCTION REACTIONS
Originally, the term oxidation was used to describe the 
addition of oxygen to an element or a compound. Because 
of the presence of dioxygen in the atmosphere (~20%), 
many elements combine with it and this is the principal 
reason why they commonly occur on the earth in the  
form of their oxides. The following reactions represent 
oxidation processes according to the limited definition of 
oxidation: 
2 Mg (s)  +  O
2
 (g)  ?  2 MgO (s) (7.1)
S (s) + O
2
 (g)  ?  SO
2
 (g) (7.2)
After studying this unit you will be 
able to 
• identify redox reactions as a class 
of reactions in which oxidation 
and reduction reactions occur 
simultaneously; 
• define the terms oxidation, 
reduction, oxidant (oxidising 
agent) and reductant (reducing 
agent); 
• explain mechanism of redox 
reactions by electron transfer 
process;
• use  the concept of oxidation 
number to identify oxidant and 
reductant in a reaction;
• classify redox reaction into 
combination (synthesis), 
decomposition, displacement 
and disproportionation 
reactions; 
• suggest a comparative order 
among various reductants and 
oxidants; 
• balance chemical equations 
using (i) oxidation number  
(ii) half reaction method;
• learn the concept of redox 
reactions in terms of electrode 
processes.
UNIT 7
REDOX REACTIONS
Where there is oxidation, there is always reduction –  
Chemistry is essentially a study of redox systems.
Unit 7.indd   235 10/10/2022   10:37:02 AM
2024-25
236 chemistry In reactions (7.1) and (7.2), the elements 
magnesium and sulphur are oxidised on 
account of addition of oxygen to them. 
Similarly,  methane is oxidised owing to the 
addition of oxygen to it.
CH
4
 (g) + 2O
2
 (g) ? CO
2
 (g) + 2H
2
O (l) (7.3)
A careful examination of reaction (7.3) in 
which hydrogen has been replaced by oxygen 
prompted chemists to reinterpret oxidation 
in terms of removal of hydrogen from it and, 
therefore, the scope of term oxidation was 
broadened to include the removal of hydrogen 
from a substance. The following illustration is 
another reaction where removal of hydrogen 
can also be cited as an oxidation reaction. 
2 H
2
S(g) + O
2
 (g) ? 2 S (s) + 2 H
2
O (l) (7.4)
As knowledge of chemists grew, it was 
natural to extend the term oxidation for 
reactions similar to (7.1 to 7.4), which do 
not involve oxygen but other electronegative 
elements. The oxidation of magnesium with 
fluorine, chlorine and sulphur etc. occurs 
according to the following reactions : 
Mg (s) + F
2
 (g) ? MgF
2
 (s) (7.5)
Mg (s) + Cl
2
 (g) ? MgCl
2
 (s) (7.6)
Mg (s) + S (s) ? MgS (s) (7.7)
Incorporating the reactions (7.5 to 
7.7) within the fold of oxidation reactions  
encouraged chemists to consider not only 
the removal of hydrogen as oxidation, but 
also the removal of electropositive elements 
as oxidation. Thus the reaction : 
2K
4
 [Fe(CN)
6
](aq) + H
2
O
2
 (aq) ?2K
3
[Fe(CN)
6
](aq) 
+ 2 KOH (aq)
is interpreted as oxidation due to the removal 
of electropositive element potassium from 
potassium ferrocyanide before it changes to 
potassium ferricyanide. To summarise, the 
term “oxidation” is defined as the addition 
of oxygen/electronegative element to 
a substance or removal of hydrogen/
electropositive element from a substance.
In the beginning, reduction was considered 
as removal of oxygen from a compound. 
However, the term reduction has been 
broadened these days to include removal 
of oxygen/electronegative element from 
a substance or addition of hydrogen/
electropositive element to a substance. 
According to the definition given above, 
the following are the examples of reduction 
processes:
2 HgO (s)   2 Hg (l) + O
2 
(g) (7.8)
(removal of oxygen from mercuric oxide )
2 FeCl
3
 (aq) + H
2
 (g) ?2 FeCl
2
 (aq) + 2 HCl(aq)
(7.9)
(removal of electronegative element, chlorine 
from ferric chloride)
CH
2 
= CH
2
 (g) + H
2
 (g) ? H
3
C – CH
3
 (g) (7.10)
(addition of hydrogen)
2HgCl
2
 (aq) + SnCl
2
 (aq) ? Hg
2
Cl
2 
(s)+SnCl
4
 (aq)
(7.11)
(addition of mercury to mercuric chloride)
In reaction (7.11) simultaneous oxidation 
of stannous chloride to stannic chloride is 
also occurring because of the addition of 
electronegative element chlorine to it. It was 
soon realised that oxidation and reduction 
always occur simultaneously (as will be 
apparent by re-examining all the equations 
given above), hence, the word “redox” was 
coined for this class of chemical reactions. 
Problem 7.1
In the reactions given below, identify 
the species undergoing oxidation and 
reduction: 
(i)  H
2
S (g) + Cl
2
 (g) ? 2 HCl (g) + S (s)
(ii) 3Fe
3
O
4
 (s) + 8 Al (s) ? 9 Fe (s) 
                                                 + 4Al
2
O
3
 (s)
(iii) 2 Na (s) + H
2
 (g) ? 2 NaH (s)
Solution  
(i) H
2
S is oxidised because a more 
electronegative element, chlorine is added 
to hydrogen (or a more electropositive 
element, hydrogen has been removed  
from S). Chlorine is reduced due to 
addition of hydrogen to it.
(ii)  Aluminium is oxidised because 
oxygen is added to it. Ferrous ferric oxide 
Unit 7.indd   236 10/10/2022   10:37:03 AM
2024-25
237 redox reactions
(Fe
3
O
4
) is reduced because oxygen has 
been removed from it. 
(iii)  With the careful application of the 
concept of electronegativity only we 
may infer that sodium is oxidised and 
hydrogen is reduced. 
Reaction (iii) chosen here prompts us to 
think in terms of another way to define 
redox reactions. 
7.2 REDOX REACTIONS IN TERMS OF 
ELECTRON  TRANSFER REACTIONS
We have already learnt that the reactions 
2Na(s) + Cl
2
(g) ?  2NaCl (s) (7.12)
4Na(s) + O
2
(g) ?  2Na
2
O(s) (7.13)
2Na(s) + S(s) ?  Na
2
S(s) (7.14)
are redox reactions because in each of these 
reactions sodium is oxidised due to the addition 
of either oxygen or more electronegative 
element to sodium. Simultaneously, chlorine, 
oxygen and sulphur are reduced because to 
each of these, the electropositive element 
sodium has been added. From our knowledge 
of chemical bonding we also know that sodium 
chloride, sodium oxide and sodium sulphide 
are ionic compounds and perhaps better 
written as Na
+
Cl
–
 (s), (Na
+
)
2
O
2–
(s), and (Na
+
)
2
 
S
2–
(s). Development of charges on the species 
produced suggests us to rewrite the reactions 
(7.12 to 7.14) in the following manner : 
For convenience, each of the above 
processes can be considered as two separate 
steps, one involving the loss of electrons 
and the other the gain of electrons. As an 
illustration, we may further elaborate one of 
these, say, the formation of  sodium chloride. 
2 Na(s) ? 2 Na
+
(g)
 
  +   2e
–
Cl
2
(g) + 2e
–
 ? 2 Cl
–
(g)
Each of the above steps is called a half 
reaction, which explicitly shows involvement 
of electrons. Sum of the half reactions gives 
the overall reaction : 
2 Na(s) + Cl
2
 (g)  ? 2 Na
+
 Cl
–
 (s) or 2 NaCl (s)
Reactions 7.12 to 7.14 suggest that half 
reactions that involve loss of electrons are 
called oxidation reactions. Similarly, the 
half reactions that involve gain of electrons 
are called reduction reactions.  It may not 
be out of context to mention here that the 
new way of defining oxidation and reduction 
has been achieved only by establishing a 
correlation between the behaviour of species 
as per the classical idea and their interplay 
in electron-transfer change. In reactions (7.12 
to 7.14) sodium, which is oxidised, acts as  
a reducing agent because it donates electron 
to each of the elements interacting with it and 
thus helps in reducing them. Chlorine, oxygen 
and sulphur are reduced and act as oxidising 
agents because these accept electrons from 
sodium. To summarise, we may mention that 
Oxidation : Loss of electron(s) by any species.
Reduction : Gain of electron(s) by any species.
Oxidising agent : Acceptor of electron(s).
Reducing agent : Donor of electron(s).
Problem 7.2  Justify that the reaction:
2 Na(s) + H
2
(g) ?  2 NaH (s) is a redox 
change.
Solution
Since in the above reaction the compound 
formed is an ionic compound, which may 
also be represented as Na
+
H
–
 (s), this 
suggests that one half reaction in this 
process is : 
2 Na (s) ? 2 Na
+
(g)  +   2e
–
Unit 7.indd   237 10/10/2022   10:37:03 AM
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238 chemistry and the other half reaction is: 
H
2
 (g) + 2e
–
 ?  2 H
–
(g)
This splitting of the reaction under 
examination into two half reactions 
automatically reveals that here sodium 
is oxidised and hydrogen is reduced, 
therefore, the complete reaction is a 
redox change. 
7.2.1 Competitive Electron Transfer 
Reactions
Place a strip of metallic zinc in an aqueous 
solution of copper nitrate as shown in Fig. 
7.1, for about one hour. You may notice 
that the  strip becomes coated with reddish 
metallic copper and the blue colour of the 
solution disappears. Formation of  Zn
2+
 ions 
among the products can easily be  judged 
when the blue colour of the solution due to 
Cu
2+
 has disappeared. If hydrogen sulphide 
gas is passed through the colourless solution 
containing Zn
2+
 ions, appearance of white zinc 
sulphide, ZnS can be seen on making the 
solution alkaline with ammonia. 
The reaction between metallic zinc and the 
aqueous solution of copper nitrate is : 
Zn(s) + Cu
2+
 (aq) ? Zn
2+
 (aq) + Cu(s)     (7.15)
In reaction (7.15), zinc has lost electrons 
to form Zn
2+ 
and, therefore, zinc is oxidised. 
Evidently, now if zinc is oxidised, releasing 
electrons, something must be reduced, 
accepting the electrons lost by zinc. Copper 
ion is reduced by gaining electrons from the zinc. 
Reaction (7.15) may be rewritten as : 
At this stage we may investigate the state 
of equilibrium for the reaction represented by 
equation (7.15). For this purpose, let us place 
a strip of metallic copper in a zinc sulphate 
solution. No visible reaction is noticed and 
attempt to detect the presence of Cu
2+
 ions 
by passing H
2
S gas through the solution to 
produce the black colour of cupric sulphide, 
CuS, does not succeed. Cupric sulphide has 
such a low solubility that this is an extremely 
sensitive test; yet the amount of Cu
2+
 formed 
cannot be detected. We thus conclude that 
the state of equilibrium for the reaction (7.15) 
greatly favours the products over the reactants. 
Let us extend electron transfer reaction 
now to copper metal and silver nitrate solution 
in water and arrange a set-up as shown in  
Fig. 7.2. The solution develops blue colour 
due to the formation of Cu
2+
 ions on account 
of the reaction:
Fig. 7.1  Redox reaction between zinc and aqueous solution of copper nitrate occurring in a beaker.
(7.16)
Here, Cu(s) is oxidised to Cu
2+
(aq) and  
Ag
+
(aq) is reduced to Ag(s). Equilibrium greatly 
favours the products Cu
2+
 (aq) and Ag(s).
By way of contrast, let us also compare 
the reaction of metallic cobalt placed in nickel 
sulphate solution. The reaction that occurs 
here is : 
(7.17)
Unit 7.indd   238 11/11/2022   09:48:49
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Page 5


235 redox reactions
Chemistry deals with varieties of matter and change 
of one kind of matter into the other. Transformation of 
matter from one kind into another occurs through the 
various types of reactions. One important category of such 
reactions is Redox Reactions. A number of phenomena, 
both physical as well as biological, are concerned with 
redox reactions. These reactions find extensive use in 
pharmaceutical, biological, industrial, metallurgical and 
agricultural areas. The importance of these reactions is 
apparent from the fact that burning of different types of 
fuels for obtaining energy for domestic, transport and 
other commercial purposes, electrochemical processes 
for extraction of highly reactive metals and non-metals, 
manufacturing of  chemical  compounds like  caustic 
soda, operation of dry and wet batteries and corrosion of 
metals fall within the purview of redox processes. Of late, 
environmental issues like Hydrogen Economy (use of 
liquid hydrogen as fuel) and development of ‘Ozone Hole’ 
have started figuring under redox phenomenon. 
7.1	 CLASSICAL 	 IDEA 	 OF 	 REDOX 	 REACTIONS 	 –	 	
OXIDATION AND REDUCTION REACTIONS
Originally, the term oxidation was used to describe the 
addition of oxygen to an element or a compound. Because 
of the presence of dioxygen in the atmosphere (~20%), 
many elements combine with it and this is the principal 
reason why they commonly occur on the earth in the  
form of their oxides. The following reactions represent 
oxidation processes according to the limited definition of 
oxidation: 
2 Mg (s)  +  O
2
 (g)  ?  2 MgO (s) (7.1)
S (s) + O
2
 (g)  ?  SO
2
 (g) (7.2)
After studying this unit you will be 
able to 
• identify redox reactions as a class 
of reactions in which oxidation 
and reduction reactions occur 
simultaneously; 
• define the terms oxidation, 
reduction, oxidant (oxidising 
agent) and reductant (reducing 
agent); 
• explain mechanism of redox 
reactions by electron transfer 
process;
• use  the concept of oxidation 
number to identify oxidant and 
reductant in a reaction;
• classify redox reaction into 
combination (synthesis), 
decomposition, displacement 
and disproportionation 
reactions; 
• suggest a comparative order 
among various reductants and 
oxidants; 
• balance chemical equations 
using (i) oxidation number  
(ii) half reaction method;
• learn the concept of redox 
reactions in terms of electrode 
processes.
UNIT 7
REDOX REACTIONS
Where there is oxidation, there is always reduction –  
Chemistry is essentially a study of redox systems.
Unit 7.indd   235 10/10/2022   10:37:02 AM
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236 chemistry In reactions (7.1) and (7.2), the elements 
magnesium and sulphur are oxidised on 
account of addition of oxygen to them. 
Similarly,  methane is oxidised owing to the 
addition of oxygen to it.
CH
4
 (g) + 2O
2
 (g) ? CO
2
 (g) + 2H
2
O (l) (7.3)
A careful examination of reaction (7.3) in 
which hydrogen has been replaced by oxygen 
prompted chemists to reinterpret oxidation 
in terms of removal of hydrogen from it and, 
therefore, the scope of term oxidation was 
broadened to include the removal of hydrogen 
from a substance. The following illustration is 
another reaction where removal of hydrogen 
can also be cited as an oxidation reaction. 
2 H
2
S(g) + O
2
 (g) ? 2 S (s) + 2 H
2
O (l) (7.4)
As knowledge of chemists grew, it was 
natural to extend the term oxidation for 
reactions similar to (7.1 to 7.4), which do 
not involve oxygen but other electronegative 
elements. The oxidation of magnesium with 
fluorine, chlorine and sulphur etc. occurs 
according to the following reactions : 
Mg (s) + F
2
 (g) ? MgF
2
 (s) (7.5)
Mg (s) + Cl
2
 (g) ? MgCl
2
 (s) (7.6)
Mg (s) + S (s) ? MgS (s) (7.7)
Incorporating the reactions (7.5 to 
7.7) within the fold of oxidation reactions  
encouraged chemists to consider not only 
the removal of hydrogen as oxidation, but 
also the removal of electropositive elements 
as oxidation. Thus the reaction : 
2K
4
 [Fe(CN)
6
](aq) + H
2
O
2
 (aq) ?2K
3
[Fe(CN)
6
](aq) 
+ 2 KOH (aq)
is interpreted as oxidation due to the removal 
of electropositive element potassium from 
potassium ferrocyanide before it changes to 
potassium ferricyanide. To summarise, the 
term “oxidation” is defined as the addition 
of oxygen/electronegative element to 
a substance or removal of hydrogen/
electropositive element from a substance.
In the beginning, reduction was considered 
as removal of oxygen from a compound. 
However, the term reduction has been 
broadened these days to include removal 
of oxygen/electronegative element from 
a substance or addition of hydrogen/
electropositive element to a substance. 
According to the definition given above, 
the following are the examples of reduction 
processes:
2 HgO (s)   2 Hg (l) + O
2 
(g) (7.8)
(removal of oxygen from mercuric oxide )
2 FeCl
3
 (aq) + H
2
 (g) ?2 FeCl
2
 (aq) + 2 HCl(aq)
(7.9)
(removal of electronegative element, chlorine 
from ferric chloride)
CH
2 
= CH
2
 (g) + H
2
 (g) ? H
3
C – CH
3
 (g) (7.10)
(addition of hydrogen)
2HgCl
2
 (aq) + SnCl
2
 (aq) ? Hg
2
Cl
2 
(s)+SnCl
4
 (aq)
(7.11)
(addition of mercury to mercuric chloride)
In reaction (7.11) simultaneous oxidation 
of stannous chloride to stannic chloride is 
also occurring because of the addition of 
electronegative element chlorine to it. It was 
soon realised that oxidation and reduction 
always occur simultaneously (as will be 
apparent by re-examining all the equations 
given above), hence, the word “redox” was 
coined for this class of chemical reactions. 
Problem 7.1
In the reactions given below, identify 
the species undergoing oxidation and 
reduction: 
(i)  H
2
S (g) + Cl
2
 (g) ? 2 HCl (g) + S (s)
(ii) 3Fe
3
O
4
 (s) + 8 Al (s) ? 9 Fe (s) 
                                                 + 4Al
2
O
3
 (s)
(iii) 2 Na (s) + H
2
 (g) ? 2 NaH (s)
Solution  
(i) H
2
S is oxidised because a more 
electronegative element, chlorine is added 
to hydrogen (or a more electropositive 
element, hydrogen has been removed  
from S). Chlorine is reduced due to 
addition of hydrogen to it.
(ii)  Aluminium is oxidised because 
oxygen is added to it. Ferrous ferric oxide 
Unit 7.indd   236 10/10/2022   10:37:03 AM
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237 redox reactions
(Fe
3
O
4
) is reduced because oxygen has 
been removed from it. 
(iii)  With the careful application of the 
concept of electronegativity only we 
may infer that sodium is oxidised and 
hydrogen is reduced. 
Reaction (iii) chosen here prompts us to 
think in terms of another way to define 
redox reactions. 
7.2 REDOX REACTIONS IN TERMS OF 
ELECTRON  TRANSFER REACTIONS
We have already learnt that the reactions 
2Na(s) + Cl
2
(g) ?  2NaCl (s) (7.12)
4Na(s) + O
2
(g) ?  2Na
2
O(s) (7.13)
2Na(s) + S(s) ?  Na
2
S(s) (7.14)
are redox reactions because in each of these 
reactions sodium is oxidised due to the addition 
of either oxygen or more electronegative 
element to sodium. Simultaneously, chlorine, 
oxygen and sulphur are reduced because to 
each of these, the electropositive element 
sodium has been added. From our knowledge 
of chemical bonding we also know that sodium 
chloride, sodium oxide and sodium sulphide 
are ionic compounds and perhaps better 
written as Na
+
Cl
–
 (s), (Na
+
)
2
O
2–
(s), and (Na
+
)
2
 
S
2–
(s). Development of charges on the species 
produced suggests us to rewrite the reactions 
(7.12 to 7.14) in the following manner : 
For convenience, each of the above 
processes can be considered as two separate 
steps, one involving the loss of electrons 
and the other the gain of electrons. As an 
illustration, we may further elaborate one of 
these, say, the formation of  sodium chloride. 
2 Na(s) ? 2 Na
+
(g)
 
  +   2e
–
Cl
2
(g) + 2e
–
 ? 2 Cl
–
(g)
Each of the above steps is called a half 
reaction, which explicitly shows involvement 
of electrons. Sum of the half reactions gives 
the overall reaction : 
2 Na(s) + Cl
2
 (g)  ? 2 Na
+
 Cl
–
 (s) or 2 NaCl (s)
Reactions 7.12 to 7.14 suggest that half 
reactions that involve loss of electrons are 
called oxidation reactions. Similarly, the 
half reactions that involve gain of electrons 
are called reduction reactions.  It may not 
be out of context to mention here that the 
new way of defining oxidation and reduction 
has been achieved only by establishing a 
correlation between the behaviour of species 
as per the classical idea and their interplay 
in electron-transfer change. In reactions (7.12 
to 7.14) sodium, which is oxidised, acts as  
a reducing agent because it donates electron 
to each of the elements interacting with it and 
thus helps in reducing them. Chlorine, oxygen 
and sulphur are reduced and act as oxidising 
agents because these accept electrons from 
sodium. To summarise, we may mention that 
Oxidation : Loss of electron(s) by any species.
Reduction : Gain of electron(s) by any species.
Oxidising agent : Acceptor of electron(s).
Reducing agent : Donor of electron(s).
Problem 7.2  Justify that the reaction:
2 Na(s) + H
2
(g) ?  2 NaH (s) is a redox 
change.
Solution
Since in the above reaction the compound 
formed is an ionic compound, which may 
also be represented as Na
+
H
–
 (s), this 
suggests that one half reaction in this 
process is : 
2 Na (s) ? 2 Na
+
(g)  +   2e
–
Unit 7.indd   237 10/10/2022   10:37:03 AM
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238 chemistry and the other half reaction is: 
H
2
 (g) + 2e
–
 ?  2 H
–
(g)
This splitting of the reaction under 
examination into two half reactions 
automatically reveals that here sodium 
is oxidised and hydrogen is reduced, 
therefore, the complete reaction is a 
redox change. 
7.2.1 Competitive Electron Transfer 
Reactions
Place a strip of metallic zinc in an aqueous 
solution of copper nitrate as shown in Fig. 
7.1, for about one hour. You may notice 
that the  strip becomes coated with reddish 
metallic copper and the blue colour of the 
solution disappears. Formation of  Zn
2+
 ions 
among the products can easily be  judged 
when the blue colour of the solution due to 
Cu
2+
 has disappeared. If hydrogen sulphide 
gas is passed through the colourless solution 
containing Zn
2+
 ions, appearance of white zinc 
sulphide, ZnS can be seen on making the 
solution alkaline with ammonia. 
The reaction between metallic zinc and the 
aqueous solution of copper nitrate is : 
Zn(s) + Cu
2+
 (aq) ? Zn
2+
 (aq) + Cu(s)     (7.15)
In reaction (7.15), zinc has lost electrons 
to form Zn
2+ 
and, therefore, zinc is oxidised. 
Evidently, now if zinc is oxidised, releasing 
electrons, something must be reduced, 
accepting the electrons lost by zinc. Copper 
ion is reduced by gaining electrons from the zinc. 
Reaction (7.15) may be rewritten as : 
At this stage we may investigate the state 
of equilibrium for the reaction represented by 
equation (7.15). For this purpose, let us place 
a strip of metallic copper in a zinc sulphate 
solution. No visible reaction is noticed and 
attempt to detect the presence of Cu
2+
 ions 
by passing H
2
S gas through the solution to 
produce the black colour of cupric sulphide, 
CuS, does not succeed. Cupric sulphide has 
such a low solubility that this is an extremely 
sensitive test; yet the amount of Cu
2+
 formed 
cannot be detected. We thus conclude that 
the state of equilibrium for the reaction (7.15) 
greatly favours the products over the reactants. 
Let us extend electron transfer reaction 
now to copper metal and silver nitrate solution 
in water and arrange a set-up as shown in  
Fig. 7.2. The solution develops blue colour 
due to the formation of Cu
2+
 ions on account 
of the reaction:
Fig. 7.1  Redox reaction between zinc and aqueous solution of copper nitrate occurring in a beaker.
(7.16)
Here, Cu(s) is oxidised to Cu
2+
(aq) and  
Ag
+
(aq) is reduced to Ag(s). Equilibrium greatly 
favours the products Cu
2+
 (aq) and Ag(s).
By way of contrast, let us also compare 
the reaction of metallic cobalt placed in nickel 
sulphate solution. The reaction that occurs 
here is : 
(7.17)
Unit 7.indd   238 11/11/2022   09:48:49
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239 redox reactions
Fig. 7.2  Redox reaction between copper and aqueous solution of silver nitrate occurring in a beaker. 
At equilibrium, chemical tests reveal that both 
Ni
2+
(aq) and Co
2+
(aq)
 
 are present at moderate 
concentrations. In this case, neither the 
reactants [Co(s) and Ni
2+
(aq)] nor the products 
[Co
2+
(aq) and Ni (s)] are greatly favoured. 
This competition for release of electrons 
incidently reminds us of the competition for 
release of  protons among acids.  The similarity 
suggests that we might develop a table in 
which metals and their ions are listed  on the 
basis of their tendency to release electrons 
just as we do in the case of acids to indicate 
the strength of the acids. As a matter of fact 
we have already made certain comparisons. 
By comparison we have come to know that 
zinc releases electrons to copper and copper 
releases electrons to silver and, therefore, 
the electron releasing tendency of the metals 
is in the order: Zn>Cu>Ag. We would love to 
make our list more vast and design a metal 
activity series or electrochemical series. 
The competition for electrons between various 
metals helps us to design a class of cells, 
named as Galvanic cells in which the chemical 
reactions become the source of electrical 
energy. We would study more about these 
cells in Class XII.
7.3 OXIDATION NUMBER
A less obvious example of electron transfer is 
realised when hydrogen combines with oxygen 
to form water by the reaction:
2H
2
(g) + O
2 
(g) ?  2H
2
O (l)      (7.18)
Though not simple in its approach, yet 
we can visualise the H atom as going from a 
neutral (zero) state in H
2
 to a positive state in 
H
2
O, the O atom goes from a zero state in O
2
 
to a dinegative state in H
2
O. It is assumed that 
there is an electron transfer from H to O and 
consequently H
2
 is oxidised and O
2
 is reduced. 
However, as we shall see later, the charge 
transfer is only partial and is perhaps better 
described as an electron shift rather than a 
complete loss of electron by H and gain by 
O. What has been said here with respect 
to equation (7.18) may be true for a good 
number of other reactions involving covalent 
compounds. Two such examples of this class 
of the reactions are:
H
2
(s) + Cl
2
(g)   ?  2HCl(g)     (7.19)
and,
CH
 4
(g)  +  4Cl
2
(g)  ? CCl
4
(l)  + 4HCl(g)    (7.20)
In order to keep track of electron shifts 
in chemical reactions involving formation 
of covalent compounds, a more practical 
method of using oxidation number has 
been developed. In this method, it is always 
assumed that there is a complete transfer 
of electron from a less electronegative atom 
to a more electonegative atom. For example, 
we rewrite equations (7.18 to 7.20) to show 
charge on each of the atoms forming part of 
the reaction : 
  0           0              +1 –2
2H
2
(g) + O
2
(g) ?  2H
2
O (l) (7.21)
0            0            +1 –1
H
2
 (s) + Cl
2
(g) ?  2HCl(g) (7.22)
–4+1          0          +4 –1        +1 –1 
CH
4
(g)  + 4Cl
2
(g) ?  CCl
4
(l)  +4HCl(g) (7.23)
It may be emphasised that the assumption 
of electron transfer is made for book-keeping 
purpose only and it will become obvious at 
a later stage in this unit that it leads to the 
simple description of redox reactions. 
Oxidation number denotes the oxidation 
state of an element in a compound 
ascertained according to a set of rules 
formulated on the basis that electron pair 
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