8. REDOX REACTIONS
Redox reactions involve oxidation and reduction. The important concepts relating to redox reactions are:
I. Classical Concept: According to this concept oxidation is the process of addition of oxygen/electronegative element
to a substance or removal of hydrogen/electropositive element from a substance.
Reduction is the process of removal of oxygen/electronegative element from a substance or addition of
hydrogen/electropositive element to a substance.
Substance which is oxidised is called reducing agent and the substance which is reduced is called oxidising agent.
If oxidation and reduction take place simultaneously, the process is called Redox reaction.
i.e. reduction + oxidation → Redox reac ons
E.g. Zn + CuO → ZnO + Cu
Here Zn is converted to ZnO. i.e oxygen is added to Zn. So it is oxidised and hence the reducing agent. CuO is converted
to Cu. i.e. oxygen is removed from Cu. So it is reduced and hence it is the oxidising agent.
Other examples are:
1. FeCl3 + H2 → FeCl2 + 2HCl
Here the electronegative Cl atom is removed from FeCl3. So it is reduced. H2 is oxidised since an electronegative Cl atom
is added to it. FeCl3 is the oxidising agent and H2 is the reducing agent.
2. 2 H2S(g) + O2 (g) → 2 S (s) + 2 H2O (l)
Here H2S is oxidised and O2 is reduced.
II. Electronic Concept: According to this concept oxidation is the process of removal (losing) of electron and reduction is
the process of addition (gaining) of electron. A redox reaction is the process of exchange of electrons between two or more
substances.
A substance that accepts electron is called oxidising agent and a substance that donates electron is called a reducing
agent.
E.g. In the reaction Zn + Cu2+ → Zn2+ + Cu, Zn loses two electrons and forms Zn2+. So it is oxidised. Cu2+ gains two
electrons and forms Cu. So it is reduced. Here Zn is the reducing agent and Cu2+ is the oxidising agent.
Other examples are:
1. Reaction between Cu and Ag+
Here Cu loses two electrons. So it is oxidised and is the reducing
agent. Ag+ accepts an electron. So it is reduced and is the oxidising agent.
In the first example, Cu is reduced while in the second reaction it is oxidised. So the oxidation or reduction of a metal
depends on the nature of the metal to which it is combined.
The series in which the different metals are arranged in the decreasing order of their reactivity is called
electrochemical series or reactivity series. Generally, a metal lying above in the reactivity series can displace another metal
from its salt solution. For example Zn can displace copper from an aqueous solution of copper sulphate, since Zn lies above
Cu in the electrochemical series.
2. Reaction between cobalt and nickel ion.
Here Co is oxidised to Co2+ and Ni2+ is reduced to Ni.
Oxidation number (Oxidation state)
Oxidation number of an element in a compound is the residual charge on the element when all the other atoms are
removed from it as ions. For example oxidation number of Mn in KMnO4 is the residual charge on Mn when one K atom and
four O atoms removed from it as K+ and O2- ions respectively.
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, Rules used for the calculation of oxidation number
1. The oxidation number of all elements in the free or the uncombined state is zero. For e.g. oxidation number of H2,
O2, Cl2, O3, P4, S8, Na, Mg, Al etc. is zero.
2. For simple ions, the oxidation number is equal to the charge on the ion. Thus Na+ ion has an oxidation number of +1,
Mg2+ ion +2, Fe3+ ion +3, Cl– ion –1, O2– ion –2 and so on.
3. All alkali metals have oxidation number of +1 and all alkaline earth metals have an oxidation number of +2.
Aluminium shows an oxidation number of +3 in all of its compounds.
4. The common oxidation number of oxygen is –2. But in peroxides (e.g., H2O2, Na2O2), oxidation number of oxygen is –1
and in superoxides (e.g., KO2, RbO2), it is –½. In oxygen difluoride (OF2) and dioxygen difluoride (O2F2), the oxygen is
assigned an oxidation number of +2 and +1 respectively.
5. The common oxidation number of hydrogen is +1. But in hydrides, H shows an oxidation number of -1.
6. The common oxidation number of halogens is -1. Fluorine shows only -1 oxidation number in all of its compounds.
But other halogens show positive oxidation numbers also in their oxides and oxoacids.
7. The algebraic sum of the oxidation number of all the atoms in a compound is zero.
8. In polyatomic ion, the sum of the oxidation numbers of all the atoms is equal to the charge on the ion.
Stock Notations
Alfred Stock proposed some notations to represent the oxidation number of a metal in a compound.
According to this, the oxidation number is represented in Roman numeral in brackets after the symbol of the metal in the
molecular formula. Thus aurous chloride and auric chloride are written as Au(I)Cl and Au(III)Cl3. Similarly, stannous chloride
and stannic chloride are written as Sn(II)Cl2 and Sn(IV)Cl4.
III. Oxidation number Concept: According to this concept, oxidation is the process of increase in the oxidation number
of an element and reduction is the process of decrease in the oxidation number of an element. A reagent that can increase
the oxidation number of an element in a given substance is called oxidising agent or oxidant and a reagent which lowers the
oxidation number of an element in a given substance is called reducing agent or reductant. A redox reaction is a reaction
which involves change in oxidation number of the interacting species.
For e.g. in the reaction , the oxidation number of Zn increases from
0 to +2 and that of Cu in CuSO4 decreases from +2 to 0. So Zn is oxidised and Cu in CuSO4 is reduced.
In the reaction , the oxidation number of Cr decreases from +3 to 0.
So it is reduced and is the oxidising agent. The oxidation number of Al increases from 0 to +3. So it is oxidised and is the
reducing agent.
Types of redox reactions
1. Combination reactions: A combination reaction may be denoted as A + B → C
Here either A or B or both A and B must be in the elemental form. All combustion reactions are combination redox reactions,
since here one of the reactants is O2. Examples are:
2. Decomposition reactions
Decomposition reactions are the opposite of combination reactions. It involves the breakdown of a compound
into two or more components, in which at least one must be in the elemental state. It may be denoted as: C → A + B.
Examples are:
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Redox reactions involve oxidation and reduction. The important concepts relating to redox reactions are:
I. Classical Concept: According to this concept oxidation is the process of addition of oxygen/electronegative element
to a substance or removal of hydrogen/electropositive element from a substance.
Reduction is the process of removal of oxygen/electronegative element from a substance or addition of
hydrogen/electropositive element to a substance.
Substance which is oxidised is called reducing agent and the substance which is reduced is called oxidising agent.
If oxidation and reduction take place simultaneously, the process is called Redox reaction.
i.e. reduction + oxidation → Redox reac ons
E.g. Zn + CuO → ZnO + Cu
Here Zn is converted to ZnO. i.e oxygen is added to Zn. So it is oxidised and hence the reducing agent. CuO is converted
to Cu. i.e. oxygen is removed from Cu. So it is reduced and hence it is the oxidising agent.
Other examples are:
1. FeCl3 + H2 → FeCl2 + 2HCl
Here the electronegative Cl atom is removed from FeCl3. So it is reduced. H2 is oxidised since an electronegative Cl atom
is added to it. FeCl3 is the oxidising agent and H2 is the reducing agent.
2. 2 H2S(g) + O2 (g) → 2 S (s) + 2 H2O (l)
Here H2S is oxidised and O2 is reduced.
II. Electronic Concept: According to this concept oxidation is the process of removal (losing) of electron and reduction is
the process of addition (gaining) of electron. A redox reaction is the process of exchange of electrons between two or more
substances.
A substance that accepts electron is called oxidising agent and a substance that donates electron is called a reducing
agent.
E.g. In the reaction Zn + Cu2+ → Zn2+ + Cu, Zn loses two electrons and forms Zn2+. So it is oxidised. Cu2+ gains two
electrons and forms Cu. So it is reduced. Here Zn is the reducing agent and Cu2+ is the oxidising agent.
Other examples are:
1. Reaction between Cu and Ag+
Here Cu loses two electrons. So it is oxidised and is the reducing
agent. Ag+ accepts an electron. So it is reduced and is the oxidising agent.
In the first example, Cu is reduced while in the second reaction it is oxidised. So the oxidation or reduction of a metal
depends on the nature of the metal to which it is combined.
The series in which the different metals are arranged in the decreasing order of their reactivity is called
electrochemical series or reactivity series. Generally, a metal lying above in the reactivity series can displace another metal
from its salt solution. For example Zn can displace copper from an aqueous solution of copper sulphate, since Zn lies above
Cu in the electrochemical series.
2. Reaction between cobalt and nickel ion.
Here Co is oxidised to Co2+ and Ni2+ is reduced to Ni.
Oxidation number (Oxidation state)
Oxidation number of an element in a compound is the residual charge on the element when all the other atoms are
removed from it as ions. For example oxidation number of Mn in KMnO4 is the residual charge on Mn when one K atom and
four O atoms removed from it as K+ and O2- ions respectively.
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, Rules used for the calculation of oxidation number
1. The oxidation number of all elements in the free or the uncombined state is zero. For e.g. oxidation number of H2,
O2, Cl2, O3, P4, S8, Na, Mg, Al etc. is zero.
2. For simple ions, the oxidation number is equal to the charge on the ion. Thus Na+ ion has an oxidation number of +1,
Mg2+ ion +2, Fe3+ ion +3, Cl– ion –1, O2– ion –2 and so on.
3. All alkali metals have oxidation number of +1 and all alkaline earth metals have an oxidation number of +2.
Aluminium shows an oxidation number of +3 in all of its compounds.
4. The common oxidation number of oxygen is –2. But in peroxides (e.g., H2O2, Na2O2), oxidation number of oxygen is –1
and in superoxides (e.g., KO2, RbO2), it is –½. In oxygen difluoride (OF2) and dioxygen difluoride (O2F2), the oxygen is
assigned an oxidation number of +2 and +1 respectively.
5. The common oxidation number of hydrogen is +1. But in hydrides, H shows an oxidation number of -1.
6. The common oxidation number of halogens is -1. Fluorine shows only -1 oxidation number in all of its compounds.
But other halogens show positive oxidation numbers also in their oxides and oxoacids.
7. The algebraic sum of the oxidation number of all the atoms in a compound is zero.
8. In polyatomic ion, the sum of the oxidation numbers of all the atoms is equal to the charge on the ion.
Stock Notations
Alfred Stock proposed some notations to represent the oxidation number of a metal in a compound.
According to this, the oxidation number is represented in Roman numeral in brackets after the symbol of the metal in the
molecular formula. Thus aurous chloride and auric chloride are written as Au(I)Cl and Au(III)Cl3. Similarly, stannous chloride
and stannic chloride are written as Sn(II)Cl2 and Sn(IV)Cl4.
III. Oxidation number Concept: According to this concept, oxidation is the process of increase in the oxidation number
of an element and reduction is the process of decrease in the oxidation number of an element. A reagent that can increase
the oxidation number of an element in a given substance is called oxidising agent or oxidant and a reagent which lowers the
oxidation number of an element in a given substance is called reducing agent or reductant. A redox reaction is a reaction
which involves change in oxidation number of the interacting species.
For e.g. in the reaction , the oxidation number of Zn increases from
0 to +2 and that of Cu in CuSO4 decreases from +2 to 0. So Zn is oxidised and Cu in CuSO4 is reduced.
In the reaction , the oxidation number of Cr decreases from +3 to 0.
So it is reduced and is the oxidising agent. The oxidation number of Al increases from 0 to +3. So it is oxidised and is the
reducing agent.
Types of redox reactions
1. Combination reactions: A combination reaction may be denoted as A + B → C
Here either A or B or both A and B must be in the elemental form. All combustion reactions are combination redox reactions,
since here one of the reactants is O2. Examples are:
2. Decomposition reactions
Decomposition reactions are the opposite of combination reactions. It involves the breakdown of a compound
into two or more components, in which at least one must be in the elemental state. It may be denoted as: C → A + B.
Examples are:
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