Redox Reaction
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Redox Reactions
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Introduction
The reaction in which oxidation and reduction both occur simultaneously are
called Redox reactions.
For Example: A reaction between hydrogen and fluorine to form hydrogen
fluoride. During the reaction, hydrogen oxidizes and loses two electrons,
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wherein fluorine gains two electrons.
Types of Redox Reactions:
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The different types of redox reactions are
Decomposition Reaction
Combination Reaction
Displacement Reaction
Disproportionation Reactions
(i) Decomposition Reaction:
This kind of reaction involves the breakdown of a compound into different
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compounds. Examples of these types of reactions are:
2H 2 O →2H2 + O 2
(ii) Combination Reaction:
These reactions are the opposite of decomposition reactions and hence, involve
the combination of two compounds to form a single compound in the form of
A + B → AB
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For example,
H 2 + Cl2 -> 2HClC+O2 ->CO 2
(iii)Displacement Reaction:
In this kind of reaction, an atom or an ion in a compound is replaced by an
atom or an ion of another element. It can be represented in the form of
X + YZ → XZ + Y
For Example: Iron displaces the copper metal when added to a copper sulphate
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solution.
Oxidation:
Oxidation is defined as the addition of oxygen/electronegative
element to a substance or removal of hydrogen/ electropositive
element from a substance. 2
Example: 2Mg(s) + O 2(g)----> 2Mg)(s)
Mg(s) + Cl2 (g)----> MgCl 2 (s)
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Reduction:
Reduction is defined as the removal of oxygen/electronegative element from a
substance or addition of hydrogen or electropositive element to a substance.
Example: 2FeCl3 (aq) + H2 (g) ------> 2FeCl2 (aq) + 2HCL(aq)
Redox Reaction in Terms of Electron Transfer Reaction:
A few examples of redox reaction on the basis of electronic concept are given
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below:
According to electronic concept every redox reaction consists of two steps known
as half reactions.
(i) Oxidation reaction: Half reactions that involve loss of electrons are called
oxidation reactions.
(ii) Reduction reaction: Half reactions that involve gain of electrons are called
reduction reactions.
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the substance that undergo oxidation itself and reduce other is called reducing
agent whereas the substance that undergo reduction itself and oxidize other is
called oxidizing agent.
For example:- CuO + H 2 --> Cu + H 2 O
Copper Oxide Hydrogen Copper Water
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In reaction, Cu is undergoing reduction and oxidizing hydrogen. Therefore, Cu is
oxidizing agent and H is reducing agent.
Oxidizing and Reducing Agents:
1. The substance (atom, ion and molecule) that gains electrons and is thereby
reduced to a low valency state is called an oxidising agent.
2. The substance that loses electrons and is thereby oxidised to a higher
valency state is called a reducing agent.
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Remember
Important Oxidizing Agents:
1. Molecules are made up of electronegative elements. Eg: O 2 , O 3 , and X 2
(halogens)
2. Compounds containing an element that is in the higher oxidized state. Eg:
KMnO 4, K 2 Cv 2O7 , HNO 3 , KClo 3
3. Oxides of metals and non-metals. Eg: MgO, CuO, CrO 3
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4. Fluorine is the strongest oxidizing agent.
Important Reducing Agents:
1. All metals, for example, Na, Zn, Fe, Al
2. A few non-metals, such as C, Hydrogen, S, P
3. Metallic hydrides, including NaH, LiH, CaH2 , etc.
Identification of Oxidizing
and Reducing Agents
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If an element is in its higher possible oxidation state in a compound, it can
function as an oxidising agent.
Eg: KMnO4 , K 2 Cr2 O7 , HNO3 , H 2 SO4 and HClO4
If an element is in its possible lower oxidation state in a compound, it can
function as a reducing agent. Eg: H 2 S, H 2 C2 O4 , FeSO 4 and SnCl 2 .
The compound will act as an oxidising agent if a highly electronegative
element is in its highest oxidation state.
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Oxidation number
It is the oxidation state of an element in a compound which is the charge
assigned to an atom of a compound is equal to the number of electrons in the
valence shell of an atom that are gained or lost completely or to a large extent
by that atom while forming a bond in a compound.
Rules to assign and calculate oxidation number:
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1. The oxidation number of atoms in their elemental state is taken as zero.
2. The oxidation number of mono-atomic atoms like Na+ etc is taken as 1.
3. The oxidation number of Hydrogen is +1 when present with non metals and -1
when present with metals.
4. The oxidation number of oxygen is -2 in most of the compounds but in
peroxides it is -1.
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5. The metals always have oxidation number in positive and non metal in
negative when present together in ionic compounds.
6. In compounds that have two atoms with different electro negativities, the
oxidation number of more electronegative is taken as –ve and for less
electronegative it is taken as positive.
7. In neutral compounds, the sum of all oxidation numbers is equal to zero.
8. In complex ions, the sum of oxidation states of all the atoms is equal to the
charge present on the complex.
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Redox reaction in terms of oxidation number:
In this, Zinc is undergoing oxidation and act as reducing agent whereas
Hydrogen is undergoing reduction and act as an oxidizing agent.
2+
Zn 0 + H + Cl - -------> Zn + H 2
Balancing of Redox reactions
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There are two ways of balancing Redox reactions:
Oxidation number method
Half equation method
(i) Oxidation Number Method. Following steps are involved:
(ii) Write the correct formula for each reactant and product.
(b) By assigning the oxidation change in oxidation number can be identified.
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(c) Calculate the increase and decrease in oxidation number per atom with
respect to the reactants. If more than one atom is present then multiply by
suitable coefficient.
(d) Balance the equation with respect to all atoms. Balance hydrogen and
oxygen atoms also.
(e) If the reaction is carried out in acidic medium, use H+ ions in the equation.
If it is in basic medium use OH– ions.
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(f) Hydrogen atoms in the expression can be balanced by adding (H 20) molecules
to the reactants or products.
If there are the same number of oxygen atoms on the both side of equation
then it represents the balanced redox reaction.
(ii) Half Reaction Method. In this method two half equation are balanced
separately and than added together to give balanced equation.
Example:
2FeCl 3 (aq) + H2 (g)-----> 2FeCl 2 (aq) + 2HCL(aq)
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Redox couple
It is defined as having together the oxidized and reduced forms of a
substance taking part in oxidation or reduction half reactions .they are
actually conjugate acid base pairs .
Redox reaction as the basis for Titrations:
Titration is the process in which the solutions of two reagents are allowed to
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react with each other.
Procedure:
In it one solution (known volume) is taken in Burette and the solution is
called titrant.
The other reagent is taken in flask called titration flask and the solution is
called as analyte.
The titration is carried out till both the reagents mix completely.
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The stage at which both the reagent mix completely is called end point.
The end point is detected by indicator.
The objective of these titrations is to find out the exact amount of an acid (or
the base) present in a given solution by reacting it against the solution of
standard base (or an acid) .
Application of Redox Reaction:
Chemicals like Caustic Soda and Chlorine are made using redox reactions
Ammonia undergoes an oxidation reaction to form a nitric oxide which is an YouTube/@padhleakshay
essential component of fertilizers
Cutting an apple into half oxidizes it and changes the colour to brown.
NCERT QUESTION
Q1.Assign oxidation number to the underline element of each of
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the following component-
1. K2 MnO4
Ans: Let X be the oxidation number of Mn
Oxidation no. of K= +1
Oxidation no. of O= -2
Then,
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2(+1) + x + 4(-2) = 0
2+x–8=0
x= +6
Therefore, the oxidation number of Mn is +6
Q2. Nitric acid is an oxidising agent and reacts with PbO, but it
does not react with PbO2. Explain why?
Ans: Nitric acid is an oxidizing agent and reacts with PbO to give a
simple acid-base reaction without any change in oxidation state. In YouTube/@padhleakshay
PbO2 , Pb is in +4 oxidation state and cannot be oxidized further;
hence no reaction takes place between PbO 2 and HNO 3 .
Q3. Explain redox reactions based on electron transfer. Give suitable
examples.
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Ans: In a redox reaction if one species loses electrons it’s considered to be
undergoing oxidation reaction and acts as oxidizing agent or oxidant,
and for species which accept electrons is said to undergo reduction and
behave as reductant.
For example, Zinc and HCl reaction
Zn + 2HCl ----> ZnCl2 + H 2
zinc loses electrons to the electronegative atom Cl with the reaction for
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oxidation and reduction as follows:
Oxidation: Zn--->Zn 2+ + 2e-
Reduction: 2H + + 2e- ---->H 2
Thus the transfer of electrons causes the redox reaction to occur.
Q4. MnO42- undergoes disproportionation reaction in acidic medium but MnO4-
does not. Give reason.
Ans: In – MnO 2 , Mn is in the highest oxidation state i.e. +7. YouTube/@padhleakshay
Therefore, it does
not undergo disproportionation. MnO 2-
4 undergoes disproportionation
as follows :
2- - -
3MnO4 + 4H+ --->2MnO 4 + MnO 2 2H 2 O
Q5. Calculate the oxidation number of phosphorus in the following species.
(a) HPO-3
2
and PO 2-
4
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Ans:
HPO -32 x= +3
PO 2-4 x=+5
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After studying from this notes u got
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