This slide will help the students to learn about the basic concepts of redox reactions and their uses.

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Oxidation and reduction
Sheikh Mahatabuddin, Ph. D.
Associate Professor
Department of NFE
Daffodil International University
smuddin.nfe@diu.edu.bd
mahatab.chem@gmail.com

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Oxidation and Reduction
Oxidation: (Classical Concepts)
Addition of oxygen => C + O2 ⇒ CO2
Addition of electronegative elements => Na + Cl2 ⇒ NaCl
Removal of electropositive elements => Zn + CuSO4 ⇒ ZnSO4 + Cu
Removal of Hydrogen => Na + HCl ⇒ NaCl + H2
Electronic Concept:
Removal or Donation of electrons: Na => Na+ + e-
Reduction: (Classical Concepts)
Removal of oxygen => Fe2O3 + CO ⇒ Fe + CO2
Removal of electronegative elements => H2S + Cl2 ⇒ HCl + S
Addition of electropositive elements => Zn + CuSO4 ⇒ ZnSO4 + Cu
Addition of Hydrogen => H2 + Cl2 ⇒ HCl
Electronic Concept:
Addition or Acceptance of electrons: Cl + e
- => Cl
-
©DSM

3. 3
Fe2O3 + CO ⇒ Fe + CO2
Reduction
Oxidizing agent
Oxidation
Reducing agent
1. O removed from Fe2O3 => Reduction of Fe2O3
2. On the other hand, Fe2O3 oxidize CO therefore
Fe2O3 is a Oxidizing agent!
1. O is added to CO=> Oxidation of CO
2. On the other hand, Fe2O3 Reduced to Fe
therefore CO is a Reducing agent
Red Oxi Oxi Red
Reducing agent get
oxidized to reduce
the oxidizing agent
Oxidizing agent get
reduced to oxidize
the reducing agent
Reducing agents: All metals, NaBH4, LiAlH4, H2
Oxidizing agents: Nonmetals except H2, KMnO4, K2Cr2O7, H2O2
Oxidation Reduction Reactions
Takes Place Simultaneously
Na => Na+ + e- Oxidation
Cl + e- => Cl- Reduction
©DSM

4. 4
Oxidation Number: Oxidation number, also called Oxidation State, the total
number of electrons that an atom either gains (-ve) or loses (+ve) in order to
form a chemical bond with another atom.
Rules for calculating oxidation number:
1. The oxidation number of an atom is zero in a neutral substance that contains atoms of
only one element. Thus, the atoms in O2, O3, P4, S8, and aluminum metal all have an
oxidation number of 0.
2. The oxidation number of simple ions is equal to the charge on the ion. The oxidation
number of sodium in the Na+ ion is +1, for example, and the oxidation number of chlorine
in the Cl- ion is -1.
3. The oxidation number of hydrogen is +1 when it is combined with a nonmetal as in CH4,
NH3, H2O, and HCl.
4. The oxidation number of hydrogen is -1 when it is combined with a metal as in. LiH,
NaH, CaH2, and LiAlH4.
5. The metals in Group IA form compounds (such as Li3N and Na2S) in which the metal
atom has an oxidation number of +1.
6. The elements in Group IIA form compounds (such as Mg3N2 and CaCO3) in which the
metal atom has a +2 oxidation number.
©DSM

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7. Oxygen usually has an oxidation number of -2. Exceptions include molecules and polyatomic
ions that contain O-O bonds, such as O2, O3, H2O2, and the O2
2- ion.
8. The elements in Group VIIA often form compounds (such as AlF3, HCl, and ZnBr2) in which the
nonmetal has a -1 oxidation number.
9. The sum of the oxidation numbers in a neutral compound is zero. H2O: 2(+1) + (-2) = 0
10. The sum of the oxidation numbers in a polyatomic ion is equal to the charge on the ion. The
oxidation number of the sulfur atom in the SO4
2- ion must be +6, for example, because the sum of
the oxidation numbers of the atoms in this ion must equal -2. SO4
2-: (+6) + 4(-2) = -2
11. Elements toward the bottom left corner of the periodic table are more likely to have positive
oxidation numbers than those toward the upper right corner of the table. Sulfur has a positive
oxidation number in SO2, for example, because it is below oxygen in the periodic table.
SO2: (+4) + 2(-2) = 0
FeCl2 + Cl2 ⇒ FeCl3
Fe oxidation number changes to +2 to +3
Oxidation
Cl oxidation number changes to 0 to -1
Reduction
Increase in +ve oxidation number
Decrease in –ve oxidation number
Increase in – ve oxidation number
Decrease in +ve oxidation number
©DSM

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Find out the oxidation and reducing agents from following reactions:-
1. Mg + PbO MgO + Pb
2. CuO + C Cu + CO2
3. Fe2O3 + C Fe + CO
4. Mg + CuO MgO + Cu
5. Mg + O2 MgO
6. Al + Fe2O3 Fe + Al2O3
7. CuO + H2 Cu + H2O
8. Cu + AgNO3 CuNO3 + Ag
9. Ce4+ + Fe2+ Ce3+ + Fe3+
10. Sn4+ + Cr2+ Sn2+ + Cr3+
©DSM

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The method used to balance redox reactions is called the Half Equation Method. The
equation is separated into two half-equations; one for oxidation and one for reduction.
Each equation is balanced by adjusting coefficients and adding H2O, H+, and e- in this
order:
1. Balance elements in the equation other than O and H.
2. Balance the oxygen atoms by adding the appropriate number of water (H2O) molecules
to the opposite side of the equation.
3. Balance the hydrogen atoms (including those added in step 2 to balance the oxygen
atom) by adding H+ ions to the opposite side of the equation.
4. Add up the charges on each side. Make them equal by adding enough electrons (e-) to
the more positive side. (Rule of thumb: e- and H+ are almost always on the same side.)
5. The e- on each side must be made equal; if they are not equal, they must be multiplied
by appropriate integers (the lowest common multiple) to be made the same.
6. The half-equations are added together, canceling out the electrons to form one balanced
equation. Common terms should also be canceled out.
• (If the equation is being balanced in a basic solution, through the addition of one more
step, the appropriate number of OH- must be added to turn the remaining H+ into water
molecules.)
• The equation can now be checked to make sure that it is balanced.
Balancing Oxidation-Reduction reactions:
©DSM

8. 8
Balance the following redox reaction in acidic conditions.
FeSO4 (aq) + Cr2O7
2- (aq) → Cr3+ (aq) + Fe3+ (aq)
Here the change of oxidation state:
Fe2+ Fe3+ increase in positive oxidation number, therefore oxidation
Cr6+ Cr3+ Decrease in positive oxidation number, therefore reduction
Reduction half reaction
Cr2O7
2- + 6e- => 2Cr3+ + 7H2O – (i)
Oxidation half reaction
Fe2+ => Fe3+ + e- as we need 6 e- while the oxidation
process provide 1e- therefore multiplying the oxidation
half reaction by 6 => 6 Fe2+ => 6 Fe3+ + 6 e- ------ (ii)
Addition of equation i with ii
Cr2O7
2- + 6e- => 2Cr3+ + 7H2O
6 Fe2+ => 6 Fe3+ + 6 e- (Addition)
Cr2O7
2- + 6 Fe2+ = 2 Cr3+ + 6 Fe3+ + 7H2O---(iii)
As it was a acidic solution and FeSO4
used as oxidizing agent, therefore to
provide opposite ion to the equation iii
K2Cr2O7 + 6 FeSO4 + 7 H2SO4 = K2SO4 + Cr2(SO4)3 + 3 Fe2(SO4)3 + 7H2O
After providing opposite ions the balanced equation will be--
©DSM

9. 9
ASSIGNMENTS
For each of the following, separate the skeletal (unbalanced) equation into two half reactions. For each half reaction,
balance the elements (mass balance), and then add electrons to the right or left side to make a net charge balance.
Identify which half reaction is the oxidation and which is the reduction. Then, multiply each half reaction by an
appropriate factor so that the two multiplied half reactions add together to make a balanced redox equation.
1.Hg2
+2+S2O2
–3→ Hg + S4O2
–6 Hg2
2++ S2O3
2–→ Hg+S4O6
2–
2. Al + Cr3+→Al3++Cr2+ Al+ Cr3+→Al3+ + Cr2+
3. Au3+ + I– → Au + I2
4.
5. a) Mn2+ (aq) + BiO3
- (aq) ⇒ MnO4
- (aq) + Bi3+(aq) [acidic]
b) Fe(CN)6
3- (aq) + Re(s) ⇒ Fe(CN)6
4- (aq) + ReO4
- (aq) [basic]
Balance I– + MnO4
2 –→ IO–3+MnO2I–+MnO4
2– → IO3– + MnO2 in basic aqueous solution.
©DSM

10. 10
Dependence of the electropositivity and oxidation potential for the displacement
of any metal ions in a solution
The electrochemical series
Equilibrium Eo (Volts)
Li (aq) + e- ⇌ Li (s) - 3.03
K+ (aq) + e- ⇌ K (s) - 2.92
Ca2+ (aq) + 2e- ⇌ Ca (s) - 2.87
Na+ (aq) + e- ⇌ Na (s) - 2.71
Mg2+ (aq) + 2e- ⇌ Mg (s) - 2.37
Al3+ (aq) + 3e- ⇌ Al (s) - 1.66
Zn2+ (aq) + 2e- ⇌ Zn (s) - 0.76
Fe2+ (aq) + 2e- ⇌ Fe (s) - 0.44
Pb2+ (aq) + 2e- ⇌ Pb (s) - 0.13
2 H+ (aq) + 2e- ⇌ Ca (s) 0
Cu2+ (aq) + 2e- ⇌ Cu (s) + 0.34
Ag+ (aq) + e- ⇌ Ag (s) + 0.80
Au3+ (aq) + 3e- ⇌ Au (s) + 1.50
A series of electrodes or half cells
arranged in order of their increasing
standard oxidation potentials or in the
decreasing order of their standard
reduction potentials is called an
electromotive force series or
electrochemical series is also known as
electromotive force series e. m. f. series.
©DSM
https://thefactfactor.com/facts/pure_science/chemistry/physical-chemistry/electrochemical-series/5877/

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