Study

1. Lecture- 1
Chemical Bonding
Oxidation States

2. Lets Learn

3. Lecture- 1
Lets Learn

4. Lecture- 1
Its there in humans too !!!

5. Lecture- 1
PHOTOSYNTHESIS

6. For Example :
(i)
(ii)
(iii)
(iv)
C + O2
2Mg+O2
Zn + S
Mg + Cl2
-------------> CO2
-------------> 2MgO
-------------> ZnS
-------------> MgCl2
}Addition of oxygen
}Addition of electronegative element
Lecture- 1
REDOX REACTIONS-CLASSICAL IDEA
OXIDATION:
The addition of oxygen or any other electronegative element or removal of
hydrogen or any other electropositive element.

7. For Example :
(i)
(ii)
(iii)
(iv)
C + O2
2Mg+O2
Zn + S
Mg + Cl2
-------------> CO2
-------------> 2MgO
-------------> ZnS
-------------> MgCl2
}Addition of oxygen
}Addition of electronegative element
Lecture- 1
REDOX REACTIONS-CLASSICAL IDEA
OXIDATION:
The addition of oxygen or any other electronegative element or removal of
hydrogen or any other electropositive element.

8. Lecture- 1
(i) Cl2 + H2
(ii) Br2 + H2S
------------->
------------->
2HCl
2HBr + S
(iii)
(iv)
2FeCl3 + Fe
2HgCl2 + SnCl2
------------->
------------->
2FeCl2
Hg2Cl2 + SnCl4
}Addition of hydrogen
}Addition of electropositive element
Reduction
The addition of hydrogen or any other electropositive element or removal of
oxygen or any other electronegative element.
For example, the reduction reactions are :

9. Lecture- 1
Oxidizing Agent or Oxidant
A substance which gives oxygen or removes hydrogen is called an
oxidizing agent or oxidant.

10. Lecture- 1
A substance which provides hydrogen or removes oxygen is called a
reducing agent or reductant.
ReducingAgent

11. Lecture- 1
Oxidation-Reduction
Oxidation-Reduction Reactions are Complementary : go side by side
2HgCl2 + SnCl2 -----------------> Hg2Cl2 + SnCl4
Thus, oxidation and reduction reactions are collectively called redox
reactions.

12. Lecture- 1
Oxidation-Reduction
Oxidation-Reduction Reactions are Complementary : go side by side
2HgCl2 + SnCl2 -----------------> Hg2Cl2 + SnCl4
Thus, oxidation and reduction reactions are collectively called redox
reactions.

13. Lecture- 1
Na+ + e–
Mg2+ + 2e–
Sn4+ + 2e–
Na
Mg
Sn2+
Fe2+
--------------->
--------------->
--------------->
---------------> Fe3+ + e–
(ii) Loss of electrons resulting in decrease of negative charge.
MnO
[Fe(CN)6]4–
--------------->
--------------->
MnO + e–
[Fe(CN)6]3– + e–
Electronic Concept
Oxidation is a process which involves loss of electrons by an atom or group
of atoms.
For example,
(i) Loss of electrons resulting in increase in positive charge.

14. Lecture- 1
--------------->
--------------->
--------------->
Fe2+
Sn2+
Sb3+
(iii) Gain of electrons resulting in increase of negative charge.
Cl 2e–
S + 2e–
---------------> 2Cl–
---------------> S2–
Electronic Concept
Reduction : A process which involves gain of electrons by an atom or group
of atoms.
(i) Gain of electrons resulting in decrease of positive charge.
Fe3+ + e–
Sn4+ + 2e–
Sb5+ + 2e–

15. Lecture- 1
Lets Understand !!
OXIDATION-REDUCTION REACTIONS AS ELECTRONS TRANSFER
REACTIONS
Oxidation takes place at the cost of reduction and vice versa.

16. Lecture- 1
Lets Learn

17. Lecture- 1
Lets Understand !!
OXIDATION NUMBER
The charge which an atom of the element has in its ion or appears to have
when present in the combined state with other atoms.
Oxidation numbers are also called oxidation states.
Oxidation number also gives the charge which an atom appears to have
when all other atoms are removed from it as ions.

18. Lecture- 1
Rules for the Determination of Oxidation Number of an Atom
1. The oxidation number of an element in the free or elementary state or in
any of its allotropic forms is always zero.

19. Lecture- 1
Rules for the Determination of Oxidation Number of an Atom
1. The oxidation number of an element in the free or elementary state or in
any of its allotropic forms is always zero.

20. Lecture- 1
Rules for the Determination of Oxidation Number of an Atom
2. The oxidation number of an element in a single (monoatomic) ion is the
same as the charge on the ion.

21. Lecture- 1
Rules for the Determination of Oxidation Number of an Atom
2. The oxidation number of an element in a single (monoatomic) ion is the
same as the charge on the ion.

22. Lecture- 1
Rules for the Determination of Oxidation Number of an Atom
3. In compounds formed by the combination of non-metallic atoms, the atom
with higher electronegativity is given negative oxidation number.

23. Lecture- 1
Rules for the Determination of Oxidation Number of an Atom
4. In all compounds of hydrogen, the oxidation number of hydrogen is +1
except in hydrides of active metals such as LiH, NaH, KH, MgH2, CaH2,
etc., where hydrogen has the oxidation number of –1.

24. Lecture- 1
7.
8. The second exception is found in compounds in which oxygen is bonded to
fluorine.
Rules for the Determination of Oxidation Number of an Atom
5. The oxidation number of oxygen is –2 in most of the compounds.
6. However, there are two exceptions. The first exception is peroxides and
superoxides

25. Lecture- 1
Rules for the Determination of Oxidation Number of an Atom
7. The most electronegative element, fluorine has oxidation number –1. For
other halogens, the oxidation number is generally – 1, but there are
exceptions when these are bonded to a more electronegative halogen atom
or oxygen.

26. Lecture- 1
For a polyatomic ion, the sum of the oxidation numbers of all the atoms is
equal to charge on the ion.
Rules for the Determination of Oxidation Number of an Atom
8. For neutral molecule, the sum of the oxidation numbers of all the atoms is
equal to zero.

27. Lecture- 1
For a polyatomic ion, the sum of the oxidation numbers of all the atoms is
equal to charge on the ion.
Rules for the Determination of Oxidation Number of an Atom
8. For neutral molecule, the sum of the oxidation numbers of all the atoms is
equal to zero.

28. Lecture- 1
Oxidation Number of
Free elements
Fluorine
Simple ions
Oxygen
F2O (+ 2); F2O2 (+ 1)
Hydrogen = + 1;
metal hydrides (–1)
Sum of O.N. of atoms in molecules = 0
Sum of O.N. of atoms in polyatomic ions = (Charge on them).
Watch out !!
= 0
= –1
= Charge on them
= OXIDES (–2);
peroxides (–1);