NCERT Textbook: Chemical Bonding & Molecular Structure | Chemistry Class 11 - NEET PDF Download

 Page 1


Scientists are constantly discovering new compounds, 
orderly arranging the facts about them, trying to explain 
with the existing knowledge, organising to modify the 
earlier views or evolve theories for explaining the newly 
observed facts.
Unit 4
After studying this Unit, you will be 
able to 
•	 unde r s t a nd 	 Kös s e l - Le w i s	
approach to chemical bonding; 
•	 e x pl a i n 	 t he 	 oc t e t 	 r ul e 	 and 	 i t s	
limitations, 	 draw 	 Lewis 	 structures	
of simple molecules;
•	 explain 	 the 	 formation 	 of	 different	
types of bonds;
•	 describe 	 the	 VSEPR	 theory	 and	
predict the geometry of simple 
molecules;
•	 explain 	 the 	 valence 	 bond	
approach for the formation of 
covalent 	bonds;
•	 predict 	 the	 directional 	 properties	
of	covalent	bonds;
•	 explain 	 the 	 different 	 types 	 of	 	
hybridisation 	 	 involving 	s, p and 
d orbitals and draw shapes of 
simple 	covalent	molecules;
•	 describe 	 the 	 molecular 	 orbital	
theory of homonuclear diatomic 
molecules;
•	 explain 	 the	 concept 	 of 	 hydrogen	
bond.
CHEMiCAL BOnDinG AnD 
MOLECULAR St RUCt URE
Matter is made up of one or different type of elements. 
Under	 normal 	 conditions	 no 	 other	 element	 exists	 as	 an	
independent	 atom	 in 	 nature,	 except 	 noble	 gases.	 However,	
a	 group	 of	atoms 	 is	 found 	 to 	 exist 	 together	 as	 one	 species	
having	 characteristic	 properties. 	 Such 	 a 	 group 	 of	 atoms 	
is	 called	 a	 molecule. 	 Obviously	 there 	 must	 be 	 some	 force	
which holds these constituent atoms together in the 
molecules. the attractive force which holds various 
constituents (atoms, ions, etc.) together in different 
chemical species is called a chemical bond.	 Since 	 the	
formation of chemical compounds takes place as a result of 
combi nat i on 	 of 	 at om s 	 of 	 var i ous 	 el em ent s 	 i n 	 di f f er ent 	 w ays,	
it raises many questions. Why do atoms combine? Why are 
only certain combinations possible? Why do some atoms 
combine while certain others do not? Why do molecules 
possess 	 definite 	 shapes? 	 To 	 answer 	 such 	 questions 	 different	
theories	 and	 concepts 	 have	 been 	 put	 forward 	 from 	 time 	
to	 time. 	 These	 are 	 Kössel-Lewis 	 approach,	 Valence	 Shell 	
Electron 	 Pair 	 R epulsion 	 (VSEPR ) 	 Theory, 	 Valence 	 Bond 	 (V B)	
Theory	 and	 Molecular 	 Orbital 	 (MO) 	 Theory.	 The	 evolution	
of 	 various 	 theories 	 of 	 valence 	 and 	 the 	 interpretation 	 of	
the	 nature	 of	 chemical 	 bonds 	 have 	 closely	 been	 related	 to 	
the	 developments 	 in	 the 	 understanding	 of 	 the	 structure 	
of	 atom,	 the	 electronic	 configuration 	 of 	 elements	 and	 the	
periodic	 table. 	 Every 	 system	 tends	 to 	 be	 more 	 stable	 and 	
bonding is nature’s way of lowering the energy of the system 
to attain stability.
Unit 4.indd   100 9/12/2022   9:36:09 AM
2024-25
Page 2


Scientists are constantly discovering new compounds, 
orderly arranging the facts about them, trying to explain 
with the existing knowledge, organising to modify the 
earlier views or evolve theories for explaining the newly 
observed facts.
Unit 4
After studying this Unit, you will be 
able to 
•	 unde r s t a nd 	 Kös s e l - Le w i s	
approach to chemical bonding; 
•	 e x pl a i n 	 t he 	 oc t e t 	 r ul e 	 and 	 i t s	
limitations, 	 draw 	 Lewis 	 structures	
of simple molecules;
•	 explain 	 the 	 formation 	 of	 different	
types of bonds;
•	 describe 	 the	 VSEPR	 theory	 and	
predict the geometry of simple 
molecules;
•	 explain 	 the 	 valence 	 bond	
approach for the formation of 
covalent 	bonds;
•	 predict 	 the	 directional 	 properties	
of	covalent	bonds;
•	 explain 	 the 	 different 	 types 	 of	 	
hybridisation 	 	 involving 	s, p and 
d orbitals and draw shapes of 
simple 	covalent	molecules;
•	 describe 	 the 	 molecular 	 orbital	
theory of homonuclear diatomic 
molecules;
•	 explain 	 the	 concept 	 of 	 hydrogen	
bond.
CHEMiCAL BOnDinG AnD 
MOLECULAR St RUCt URE
Matter is made up of one or different type of elements. 
Under	 normal 	 conditions	 no 	 other	 element	 exists	 as	 an	
independent	 atom	 in 	 nature,	 except 	 noble	 gases.	 However,	
a	 group	 of	atoms 	 is	 found 	 to 	 exist 	 together	 as	 one	 species	
having	 characteristic	 properties. 	 Such 	 a 	 group 	 of	 atoms 	
is	 called	 a	 molecule. 	 Obviously	 there 	 must	 be 	 some	 force	
which holds these constituent atoms together in the 
molecules. the attractive force which holds various 
constituents (atoms, ions, etc.) together in different 
chemical species is called a chemical bond.	 Since 	 the	
formation of chemical compounds takes place as a result of 
combi nat i on 	 of 	 at om s 	 of 	 var i ous 	 el em ent s 	 i n 	 di f f er ent 	 w ays,	
it raises many questions. Why do atoms combine? Why are 
only certain combinations possible? Why do some atoms 
combine while certain others do not? Why do molecules 
possess 	 definite 	 shapes? 	 To 	 answer 	 such 	 questions 	 different	
theories	 and	 concepts 	 have	 been 	 put	 forward 	 from 	 time 	
to	 time. 	 These	 are 	 Kössel-Lewis 	 approach,	 Valence	 Shell 	
Electron 	 Pair 	 R epulsion 	 (VSEPR ) 	 Theory, 	 Valence 	 Bond 	 (V B)	
Theory	 and	 Molecular 	 Orbital 	 (MO) 	 Theory.	 The	 evolution	
of 	 various 	 theories 	 of 	 valence 	 and 	 the 	 interpretation 	 of	
the	 nature	 of	 chemical 	 bonds 	 have 	 closely	 been	 related	 to 	
the	 developments 	 in	 the 	 understanding	 of 	 the	 structure 	
of	 atom,	 the	 electronic	 configuration 	 of 	 elements	 and	 the	
periodic	 table. 	 Every 	 system	 tends	 to 	 be	 more 	 stable	 and 	
bonding is nature’s way of lowering the energy of the system 
to attain stability.
Unit 4.indd   100 9/12/2022   9:36:09 AM
2024-25
101
Chemi Cal Bonding a nd m ole Cular Stru Cture 4.1 KÖSSEL-LEwiS AppROACH t O 
CHEMiCAL BOnDinG  
In	 order 	 to	 explain 	 the 	 formation 	 of	 chemical	
bond in terms of electrons, a number of 
attempts were made, but it was only in 
1916 	 when 	 Kössel 	 and 	 Lewis 	 succeeded	
independently 	 in 	 giving 	 a 	 satisfactory	
explan ation .	 Th ey 	 were 	 th e 	 first 	 to 	 provide 	
some 	 logical 	 explanation 	 of 	 valence 	 which 	 was	
based on the inertness of noble gases. 
Lewis 	 pictured 	 the 	 atom 	 in 	 terms 	 of 	 a	
positively 	 charged 	 ‘Kerne l’ 	 (the 	 nucleus	 plus	
the	 inner 	 electrons) 	 and	 the	 outer	 shell	 that	
could 	 accommodate 	 a 	 maximum 	 of 	 eight	
e l e c t r ons . 	 H e , 	 f ur t he r 	 as s um e d 	 t ha t 	 t he s e	
eight electrons occupy the corners of a cube 
which 	 surround	 the	 ‘Kernel’.	 Thus	 the	 single	
outer shell electron of sodium would occupy 
one corner of the cube, while in the case of 
a noble gas all the eight corners would be 
occupied. 	 This	 octet	 of 	 electrons,	 represents	
a particularly stable electronic arrangement. 
Lewis postulated that atoms achieve 
the stable octet when they are linked by 
chemical bonds. In the case of sodium and 
chlorine, this can happen by the transfer of 
an electron from sodium to chlorine thereby 
giving 	 the 	 Na
+
 and Cl
–
 ions. In the case of 
other molecules like Cl
2
,	 H
2
, F
2
, etc., the bond 
is formed by the sharing of a pair of electrons 
between the atoms. In the process each atom 
attains a stable outer octet of electrons.
Lewis Symbols: In the formation of a 
molecule, only the outer shell electrons take 
part in chemical combination and they are 
known as valence electrons .	 The	 inner	 shell	
electrons are well protected and are generally 
not 	 i nvol ved 	 i n 	 t he 	 combi nat i on 	 pr ocess. 
G.N.	 Lewis, 	 an	 American	 chemist 	 introduced	
simple 	 notations 	 to 	 represent 	 valence 	 electrons	
in 	 an 	 atom.	 These	 notations 	 are 	 called 	Lewis 
symbols . 	 For	 example,	 the	 Lewis 	 symbols	 for	
the elements of second period are as under:
Significance of Lewis Symbols : 	 The	
number of dots around the symbol represents 
the 	 number 	 of 	 valence 	 electrons. 	 This 	 number	
of 	 valence 	 electrons 	 helps 	 to 	 calculate 	 the	
common or group valence of the element. 
The 	 gr o up 	 val e nce 	 of 	 t he 	 el e m e nt s 	 i s 	 ge ne r al l y	
either	 equal	 to 	 the 	 number	 of 	 dots	 in	 Lewis	
symbols or 8 minus the number of dots or 
valence 	 electrons.
Kössel, in relation to chemical bonding, 
drew attention to the following facts:
•	 In the periodic table, the highly 
electronegative 	 halogens	 and	 the	 highly	
electropositive 	 alkali 	 metals 	 are 	 separated	
by the noble gases;
•	 The 	 formation 	 of 	 a 	 negative 	 ion 	 from 	 a	
halogen 	 atom 	 and 	 a 	 positive 	 ion 	 from	
an alkali metal atom is associated with 
the gain and loss of an electron by the 
respective 	atoms;
•	 The 	 negative 	 and 	 positive 	 ions 	 thus	
formed attain stable noble gas electronic 
configurations. 	 The	 noble	 gases 	 (with 	 the	
exception 	 of	 helium 	 which 	 has 	 a 	 duplet 	
of 	 electrons) 	 have 	 a 	 particul arl y 	 stabl e	
outer 	 shell	 configuration 	 of 	 eight 	 (octet) 	
electrons, ns
2
np
6
. 
•	 The 	 negative 	 and 	 positive 	 ions 	 are 	 stabilized	
by electrostatic attraction.
For	 example, 	 the 	 formation 	 of	 NaCl 	 from	
sodium	 and	 chlorine,	 according 	 to 	 the	 above 	
scheme,	can	 be 	 explained	 as:	
Na 																? 							 Na
+
   +    e
–
[Ne] 	3s
1
																			 [Ne] 	
Cl  +  e
–
          ?        Cl
–
[Ne] 	3s
2
	 3p
5
													[Ne] 	3s
2
	3p
6
	or	[Ar] 	
Na
+
  +  Cl
–
     ? 								 NaCl 	or	Na
+
Cl
–
Similarly 	 th e 	 formation 	 of 	 CaF
2
 may be 
shown as:
Ca                ?       Ca
2+
  +  2e
–
[Ar]4s
2
																				[Ar]
F  + e
–
            ?       F
–
[He] 	2s
2
 2p
5
													[He] 	 2s
2
 2p
6
		or	[Ne]
Ca
2+
 + 2F
–
     ?       CaF
2
   or  Ca
2+
(F
– 
)
2
the bond formed, as a result of the 
electrostatic attraction between the 
positive and negative ions was termed as 
Unit 4.indd   101 9/12/2022   9:36:09 AM
2024-25
Page 3


Scientists are constantly discovering new compounds, 
orderly arranging the facts about them, trying to explain 
with the existing knowledge, organising to modify the 
earlier views or evolve theories for explaining the newly 
observed facts.
Unit 4
After studying this Unit, you will be 
able to 
•	 unde r s t a nd 	 Kös s e l - Le w i s	
approach to chemical bonding; 
•	 e x pl a i n 	 t he 	 oc t e t 	 r ul e 	 and 	 i t s	
limitations, 	 draw 	 Lewis 	 structures	
of simple molecules;
•	 explain 	 the 	 formation 	 of	 different	
types of bonds;
•	 describe 	 the	 VSEPR	 theory	 and	
predict the geometry of simple 
molecules;
•	 explain 	 the 	 valence 	 bond	
approach for the formation of 
covalent 	bonds;
•	 predict 	 the	 directional 	 properties	
of	covalent	bonds;
•	 explain 	 the 	 different 	 types 	 of	 	
hybridisation 	 	 involving 	s, p and 
d orbitals and draw shapes of 
simple 	covalent	molecules;
•	 describe 	 the 	 molecular 	 orbital	
theory of homonuclear diatomic 
molecules;
•	 explain 	 the	 concept 	 of 	 hydrogen	
bond.
CHEMiCAL BOnDinG AnD 
MOLECULAR St RUCt URE
Matter is made up of one or different type of elements. 
Under	 normal 	 conditions	 no 	 other	 element	 exists	 as	 an	
independent	 atom	 in 	 nature,	 except 	 noble	 gases.	 However,	
a	 group	 of	atoms 	 is	 found 	 to 	 exist 	 together	 as	 one	 species	
having	 characteristic	 properties. 	 Such 	 a 	 group 	 of	 atoms 	
is	 called	 a	 molecule. 	 Obviously	 there 	 must	 be 	 some	 force	
which holds these constituent atoms together in the 
molecules. the attractive force which holds various 
constituents (atoms, ions, etc.) together in different 
chemical species is called a chemical bond.	 Since 	 the	
formation of chemical compounds takes place as a result of 
combi nat i on 	 of 	 at om s 	 of 	 var i ous 	 el em ent s 	 i n 	 di f f er ent 	 w ays,	
it raises many questions. Why do atoms combine? Why are 
only certain combinations possible? Why do some atoms 
combine while certain others do not? Why do molecules 
possess 	 definite 	 shapes? 	 To 	 answer 	 such 	 questions 	 different	
theories	 and	 concepts 	 have	 been 	 put	 forward 	 from 	 time 	
to	 time. 	 These	 are 	 Kössel-Lewis 	 approach,	 Valence	 Shell 	
Electron 	 Pair 	 R epulsion 	 (VSEPR ) 	 Theory, 	 Valence 	 Bond 	 (V B)	
Theory	 and	 Molecular 	 Orbital 	 (MO) 	 Theory.	 The	 evolution	
of 	 various 	 theories 	 of 	 valence 	 and 	 the 	 interpretation 	 of	
the	 nature	 of	 chemical 	 bonds 	 have 	 closely	 been	 related	 to 	
the	 developments 	 in	 the 	 understanding	 of 	 the	 structure 	
of	 atom,	 the	 electronic	 configuration 	 of 	 elements	 and	 the	
periodic	 table. 	 Every 	 system	 tends	 to 	 be	 more 	 stable	 and 	
bonding is nature’s way of lowering the energy of the system 
to attain stability.
Unit 4.indd   100 9/12/2022   9:36:09 AM
2024-25
101
Chemi Cal Bonding a nd m ole Cular Stru Cture 4.1 KÖSSEL-LEwiS AppROACH t O 
CHEMiCAL BOnDinG  
In	 order 	 to	 explain 	 the 	 formation 	 of	 chemical	
bond in terms of electrons, a number of 
attempts were made, but it was only in 
1916 	 when 	 Kössel 	 and 	 Lewis 	 succeeded	
independently 	 in 	 giving 	 a 	 satisfactory	
explan ation .	 Th ey 	 were 	 th e 	 first 	 to 	 provide 	
some 	 logical 	 explanation 	 of 	 valence 	 which 	 was	
based on the inertness of noble gases. 
Lewis 	 pictured 	 the 	 atom 	 in 	 terms 	 of 	 a	
positively 	 charged 	 ‘Kerne l’ 	 (the 	 nucleus	 plus	
the	 inner 	 electrons) 	 and	 the	 outer	 shell	 that	
could 	 accommodate 	 a 	 maximum 	 of 	 eight	
e l e c t r ons . 	 H e , 	 f ur t he r 	 as s um e d 	 t ha t 	 t he s e	
eight electrons occupy the corners of a cube 
which 	 surround	 the	 ‘Kernel’.	 Thus	 the	 single	
outer shell electron of sodium would occupy 
one corner of the cube, while in the case of 
a noble gas all the eight corners would be 
occupied. 	 This	 octet	 of 	 electrons,	 represents	
a particularly stable electronic arrangement. 
Lewis postulated that atoms achieve 
the stable octet when they are linked by 
chemical bonds. In the case of sodium and 
chlorine, this can happen by the transfer of 
an electron from sodium to chlorine thereby 
giving 	 the 	 Na
+
 and Cl
–
 ions. In the case of 
other molecules like Cl
2
,	 H
2
, F
2
, etc., the bond 
is formed by the sharing of a pair of electrons 
between the atoms. In the process each atom 
attains a stable outer octet of electrons.
Lewis Symbols: In the formation of a 
molecule, only the outer shell electrons take 
part in chemical combination and they are 
known as valence electrons .	 The	 inner	 shell	
electrons are well protected and are generally 
not 	 i nvol ved 	 i n 	 t he 	 combi nat i on 	 pr ocess. 
G.N.	 Lewis, 	 an	 American	 chemist 	 introduced	
simple 	 notations 	 to 	 represent 	 valence 	 electrons	
in 	 an 	 atom.	 These	 notations 	 are 	 called 	Lewis 
symbols . 	 For	 example,	 the	 Lewis 	 symbols	 for	
the elements of second period are as under:
Significance of Lewis Symbols : 	 The	
number of dots around the symbol represents 
the 	 number 	 of 	 valence 	 electrons. 	 This 	 number	
of 	 valence 	 electrons 	 helps 	 to 	 calculate 	 the	
common or group valence of the element. 
The 	 gr o up 	 val e nce 	 of 	 t he 	 el e m e nt s 	 i s 	 ge ne r al l y	
either	 equal	 to 	 the 	 number	 of 	 dots	 in	 Lewis	
symbols or 8 minus the number of dots or 
valence 	 electrons.
Kössel, in relation to chemical bonding, 
drew attention to the following facts:
•	 In the periodic table, the highly 
electronegative 	 halogens	 and	 the	 highly	
electropositive 	 alkali 	 metals 	 are 	 separated	
by the noble gases;
•	 The 	 formation 	 of 	 a 	 negative 	 ion 	 from 	 a	
halogen 	 atom 	 and 	 a 	 positive 	 ion 	 from	
an alkali metal atom is associated with 
the gain and loss of an electron by the 
respective 	atoms;
•	 The 	 negative 	 and 	 positive 	 ions 	 thus	
formed attain stable noble gas electronic 
configurations. 	 The	 noble	 gases 	 (with 	 the	
exception 	 of	 helium 	 which 	 has 	 a 	 duplet 	
of 	 electrons) 	 have 	 a 	 particul arl y 	 stabl e	
outer 	 shell	 configuration 	 of 	 eight 	 (octet) 	
electrons, ns
2
np
6
. 
•	 The 	 negative 	 and 	 positive 	 ions 	 are 	 stabilized	
by electrostatic attraction.
For	 example, 	 the 	 formation 	 of	 NaCl 	 from	
sodium	 and	 chlorine,	 according 	 to 	 the	 above 	
scheme,	can	 be 	 explained	 as:	
Na 																? 							 Na
+
   +    e
–
[Ne] 	3s
1
																			 [Ne] 	
Cl  +  e
–
          ?        Cl
–
[Ne] 	3s
2
	 3p
5
													[Ne] 	3s
2
	3p
6
	or	[Ar] 	
Na
+
  +  Cl
–
     ? 								 NaCl 	or	Na
+
Cl
–
Similarly 	 th e 	 formation 	 of 	 CaF
2
 may be 
shown as:
Ca                ?       Ca
2+
  +  2e
–
[Ar]4s
2
																				[Ar]
F  + e
–
            ?       F
–
[He] 	2s
2
 2p
5
													[He] 	 2s
2
 2p
6
		or	[Ne]
Ca
2+
 + 2F
–
     ?       CaF
2
   or  Ca
2+
(F
– 
)
2
the bond formed, as a result of the 
electrostatic attraction between the 
positive and negative ions was termed as 
Unit 4.indd   101 9/12/2022   9:36:09 AM
2024-25
102 chemistry the electrovalent bond. t he electrovalence 
is thus equal to the number of unit charge(s) 
on the ion. 	 Thus, 	 calcium 	 is 	 assigned 	 a	
positive 	 electrovalence 	 of	 two,	 while	 chlorine	
a 	 negative 	electrovalence	of	one. 	
Kösse l ’s 	 post ul at i ons 	 pr ovi de 	 t he 	 basi s 	 f or	
th e 	 modern	 con cepts 	 regard ing 	 ion-formation 	
by electron transfer and the formation of ionic 
crystallin e 	 compou nd s. 	 His 	 views 	 h ave 	 proved	
to 	 be 	 of 	 great	 value 	 in 	 the 	 understanding	 and	
systematisation of the ionic compounds. At 
the same time he did recognise the fact that 
a 	 large 	 number 	 of	 compounds 	 did 	 not	 fit 	 into	
these concepts.
4.1.1 Octet Rule
Kössel 	 and 	 Lewis 	 in 	 1916 	 developed 	 an	
important theory of chemical combination 
between atoms known as electronic theory 
of chemical bonding. According to this, 
atoms  can combine either by transfer of 
valence 	 electrons	 from 	 one	 atom 	 to 	 another	
(gaining 	 or	 losing)	 or	 by 	 sharing 	 of	 valence	
electrons 	 in	 order	 to	 have 	 an 	 octet	 in	 their	
valence 	shells.	This 	is 	known 	 as 	octet rule.
4.1.2 Covalent Bond
Langmuir (1919) 	 refined 	 the 	 Lewis	
postulations by abandoning the idea of 
the stationary cubical arrangement of the 
octet, and by introducing the term covalent 
bond . 	 The 	 Lewis-Langmuir 	 theory 	 can 	 be	
understood by considering the formation of 
the chlorine molecule, Cl
2
. 	 The	 Cl 	 atom	 with	
electronic 	 configuration, 	 [Ne]3s
2
	 3p
5
, is one 
electron 	 short 	 of 	 the 	 argon 	 configuration. 	
The 	 formation 	 of 	 the 	 Cl 	
2
 molecule can be 
understood in terms of the sharing of a pair 
of electrons between the two chlorine atoms, 
each chlorine atom contributing one electron 
to the shared pair. In the process both 
chlorine atoms attain the outer shell octet of 
the	 nearest 	 noble	gas	 (i.e.,	 argon).	
t he dots represent electrons. Such 
structures are referred to as Lewis dot 
structures.
The 	 Le w i s 	 dot 	 s t r uc t ur es 	 c an 	 be 	 w r i t t en 	 f or	
other molecules also, in which the combining 
atoms 	 may 	 be 	 identical 	 or 	 d i f f e r e n t . 	 T h e	
important conditions being that:
•	 Ea c h 	 bo nd 	 i s 	 f o r m e d 	 a s 	 a 	 r e s ul t 	 o f 	 s ha r i ng	
of an electron pair between the atoms.
•	 Each 	 combining 	 atom 	 contri butes 	 at 	 least	
one electron to the shared pair.
•	 The 	 combining 	 atoms 	 attain 	 the 	 outer-
shell	 noble	 gas	 configurations 	 as 	 a	 result	
of the sharing of electrons.
•	 Thus 	 in	 water	 and	 carbon 	 tetrachloride	
mol ecules, 	 formation 	 of 	 covalent 	 bonds	
can be represented as:
   or Cl – Cl
Covalent bond between two Cl atoms
t hus, when two atoms share one 
electron pair they are said to be joined by 
a single covalent bond. In many compounds 
we	 have	multiple bonds 	 between	 atoms.	 The	
formation 	 of 	 mu ltip le 	 b on ds 	 envisages 	 sh a rin g	
of more than one electron pair between two 
atoms. if two atoms share two pairs of 
electrons, the covalent bond between them 
is called a double bond. For 	 example,	 in	 the	
carbon 	 dioxide	 molecule,	 we 	 have 	 two	 double	
bonds 	 between 	 the 	 carbon 	 and 	 oxygen 	 atoms.	
Similarly 	 in 	 ethene 	 molecule	 the	 two	 carbon	
atoms are joined by a double bond.
Double bonds in CO
2
 molecule
Unit 4.indd   102 9/12/2022   9:36:09 AM
2024-25
Page 4


Scientists are constantly discovering new compounds, 
orderly arranging the facts about them, trying to explain 
with the existing knowledge, organising to modify the 
earlier views or evolve theories for explaining the newly 
observed facts.
Unit 4
After studying this Unit, you will be 
able to 
•	 unde r s t a nd 	 Kös s e l - Le w i s	
approach to chemical bonding; 
•	 e x pl a i n 	 t he 	 oc t e t 	 r ul e 	 and 	 i t s	
limitations, 	 draw 	 Lewis 	 structures	
of simple molecules;
•	 explain 	 the 	 formation 	 of	 different	
types of bonds;
•	 describe 	 the	 VSEPR	 theory	 and	
predict the geometry of simple 
molecules;
•	 explain 	 the 	 valence 	 bond	
approach for the formation of 
covalent 	bonds;
•	 predict 	 the	 directional 	 properties	
of	covalent	bonds;
•	 explain 	 the 	 different 	 types 	 of	 	
hybridisation 	 	 involving 	s, p and 
d orbitals and draw shapes of 
simple 	covalent	molecules;
•	 describe 	 the 	 molecular 	 orbital	
theory of homonuclear diatomic 
molecules;
•	 explain 	 the	 concept 	 of 	 hydrogen	
bond.
CHEMiCAL BOnDinG AnD 
MOLECULAR St RUCt URE
Matter is made up of one or different type of elements. 
Under	 normal 	 conditions	 no 	 other	 element	 exists	 as	 an	
independent	 atom	 in 	 nature,	 except 	 noble	 gases.	 However,	
a	 group	 of	atoms 	 is	 found 	 to 	 exist 	 together	 as	 one	 species	
having	 characteristic	 properties. 	 Such 	 a 	 group 	 of	 atoms 	
is	 called	 a	 molecule. 	 Obviously	 there 	 must	 be 	 some	 force	
which holds these constituent atoms together in the 
molecules. the attractive force which holds various 
constituents (atoms, ions, etc.) together in different 
chemical species is called a chemical bond.	 Since 	 the	
formation of chemical compounds takes place as a result of 
combi nat i on 	 of 	 at om s 	 of 	 var i ous 	 el em ent s 	 i n 	 di f f er ent 	 w ays,	
it raises many questions. Why do atoms combine? Why are 
only certain combinations possible? Why do some atoms 
combine while certain others do not? Why do molecules 
possess 	 definite 	 shapes? 	 To 	 answer 	 such 	 questions 	 different	
theories	 and	 concepts 	 have	 been 	 put	 forward 	 from 	 time 	
to	 time. 	 These	 are 	 Kössel-Lewis 	 approach,	 Valence	 Shell 	
Electron 	 Pair 	 R epulsion 	 (VSEPR ) 	 Theory, 	 Valence 	 Bond 	 (V B)	
Theory	 and	 Molecular 	 Orbital 	 (MO) 	 Theory.	 The	 evolution	
of 	 various 	 theories 	 of 	 valence 	 and 	 the 	 interpretation 	 of	
the	 nature	 of	 chemical 	 bonds 	 have 	 closely	 been	 related	 to 	
the	 developments 	 in	 the 	 understanding	 of 	 the	 structure 	
of	 atom,	 the	 electronic	 configuration 	 of 	 elements	 and	 the	
periodic	 table. 	 Every 	 system	 tends	 to 	 be	 more 	 stable	 and 	
bonding is nature’s way of lowering the energy of the system 
to attain stability.
Unit 4.indd   100 9/12/2022   9:36:09 AM
2024-25
101
Chemi Cal Bonding a nd m ole Cular Stru Cture 4.1 KÖSSEL-LEwiS AppROACH t O 
CHEMiCAL BOnDinG  
In	 order 	 to	 explain 	 the 	 formation 	 of	 chemical	
bond in terms of electrons, a number of 
attempts were made, but it was only in 
1916 	 when 	 Kössel 	 and 	 Lewis 	 succeeded	
independently 	 in 	 giving 	 a 	 satisfactory	
explan ation .	 Th ey 	 were 	 th e 	 first 	 to 	 provide 	
some 	 logical 	 explanation 	 of 	 valence 	 which 	 was	
based on the inertness of noble gases. 
Lewis 	 pictured 	 the 	 atom 	 in 	 terms 	 of 	 a	
positively 	 charged 	 ‘Kerne l’ 	 (the 	 nucleus	 plus	
the	 inner 	 electrons) 	 and	 the	 outer	 shell	 that	
could 	 accommodate 	 a 	 maximum 	 of 	 eight	
e l e c t r ons . 	 H e , 	 f ur t he r 	 as s um e d 	 t ha t 	 t he s e	
eight electrons occupy the corners of a cube 
which 	 surround	 the	 ‘Kernel’.	 Thus	 the	 single	
outer shell electron of sodium would occupy 
one corner of the cube, while in the case of 
a noble gas all the eight corners would be 
occupied. 	 This	 octet	 of 	 electrons,	 represents	
a particularly stable electronic arrangement. 
Lewis postulated that atoms achieve 
the stable octet when they are linked by 
chemical bonds. In the case of sodium and 
chlorine, this can happen by the transfer of 
an electron from sodium to chlorine thereby 
giving 	 the 	 Na
+
 and Cl
–
 ions. In the case of 
other molecules like Cl
2
,	 H
2
, F
2
, etc., the bond 
is formed by the sharing of a pair of electrons 
between the atoms. In the process each atom 
attains a stable outer octet of electrons.
Lewis Symbols: In the formation of a 
molecule, only the outer shell electrons take 
part in chemical combination and they are 
known as valence electrons .	 The	 inner	 shell	
electrons are well protected and are generally 
not 	 i nvol ved 	 i n 	 t he 	 combi nat i on 	 pr ocess. 
G.N.	 Lewis, 	 an	 American	 chemist 	 introduced	
simple 	 notations 	 to 	 represent 	 valence 	 electrons	
in 	 an 	 atom.	 These	 notations 	 are 	 called 	Lewis 
symbols . 	 For	 example,	 the	 Lewis 	 symbols	 for	
the elements of second period are as under:
Significance of Lewis Symbols : 	 The	
number of dots around the symbol represents 
the 	 number 	 of 	 valence 	 electrons. 	 This 	 number	
of 	 valence 	 electrons 	 helps 	 to 	 calculate 	 the	
common or group valence of the element. 
The 	 gr o up 	 val e nce 	 of 	 t he 	 el e m e nt s 	 i s 	 ge ne r al l y	
either	 equal	 to 	 the 	 number	 of 	 dots	 in	 Lewis	
symbols or 8 minus the number of dots or 
valence 	 electrons.
Kössel, in relation to chemical bonding, 
drew attention to the following facts:
•	 In the periodic table, the highly 
electronegative 	 halogens	 and	 the	 highly	
electropositive 	 alkali 	 metals 	 are 	 separated	
by the noble gases;
•	 The 	 formation 	 of 	 a 	 negative 	 ion 	 from 	 a	
halogen 	 atom 	 and 	 a 	 positive 	 ion 	 from	
an alkali metal atom is associated with 
the gain and loss of an electron by the 
respective 	atoms;
•	 The 	 negative 	 and 	 positive 	 ions 	 thus	
formed attain stable noble gas electronic 
configurations. 	 The	 noble	 gases 	 (with 	 the	
exception 	 of	 helium 	 which 	 has 	 a 	 duplet 	
of 	 electrons) 	 have 	 a 	 particul arl y 	 stabl e	
outer 	 shell	 configuration 	 of 	 eight 	 (octet) 	
electrons, ns
2
np
6
. 
•	 The 	 negative 	 and 	 positive 	 ions 	 are 	 stabilized	
by electrostatic attraction.
For	 example, 	 the 	 formation 	 of	 NaCl 	 from	
sodium	 and	 chlorine,	 according 	 to 	 the	 above 	
scheme,	can	 be 	 explained	 as:	
Na 																? 							 Na
+
   +    e
–
[Ne] 	3s
1
																			 [Ne] 	
Cl  +  e
–
          ?        Cl
–
[Ne] 	3s
2
	 3p
5
													[Ne] 	3s
2
	3p
6
	or	[Ar] 	
Na
+
  +  Cl
–
     ? 								 NaCl 	or	Na
+
Cl
–
Similarly 	 th e 	 formation 	 of 	 CaF
2
 may be 
shown as:
Ca                ?       Ca
2+
  +  2e
–
[Ar]4s
2
																				[Ar]
F  + e
–
            ?       F
–
[He] 	2s
2
 2p
5
													[He] 	 2s
2
 2p
6
		or	[Ne]
Ca
2+
 + 2F
–
     ?       CaF
2
   or  Ca
2+
(F
– 
)
2
the bond formed, as a result of the 
electrostatic attraction between the 
positive and negative ions was termed as 
Unit 4.indd   101 9/12/2022   9:36:09 AM
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102 chemistry the electrovalent bond. t he electrovalence 
is thus equal to the number of unit charge(s) 
on the ion. 	 Thus, 	 calcium 	 is 	 assigned 	 a	
positive 	 electrovalence 	 of	 two,	 while	 chlorine	
a 	 negative 	electrovalence	of	one. 	
Kösse l ’s 	 post ul at i ons 	 pr ovi de 	 t he 	 basi s 	 f or	
th e 	 modern	 con cepts 	 regard ing 	 ion-formation 	
by electron transfer and the formation of ionic 
crystallin e 	 compou nd s. 	 His 	 views 	 h ave 	 proved	
to 	 be 	 of 	 great	 value 	 in 	 the 	 understanding	 and	
systematisation of the ionic compounds. At 
the same time he did recognise the fact that 
a 	 large 	 number 	 of	 compounds 	 did 	 not	 fit 	 into	
these concepts.
4.1.1 Octet Rule
Kössel 	 and 	 Lewis 	 in 	 1916 	 developed 	 an	
important theory of chemical combination 
between atoms known as electronic theory 
of chemical bonding. According to this, 
atoms  can combine either by transfer of 
valence 	 electrons	 from 	 one	 atom 	 to 	 another	
(gaining 	 or	 losing)	 or	 by 	 sharing 	 of	 valence	
electrons 	 in	 order	 to	 have 	 an 	 octet	 in	 their	
valence 	shells.	This 	is 	known 	 as 	octet rule.
4.1.2 Covalent Bond
Langmuir (1919) 	 refined 	 the 	 Lewis	
postulations by abandoning the idea of 
the stationary cubical arrangement of the 
octet, and by introducing the term covalent 
bond . 	 The 	 Lewis-Langmuir 	 theory 	 can 	 be	
understood by considering the formation of 
the chlorine molecule, Cl
2
. 	 The	 Cl 	 atom	 with	
electronic 	 configuration, 	 [Ne]3s
2
	 3p
5
, is one 
electron 	 short 	 of 	 the 	 argon 	 configuration. 	
The 	 formation 	 of 	 the 	 Cl 	
2
 molecule can be 
understood in terms of the sharing of a pair 
of electrons between the two chlorine atoms, 
each chlorine atom contributing one electron 
to the shared pair. In the process both 
chlorine atoms attain the outer shell octet of 
the	 nearest 	 noble	gas	 (i.e.,	 argon).	
t he dots represent electrons. Such 
structures are referred to as Lewis dot 
structures.
The 	 Le w i s 	 dot 	 s t r uc t ur es 	 c an 	 be 	 w r i t t en 	 f or	
other molecules also, in which the combining 
atoms 	 may 	 be 	 identical 	 or 	 d i f f e r e n t . 	 T h e	
important conditions being that:
•	 Ea c h 	 bo nd 	 i s 	 f o r m e d 	 a s 	 a 	 r e s ul t 	 o f 	 s ha r i ng	
of an electron pair between the atoms.
•	 Each 	 combining 	 atom 	 contri butes 	 at 	 least	
one electron to the shared pair.
•	 The 	 combining 	 atoms 	 attain 	 the 	 outer-
shell	 noble	 gas	 configurations 	 as 	 a	 result	
of the sharing of electrons.
•	 Thus 	 in	 water	 and	 carbon 	 tetrachloride	
mol ecules, 	 formation 	 of 	 covalent 	 bonds	
can be represented as:
   or Cl – Cl
Covalent bond between two Cl atoms
t hus, when two atoms share one 
electron pair they are said to be joined by 
a single covalent bond. In many compounds 
we	 have	multiple bonds 	 between	 atoms.	 The	
formation 	 of 	 mu ltip le 	 b on ds 	 envisages 	 sh a rin g	
of more than one electron pair between two 
atoms. if two atoms share two pairs of 
electrons, the covalent bond between them 
is called a double bond. For 	 example,	 in	 the	
carbon 	 dioxide	 molecule,	 we 	 have 	 two	 double	
bonds 	 between 	 the 	 carbon 	 and 	 oxygen 	 atoms.	
Similarly 	 in 	 ethene 	 molecule	 the	 two	 carbon	
atoms are joined by a double bond.
Double bonds in CO
2
 molecule
Unit 4.indd   102 9/12/2022   9:36:09 AM
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103
Chemi Cal Bonding a nd m ole Cular Stru Cture when combining atoms share three 
electron pairs as in the case of two nitrogen 
atoms in the n
2
 molecule and the two 
carbon atoms in the ethyne molecule, a 
triple bond is formed.
N
2
 molecule
C
2
H
2
 molecule
4.1.3 Lewis Representation of Simple 
Molecules (the Lewis Structures)
t he Lewis dot structures provide a picture 
of bonding in molecules and ions in terms of 
the shared pairs of electrons and the octet 
rule. 	 While 	 such	 a	 picture 	 may 	 not	 explain 	
the 	 bonding 	 and 	 behaviour 	 of 	 a 	 molecule	
completely, it does help in understanding the 
formation and properties of a molecule to a 
large 	 extent. 	 Writing 	 of 	 Lewis 	 dot 	 structures 	 of	
molecules 	 is, 	 therefore, 	 very 	 useful. 	 The 	 Lewis	
dot structures can be written by adopting the 
following steps:
•	 The 	 total 	 number 	 of 	 electrons 	 required	
for writing the structures are obtained 
by 	 adding 	 the 	 valence 	 electrons 	 of 	 the	
combining 	 atoms. 	 For 	 example, 	 in 	 the 	 CH
4
 
molecule 	 there 	 are 	 eight 	 valence 	 electrons 	
available 	 for 	 bonding	 (4	 from	 carbon	 and	
4 	 from	the 	 four	hydrogen	atoms).
•	 For	 anions, 	 each	 negative 	 charge	 would	
mean addition of one electron. For cations, 
each 	 positive 	 ch arge 	 wou ld 	 resu lt 	 in	
subtraction of one electron from the total 
number 	 of 	 valence 	 electrons. 	 For 	 example,	
for the CO
3
2–
	 ion,	 the	 two	 negative 	 charges	
indicate that there are two additional 
electrons 	 than 	 those 	 provided 	 by 	 the	
neutral 	 atoms.	 For 	 NH
4
+
	 ion, 	 one 	 positive	
charge indicates the loss of one electron 
from the group of neutral atoms.
•	 Knowing 	 the 	 chemical 	 symbols 	 of 	 the	
combining 	 atoms 	 and 	 having 	 knowledge 	
of the skeletal structure of the compound 
(known 	 or	 guessed 	 intelligently), 	 it	 is 	 easy 	
to distribute the total number of electrons 
as bonding shared pairs between the 
atoms in proportion to the total bonds.
•	 In	 general 	 the	 least	 electronegative	 atom 	
occupies the central position in the 
molecule/ion. 	 For 	 example 	 in 	 the 	 NF
3
 and 
CO
3
2–
, nitrogen and carbon are the central 
atoms 	 whereas 	 fluorine 	 and 	 oxygen	
occupy the terminal positions.
•	 After accounting for the shared pairs of 
electrons for single bonds, the remaining 
electron 	 pairs 	 are 	 either 	 utilized 	 for	
multiple bonding or remain as the lone 
pai r s . 	 The 	 basi c 	 r e qui r e m e nt 	 bei ng	
that each bonded atom gets an octet of 
electrons.
	 Lewis 	 representations 	 of 	 a 	 few 	 molecules/	
ions	 are 	 given 	 in	 Table 	 4.1.
table 4.1 the Lewis Representation of 
Some Molecules
* Each H atom attains the configuration of helium 
(a duplet of electrons)
C
2
H
4
 molecule
Unit 4.indd   103 9/12/2022   9:36:10 AM
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Page 5


Scientists are constantly discovering new compounds, 
orderly arranging the facts about them, trying to explain 
with the existing knowledge, organising to modify the 
earlier views or evolve theories for explaining the newly 
observed facts.
Unit 4
After studying this Unit, you will be 
able to 
•	 unde r s t a nd 	 Kös s e l - Le w i s	
approach to chemical bonding; 
•	 e x pl a i n 	 t he 	 oc t e t 	 r ul e 	 and 	 i t s	
limitations, 	 draw 	 Lewis 	 structures	
of simple molecules;
•	 explain 	 the 	 formation 	 of	 different	
types of bonds;
•	 describe 	 the	 VSEPR	 theory	 and	
predict the geometry of simple 
molecules;
•	 explain 	 the 	 valence 	 bond	
approach for the formation of 
covalent 	bonds;
•	 predict 	 the	 directional 	 properties	
of	covalent	bonds;
•	 explain 	 the 	 different 	 types 	 of	 	
hybridisation 	 	 involving 	s, p and 
d orbitals and draw shapes of 
simple 	covalent	molecules;
•	 describe 	 the 	 molecular 	 orbital	
theory of homonuclear diatomic 
molecules;
•	 explain 	 the	 concept 	 of 	 hydrogen	
bond.
CHEMiCAL BOnDinG AnD 
MOLECULAR St RUCt URE
Matter is made up of one or different type of elements. 
Under	 normal 	 conditions	 no 	 other	 element	 exists	 as	 an	
independent	 atom	 in 	 nature,	 except 	 noble	 gases.	 However,	
a	 group	 of	atoms 	 is	 found 	 to 	 exist 	 together	 as	 one	 species	
having	 characteristic	 properties. 	 Such 	 a 	 group 	 of	 atoms 	
is	 called	 a	 molecule. 	 Obviously	 there 	 must	 be 	 some	 force	
which holds these constituent atoms together in the 
molecules. the attractive force which holds various 
constituents (atoms, ions, etc.) together in different 
chemical species is called a chemical bond.	 Since 	 the	
formation of chemical compounds takes place as a result of 
combi nat i on 	 of 	 at om s 	 of 	 var i ous 	 el em ent s 	 i n 	 di f f er ent 	 w ays,	
it raises many questions. Why do atoms combine? Why are 
only certain combinations possible? Why do some atoms 
combine while certain others do not? Why do molecules 
possess 	 definite 	 shapes? 	 To 	 answer 	 such 	 questions 	 different	
theories	 and	 concepts 	 have	 been 	 put	 forward 	 from 	 time 	
to	 time. 	 These	 are 	 Kössel-Lewis 	 approach,	 Valence	 Shell 	
Electron 	 Pair 	 R epulsion 	 (VSEPR ) 	 Theory, 	 Valence 	 Bond 	 (V B)	
Theory	 and	 Molecular 	 Orbital 	 (MO) 	 Theory.	 The	 evolution	
of 	 various 	 theories 	 of 	 valence 	 and 	 the 	 interpretation 	 of	
the	 nature	 of	 chemical 	 bonds 	 have 	 closely	 been	 related	 to 	
the	 developments 	 in	 the 	 understanding	 of 	 the	 structure 	
of	 atom,	 the	 electronic	 configuration 	 of 	 elements	 and	 the	
periodic	 table. 	 Every 	 system	 tends	 to 	 be	 more 	 stable	 and 	
bonding is nature’s way of lowering the energy of the system 
to attain stability.
Unit 4.indd   100 9/12/2022   9:36:09 AM
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101
Chemi Cal Bonding a nd m ole Cular Stru Cture 4.1 KÖSSEL-LEwiS AppROACH t O 
CHEMiCAL BOnDinG  
In	 order 	 to	 explain 	 the 	 formation 	 of	 chemical	
bond in terms of electrons, a number of 
attempts were made, but it was only in 
1916 	 when 	 Kössel 	 and 	 Lewis 	 succeeded	
independently 	 in 	 giving 	 a 	 satisfactory	
explan ation .	 Th ey 	 were 	 th e 	 first 	 to 	 provide 	
some 	 logical 	 explanation 	 of 	 valence 	 which 	 was	
based on the inertness of noble gases. 
Lewis 	 pictured 	 the 	 atom 	 in 	 terms 	 of 	 a	
positively 	 charged 	 ‘Kerne l’ 	 (the 	 nucleus	 plus	
the	 inner 	 electrons) 	 and	 the	 outer	 shell	 that	
could 	 accommodate 	 a 	 maximum 	 of 	 eight	
e l e c t r ons . 	 H e , 	 f ur t he r 	 as s um e d 	 t ha t 	 t he s e	
eight electrons occupy the corners of a cube 
which 	 surround	 the	 ‘Kernel’.	 Thus	 the	 single	
outer shell electron of sodium would occupy 
one corner of the cube, while in the case of 
a noble gas all the eight corners would be 
occupied. 	 This	 octet	 of 	 electrons,	 represents	
a particularly stable electronic arrangement. 
Lewis postulated that atoms achieve 
the stable octet when they are linked by 
chemical bonds. In the case of sodium and 
chlorine, this can happen by the transfer of 
an electron from sodium to chlorine thereby 
giving 	 the 	 Na
+
 and Cl
–
 ions. In the case of 
other molecules like Cl
2
,	 H
2
, F
2
, etc., the bond 
is formed by the sharing of a pair of electrons 
between the atoms. In the process each atom 
attains a stable outer octet of electrons.
Lewis Symbols: In the formation of a 
molecule, only the outer shell electrons take 
part in chemical combination and they are 
known as valence electrons .	 The	 inner	 shell	
electrons are well protected and are generally 
not 	 i nvol ved 	 i n 	 t he 	 combi nat i on 	 pr ocess. 
G.N.	 Lewis, 	 an	 American	 chemist 	 introduced	
simple 	 notations 	 to 	 represent 	 valence 	 electrons	
in 	 an 	 atom.	 These	 notations 	 are 	 called 	Lewis 
symbols . 	 For	 example,	 the	 Lewis 	 symbols	 for	
the elements of second period are as under:
Significance of Lewis Symbols : 	 The	
number of dots around the symbol represents 
the 	 number 	 of 	 valence 	 electrons. 	 This 	 number	
of 	 valence 	 electrons 	 helps 	 to 	 calculate 	 the	
common or group valence of the element. 
The 	 gr o up 	 val e nce 	 of 	 t he 	 el e m e nt s 	 i s 	 ge ne r al l y	
either	 equal	 to 	 the 	 number	 of 	 dots	 in	 Lewis	
symbols or 8 minus the number of dots or 
valence 	 electrons.
Kössel, in relation to chemical bonding, 
drew attention to the following facts:
•	 In the periodic table, the highly 
electronegative 	 halogens	 and	 the	 highly	
electropositive 	 alkali 	 metals 	 are 	 separated	
by the noble gases;
•	 The 	 formation 	 of 	 a 	 negative 	 ion 	 from 	 a	
halogen 	 atom 	 and 	 a 	 positive 	 ion 	 from	
an alkali metal atom is associated with 
the gain and loss of an electron by the 
respective 	atoms;
•	 The 	 negative 	 and 	 positive 	 ions 	 thus	
formed attain stable noble gas electronic 
configurations. 	 The	 noble	 gases 	 (with 	 the	
exception 	 of	 helium 	 which 	 has 	 a 	 duplet 	
of 	 electrons) 	 have 	 a 	 particul arl y 	 stabl e	
outer 	 shell	 configuration 	 of 	 eight 	 (octet) 	
electrons, ns
2
np
6
. 
•	 The 	 negative 	 and 	 positive 	 ions 	 are 	 stabilized	
by electrostatic attraction.
For	 example, 	 the 	 formation 	 of	 NaCl 	 from	
sodium	 and	 chlorine,	 according 	 to 	 the	 above 	
scheme,	can	 be 	 explained	 as:	
Na 																? 							 Na
+
   +    e
–
[Ne] 	3s
1
																			 [Ne] 	
Cl  +  e
–
          ?        Cl
–
[Ne] 	3s
2
	 3p
5
													[Ne] 	3s
2
	3p
6
	or	[Ar] 	
Na
+
  +  Cl
–
     ? 								 NaCl 	or	Na
+
Cl
–
Similarly 	 th e 	 formation 	 of 	 CaF
2
 may be 
shown as:
Ca                ?       Ca
2+
  +  2e
–
[Ar]4s
2
																				[Ar]
F  + e
–
            ?       F
–
[He] 	2s
2
 2p
5
													[He] 	 2s
2
 2p
6
		or	[Ne]
Ca
2+
 + 2F
–
     ?       CaF
2
   or  Ca
2+
(F
– 
)
2
the bond formed, as a result of the 
electrostatic attraction between the 
positive and negative ions was termed as 
Unit 4.indd   101 9/12/2022   9:36:09 AM
2024-25
102 chemistry the electrovalent bond. t he electrovalence 
is thus equal to the number of unit charge(s) 
on the ion. 	 Thus, 	 calcium 	 is 	 assigned 	 a	
positive 	 electrovalence 	 of	 two,	 while	 chlorine	
a 	 negative 	electrovalence	of	one. 	
Kösse l ’s 	 post ul at i ons 	 pr ovi de 	 t he 	 basi s 	 f or	
th e 	 modern	 con cepts 	 regard ing 	 ion-formation 	
by electron transfer and the formation of ionic 
crystallin e 	 compou nd s. 	 His 	 views 	 h ave 	 proved	
to 	 be 	 of 	 great	 value 	 in 	 the 	 understanding	 and	
systematisation of the ionic compounds. At 
the same time he did recognise the fact that 
a 	 large 	 number 	 of	 compounds 	 did 	 not	 fit 	 into	
these concepts.
4.1.1 Octet Rule
Kössel 	 and 	 Lewis 	 in 	 1916 	 developed 	 an	
important theory of chemical combination 
between atoms known as electronic theory 
of chemical bonding. According to this, 
atoms  can combine either by transfer of 
valence 	 electrons	 from 	 one	 atom 	 to 	 another	
(gaining 	 or	 losing)	 or	 by 	 sharing 	 of	 valence	
electrons 	 in	 order	 to	 have 	 an 	 octet	 in	 their	
valence 	shells.	This 	is 	known 	 as 	octet rule.
4.1.2 Covalent Bond
Langmuir (1919) 	 refined 	 the 	 Lewis	
postulations by abandoning the idea of 
the stationary cubical arrangement of the 
octet, and by introducing the term covalent 
bond . 	 The 	 Lewis-Langmuir 	 theory 	 can 	 be	
understood by considering the formation of 
the chlorine molecule, Cl
2
. 	 The	 Cl 	 atom	 with	
electronic 	 configuration, 	 [Ne]3s
2
	 3p
5
, is one 
electron 	 short 	 of 	 the 	 argon 	 configuration. 	
The 	 formation 	 of 	 the 	 Cl 	
2
 molecule can be 
understood in terms of the sharing of a pair 
of electrons between the two chlorine atoms, 
each chlorine atom contributing one electron 
to the shared pair. In the process both 
chlorine atoms attain the outer shell octet of 
the	 nearest 	 noble	gas	 (i.e.,	 argon).	
t he dots represent electrons. Such 
structures are referred to as Lewis dot 
structures.
The 	 Le w i s 	 dot 	 s t r uc t ur es 	 c an 	 be 	 w r i t t en 	 f or	
other molecules also, in which the combining 
atoms 	 may 	 be 	 identical 	 or 	 d i f f e r e n t . 	 T h e	
important conditions being that:
•	 Ea c h 	 bo nd 	 i s 	 f o r m e d 	 a s 	 a 	 r e s ul t 	 o f 	 s ha r i ng	
of an electron pair between the atoms.
•	 Each 	 combining 	 atom 	 contri butes 	 at 	 least	
one electron to the shared pair.
•	 The 	 combining 	 atoms 	 attain 	 the 	 outer-
shell	 noble	 gas	 configurations 	 as 	 a	 result	
of the sharing of electrons.
•	 Thus 	 in	 water	 and	 carbon 	 tetrachloride	
mol ecules, 	 formation 	 of 	 covalent 	 bonds	
can be represented as:
   or Cl – Cl
Covalent bond between two Cl atoms
t hus, when two atoms share one 
electron pair they are said to be joined by 
a single covalent bond. In many compounds 
we	 have	multiple bonds 	 between	 atoms.	 The	
formation 	 of 	 mu ltip le 	 b on ds 	 envisages 	 sh a rin g	
of more than one electron pair between two 
atoms. if two atoms share two pairs of 
electrons, the covalent bond between them 
is called a double bond. For 	 example,	 in	 the	
carbon 	 dioxide	 molecule,	 we 	 have 	 two	 double	
bonds 	 between 	 the 	 carbon 	 and 	 oxygen 	 atoms.	
Similarly 	 in 	 ethene 	 molecule	 the	 two	 carbon	
atoms are joined by a double bond.
Double bonds in CO
2
 molecule
Unit 4.indd   102 9/12/2022   9:36:09 AM
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103
Chemi Cal Bonding a nd m ole Cular Stru Cture when combining atoms share three 
electron pairs as in the case of two nitrogen 
atoms in the n
2
 molecule and the two 
carbon atoms in the ethyne molecule, a 
triple bond is formed.
N
2
 molecule
C
2
H
2
 molecule
4.1.3 Lewis Representation of Simple 
Molecules (the Lewis Structures)
t he Lewis dot structures provide a picture 
of bonding in molecules and ions in terms of 
the shared pairs of electrons and the octet 
rule. 	 While 	 such	 a	 picture 	 may 	 not	 explain 	
the 	 bonding 	 and 	 behaviour 	 of 	 a 	 molecule	
completely, it does help in understanding the 
formation and properties of a molecule to a 
large 	 extent. 	 Writing 	 of 	 Lewis 	 dot 	 structures 	 of	
molecules 	 is, 	 therefore, 	 very 	 useful. 	 The 	 Lewis	
dot structures can be written by adopting the 
following steps:
•	 The 	 total 	 number 	 of 	 electrons 	 required	
for writing the structures are obtained 
by 	 adding 	 the 	 valence 	 electrons 	 of 	 the	
combining 	 atoms. 	 For 	 example, 	 in 	 the 	 CH
4
 
molecule 	 there 	 are 	 eight 	 valence 	 electrons 	
available 	 for 	 bonding	 (4	 from	 carbon	 and	
4 	 from	the 	 four	hydrogen	atoms).
•	 For	 anions, 	 each	 negative 	 charge	 would	
mean addition of one electron. For cations, 
each 	 positive 	 ch arge 	 wou ld 	 resu lt 	 in	
subtraction of one electron from the total 
number 	 of 	 valence 	 electrons. 	 For 	 example,	
for the CO
3
2–
	 ion,	 the	 two	 negative 	 charges	
indicate that there are two additional 
electrons 	 than 	 those 	 provided 	 by 	 the	
neutral 	 atoms.	 For 	 NH
4
+
	 ion, 	 one 	 positive	
charge indicates the loss of one electron 
from the group of neutral atoms.
•	 Knowing 	 the 	 chemical 	 symbols 	 of 	 the	
combining 	 atoms 	 and 	 having 	 knowledge 	
of the skeletal structure of the compound 
(known 	 or	 guessed 	 intelligently), 	 it	 is 	 easy 	
to distribute the total number of electrons 
as bonding shared pairs between the 
atoms in proportion to the total bonds.
•	 In	 general 	 the	 least	 electronegative	 atom 	
occupies the central position in the 
molecule/ion. 	 For 	 example 	 in 	 the 	 NF
3
 and 
CO
3
2–
, nitrogen and carbon are the central 
atoms 	 whereas 	 fluorine 	 and 	 oxygen	
occupy the terminal positions.
•	 After accounting for the shared pairs of 
electrons for single bonds, the remaining 
electron 	 pairs 	 are 	 either 	 utilized 	 for	
multiple bonding or remain as the lone 
pai r s . 	 The 	 basi c 	 r e qui r e m e nt 	 bei ng	
that each bonded atom gets an octet of 
electrons.
	 Lewis 	 representations 	 of 	 a 	 few 	 molecules/	
ions	 are 	 given 	 in	 Table 	 4.1.
table 4.1 the Lewis Representation of 
Some Molecules
* Each H atom attains the configuration of helium 
(a duplet of electrons)
C
2
H
4
 molecule
Unit 4.indd   103 9/12/2022   9:36:10 AM
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104 chemistry problem 4.1
W r i t e 	 t he 	 Le w i s 	 dot 	 st r uc t ur e 	 of 	 C O	
molecule.
Solution
Step 1. 	 Count 	 the 	 total 	 number 	 of 	 valence	
electrons 	 of	 carbon	 and 	 oxygen	 atoms.	
The 	 outer	 (valence)	 shell 	 configurations 	
of 	 carbon 	 and 	 oxygen 	 atoms 	 are: 	 2s
2
 2p
2
 
and 2s
2
 2p
4
,	 respectively. 	 The 	 valence 	
electrons 	available	are	4 	+	6 	=10.
Step 2. 	 The 	 skeletal	 structure 	 of	 CO	 is	
written as: C  O
Step 3. 	 Draw 	 a 	 single 	 bond 	 (one 	 shared 	
electron 	 pair) 	 betw een 	 C 	 and 	 O 	 and	
complete the octet on O, the remaining 
two electrons are the lone pair on C.
 Th is 	 does 	 n ot 	 complete 	 th e 	 octet 	 on 	
carbon 	 and	 hence	 we	 have 	 to	 resort 	 to 	
multiple 	 bonding 	 (in 	 this 	 case 	 a	 triple 	
bo nd) 	 bet w ee n 	 C 	 and 	 O 	 at om s . 	 Thi s	
sati sfies 	 t he 	 oct et 	 rul e 	 condi t i on 	 f or 	 bot h	
atoms.
problem 4.2
Write 	 the	 Lewis 	 structure 	 of	 the	 nitrite 	
ion, 	NO
2
– 
.
Solution
Step 1. Count the total number of 
valence 	 electrons	 of 	 the 	 nitrogen 	 atom,	
t he 	 ox y g e n 	 at o m s 	 and 	 t he 	 add i t i o na l 	 one	
negative 	charge 	 (equal 	to	one	 electron).
	 N(2s
2
 2p
3
),	O	(2s
2
 2p
4
)
	 5 	+	(2	 ×	6)	+1	 =	 18	electrons 	
Step 2.	 The 	 skeletal	 structure 	 of	 NO
2
–
 is 
written 	as 	 : 		O 			N 			O
Step 3. 	 Draw	 a	 single 	 bond 	 (one 	 shared	
electron 	 pair) 	 between	 the 	 nitrogen	 and	
e a c h 	 o f 	 t he 	 ox y g e n 	 at o m s 	 co m pl e t i ng 	 t he	
octets 	 on 	 oxygen	 atoms.	 This,	 however,	
does not complete the octet on nitrogen 
if the remaining two electrons constitute 
lone pair on it.
	 H ence 	 w e 	 have 	 t o 	 resort 	 t o 	 mul t i pl e	
bonding between nitrogen and one of 
the 	 oxygen	 atoms	 (in 	 this 	 case	 a 	 double 	
bond). 	 This	 leads 	 to	 the	 following 	 Lewis	
dot structures.
4.1.4 Formal Charge
Lewis 	 dot 	 structures, 	 in 	 general, 	 do 	 not	
represent the actual shapes of the molecules. 
In case of polyatomic ions, the net charge is 
possessed by the ion as a whole and not by 
a	 particular	 atom. 	 It 	 is,	 however,	 feasible	 to	
assign 	 a 	 formal 	 charge 	 on 	 each 	 atom. 	 The	
formal charge of an atom in a polyatomic 
molecule or ion may be defined as the 
di f f er ence 	 betw een 	 the 	 number 	 of 	 val ence	
electrons of that atom in an isolated or free 
state and the number of electrons assigned 
to 	 that 	 atom 	 in 	 the 	 Lew is 	 structure. 	 It 	 is	
expressed	 as 	:
Formal 	 charge 	(F.C.)		
on 	 an 	atom	in	 a	Lewis 	
structure
=
total 	number 	 of 	valence 	
electrons in the free 
atom
—
total number of non 
bonding	(lone	pair)	
electrons
—	(1/2)	
total number of  
bonding 	 (shared)	
electrons
Unit 4.indd   104 9/12/2022   9:36:10 AM
2024-25