NCERT Textbook - Chemical Bonding and Molecular Class 11 Notes | EduRev

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Class 11 : NCERT Textbook - Chemical Bonding and Molecular Class 11 Notes | EduRev

 Page 1


96 CHEMISTRY
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
• understand KÖssel-Lewis
approach to chemical bonding;
•
explain the octet rule and its
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 STRUCTURE
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 combination of atoms of various elements in different
ways, 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 Repulsion (VSEPR)
Theory, Valence Bond (VB) 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.
© NCERT
not to be republished
Page 2


96 CHEMISTRY
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
• understand KÖssel-Lewis
approach to chemical bonding;
•
explain the octet rule and its
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 STRUCTURE
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 combination of atoms of various elements in different
ways, 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 Repulsion (VSEPR)
Theory, Valence Bond (VB) 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.
© NCERT
not to be republished
97
CHEMICAL BONDING AND MOLECULAR STRUCTURE
4.1 KÖSSEL-LEWIS APPROACH TO
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
explanation. They were the 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 ‘Kernel’ (the nucleus plus
the inner electrons) and the outer shell that
could accommodate a maximum of eight
electrons. He, further assumed that these
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 involved in the combination process.
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
group valence of the elements is generally
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 particularly stable
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 the 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
© NCERT
not to be republished
Page 3


96 CHEMISTRY
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
• understand KÖssel-Lewis
approach to chemical bonding;
•
explain the octet rule and its
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 STRUCTURE
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 combination of atoms of various elements in different
ways, 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 Repulsion (VSEPR)
Theory, Valence Bond (VB) 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.
© NCERT
not to be republished
97
CHEMICAL BONDING AND MOLECULAR STRUCTURE
4.1 KÖSSEL-LEWIS APPROACH TO
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
explanation. They were the 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 ‘Kernel’ (the nucleus plus
the inner electrons) and the outer shell that
could accommodate a maximum of eight
electrons. He, further assumed that these
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 involved in the combination process.
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
group valence of the elements is generally
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 particularly stable
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 the 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
© NCERT
not to be republished
98 CHEMISTRY
the electrovalent bond. The 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.
Kössel’s postulations provide the basis for
the modern concepts regarding ion-formation
by electron transfer and the formation of ionic
crystalline compounds. His views have 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).
The dots represent electrons. Such
structures are referred to as Lewis dot
structures.
The Lewis dot structures can be written
for other molecules also, in which the
combining atoms may be identical or
different. The important conditions being that:
• Each bond is formed as a result of sharing
of an electron pair between the atoms.
• Each combining atom contributes 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
molecules, formation of covalent bonds
can be represented as:
                                             or Cl – Cl
        Covalent bond between two Cl atoms
Thus, 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 multiple bonds envisages
sharing 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
© NCERT
not to be republished
Page 4


96 CHEMISTRY
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
• understand KÖssel-Lewis
approach to chemical bonding;
•
explain the octet rule and its
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 STRUCTURE
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 combination of atoms of various elements in different
ways, 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 Repulsion (VSEPR)
Theory, Valence Bond (VB) 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.
© NCERT
not to be republished
97
CHEMICAL BONDING AND MOLECULAR STRUCTURE
4.1 KÖSSEL-LEWIS APPROACH TO
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
explanation. They were the 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 ‘Kernel’ (the nucleus plus
the inner electrons) and the outer shell that
could accommodate a maximum of eight
electrons. He, further assumed that these
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 involved in the combination process.
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
group valence of the elements is generally
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 particularly stable
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 the 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
© NCERT
not to be republished
98 CHEMISTRY
the electrovalent bond. The 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.
Kössel’s postulations provide the basis for
the modern concepts regarding ion-formation
by electron transfer and the formation of ionic
crystalline compounds. His views have 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).
The dots represent electrons. Such
structures are referred to as Lewis dot
structures.
The Lewis dot structures can be written
for other molecules also, in which the
combining atoms may be identical or
different. The important conditions being that:
• Each bond is formed as a result of sharing
of an electron pair between the atoms.
• Each combining atom contributes 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
molecules, formation of covalent bonds
can be represented as:
                                             or Cl – Cl
        Covalent bond between two Cl atoms
Thus, 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 multiple bonds envisages
sharing 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
© NCERT
not to be republished
99
CHEMICAL BONDING AND MOLECULAR STRUCTURE
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.
4.1.3 Lewis Representation of Simple
Molecules (the Lewis Structures)
The 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 charge would result
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
pairs. The basic requirement being 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
N
2
 molecule
C
2
H
2
 molecule
© NCERT
not to be republished
Page 5


96 CHEMISTRY
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
• understand KÖssel-Lewis
approach to chemical bonding;
•
explain the octet rule and its
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 STRUCTURE
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 combination of atoms of various elements in different
ways, 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 Repulsion (VSEPR)
Theory, Valence Bond (VB) 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.
© NCERT
not to be republished
97
CHEMICAL BONDING AND MOLECULAR STRUCTURE
4.1 KÖSSEL-LEWIS APPROACH TO
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
explanation. They were the 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 ‘Kernel’ (the nucleus plus
the inner electrons) and the outer shell that
could accommodate a maximum of eight
electrons. He, further assumed that these
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 involved in the combination process.
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
group valence of the elements is generally
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 particularly stable
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 the 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
© NCERT
not to be republished
98 CHEMISTRY
the electrovalent bond. The 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.
Kössel’s postulations provide the basis for
the modern concepts regarding ion-formation
by electron transfer and the formation of ionic
crystalline compounds. His views have 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).
The dots represent electrons. Such
structures are referred to as Lewis dot
structures.
The Lewis dot structures can be written
for other molecules also, in which the
combining atoms may be identical or
different. The important conditions being that:
• Each bond is formed as a result of sharing
of an electron pair between the atoms.
• Each combining atom contributes 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
molecules, formation of covalent bonds
can be represented as:
                                             or Cl – Cl
        Covalent bond between two Cl atoms
Thus, 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 multiple bonds envisages
sharing 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
© NCERT
not to be republished
99
CHEMICAL BONDING AND MOLECULAR STRUCTURE
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.
4.1.3 Lewis Representation of Simple
Molecules (the Lewis Structures)
The 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 charge would result
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
pairs. The basic requirement being 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
N
2
 molecule
C
2
H
2
 molecule
© NCERT
not to be republished
100 CHEMISTRY
Problem 4.1
Write the Lewis dot structure of CO
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) between C and O and
complete the octet on O, the remaining
two electrons are the lone pair on C.
This does not complete the octet on
carbon and hence we have to resort to
multiple bonding (in this case a triple
bond) between C and O atoms. This
satisfies the octet rule condition for both
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,
the oxygen atoms and the additional 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
each of the oxygen atoms completing the
octets on oxygen atoms. This, however,
does not complete the octet on nitrogen
if the remaining two electrons constitute
lone pair on it.
Hence we have to resort to multiple
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
difference between the number of valence
electrons of that atom in an isolated or free
state and the number of electrons assigned
to that atom in the Lewis 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
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