1.  NOTES BY
SAMPATHU ARUNKUMAR
WALAJAPET
BMV SR.SEC SCHOOL
CLASS:XI

2.  The attractive force which holds various
constituents(atoms, ions,etc) together in
different chemical species is called a chemical
bond.
 Why do atoms combine?
 Why do molecules posses definite shape?
 Why do some atoms combine , certain others do
not?

3.  Valence electrons, ionic bond, covalent bond,
bond parameters, Lewis structure, polar
character of covalent bond, covalent character
of ionic bond, valence bond theory, resonance.
 Geometry of covalent molecules, VSEPR theory,
concept of hybridization, involving s, p and d
orbitals .
 Shapes of some simple molecules, molecular
orbital theory of homonuclear diatomic
Molecules(qualitative idea only), Hydrogen
bond.

4.  Atoms combine together in order to complete their
respective octets so as to acquire the stable inert gas
configuration. This can occur in two ways:
 1)By complete transference of one or more electrons
from one atom to another. Chemical bond formed is
called as electrovalent Bond or ionic bond.
 2)By sharing of electrons
 This can occur in two ways:
 a)When the shared electrons are contributed by the
two combining atoms equally ,the bond formed is
called covalent bond.
 b)When these electrons are contributed entirely by
one of the atoms but shared by both ,the bond form
is known as co-ordinate bond or dative bond.

5.  Ex:NaCl

6.  Ex:H2, F2,Cl2

7.  EX:NH4

8.  The outer shell electrons are known as valence
electrons. The outer shell electrons are shown
as dots surrounding the symbol of the atom.
 These symbols are known as Lewis symbols are
electron dot symbols.

9.  The number of dots around the symbol give
the number of electrons present in the
outermost shell. This number of electrons
helps to calculate the common valency of the
element.
 The common valency of the element is either
equal to the number of dots in the Lewis
structure or 8 minus the number of dots.
 Li, Be, B and C have valency 1 ,2, 3 and 4 i.e.
equal to the number of dots.
 Valencies of N, O, F and Ne are 3 ,2, 1 and 0
i.e. 8 minus the number of dots.

11. Kossel in relation to chemical bonding:
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 are stabilized by
electrostatic attraction.

13.  The bond formed, as a result of the
electrostatic attraction between the positive
and negative ions was termed as the
electrovalent bond.
 Octet Rule:
Kossel 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 or by sharing of valence electrons in
order to have an octet in their valence shells.

14. Covalent bond:
Langmuir refined the lewis postulation by
introducing the term covalent bond.
Ex: Formation of the chlorine molecule.
The dots represent electrons, such structures
are referred to as lewis dot structures.

15.  Conditions for Lewis dot structure:
1. Each bond is formed as a result of sharing of
an electron pair between the atoms.
2. Each combining atom contributes at least
one electron to the shared pair.
3. The combining atoms attain the outer-shell
noble gas configurations as a result of the
sharing of electrons.
4. Thus in water and carbon tetrachloride
molecules, formation of covalent bond can
be represented as

16.  When two atoms share one electron pair they
are said to be joined by a single covalent
bond.
 If two atoms share two pairs of electrons, the
covalent bond between them is called a
double bond.
 When combining atoms share three electron a
triple bond is formed.

18.  Lewis representation of simple molecules and ions
 Step 1 Calculate the total number of valence
electrons of the atoms present.
 Step 2 If the species is an anion ,add number of
electrons equal to the units of negative charge and if
the species is a cation ,subtract number of electrons
equal to the units of positive charge.
 Step 3 Select the central atom and draw the skeleton
structure by guess to indicate which atom is linked
to which other atom. Hydrogen and fluorine usually
occupy terminal position.
 Step 4 Put one shared pair of electrons between
every two atoms to represent a single bond between
them. Use the remaining pairs of electrons either for
multiple bonding or to show them as lone pair,
keeping in mind that octet of each atom is
completed.

20. Lewis dot structure of CO:
1. Count the total number of valence electrons
C-2s22p2 O-2s22p4 =>4+6=10
2. The skeletal structure of CO is written as
C O
3. Complete octet for each atom

21.  Lewis Structure of the nitrite:
Step:1 Count the total number of valence
electrons of nitrogen and oxygen atoms
N(2s22p3), O(2s22p4)
5+(2x6)+1=18 electrons
Step:2 The skeletal structure of NO2
- is written
as : O N O
Step:3 Complete the octet

22.  Formal Charge:
In polyatomic ions, the net charge is the charge on
the ion as a whole and not by particular atom.
However, charges can be assigned to individual
atoms or ions. These are called formal charges.
Formal charge=
Valence electrons- Non bonding electrons-
𝟏
𝟐
[𝒔𝒉𝒂𝒓𝒆𝒅 𝒆𝒍𝒆𝒄𝒕𝒓𝒐𝒏𝒔]

24. Limitation of the octet rule:
(i) The incomplete octet of the central atoms:
In some covalent compounds central atom
has less than eight electrons, i.e., it has an
incomplete octet. For example,

25.  Li, Be and B have 1, 2, and 3 valence electrons
only.
(ii) Odd-electron molecules: There are certain
molecules which have odd number of
electrons the octet rule is not applied for all
the atoms.
The expanded Octet: In many compounds there
are more than eight valence electrons around
the central atom. It is termed as expanded octet.
For Example,

26. Other drawbacks of octet theory:
(i) Some noble gases, also combine with oxygen
and fluorine to form a number of compounds
like XeF2 ,KrF2, XeOF2 etc.
(ii) This theory does not account for the shape
of the molecule.
(iii) It does not give any idea about the energy
of the molecule and relative stability.

27.  Ionic or Electrovalent bond:
From Kossel and Lewis treatment the
formation of ionic bond would depend upon:
1. The formation of the positive and negative
ions from the respective neutral atoms.
2. The arrangement of the positive and negative
ions in the solid .
The formation of a positive ion involves
ionization i.e removal of electrons from the
neutral atom.
The negative ion involves the addition of electron
to the neutral atom.

28.  M(g)M+(g)+e- Ionization enthalpy
 X(g)+e-X-(g) Electron gain enthalpy
 M+ (g) + X-(g)MX(s)
Ionic bond will be formed more easily between
elements with comparatively low ionization
enthalpies and elements with comparatively
high negative value of electron gain enthalpy.
Most ionic compounds have cations derived
from metallic elements and anion from non-
metallic elements.

29.  The three dimensional arrangement of cations
and anions in a ionic compounds is called
lattice.
 Crystal structure depends on size of ions and
packing arrangements etc.

30.  In ionic solids the sum of the electron gain
enthalpy and the ionization enthalpy may be
positive.
 The crystal structure get stabilized due to the
energy released in the formation of crystal
lattice.
 Ex: Ionization enthalpy of
 Na-Na+ions=495.8kj/mol
 Electron gain enthalpy of
 Cl(g)-Cl- =-348.7kj/mol
 The sum of the two is 147.1kj/mol

31.  The energy required for the lattice formation
is -788kj/mole.
 The stability of an ionic compound is
provided by is enthalpy of lattice formation
and not simply by achieving octet of electrons
around the ionic species in gaseous state.
 Lattice enthalpy:
The lattice enthalpy of an ionic solid is defined
as the energy required to completely separate
one mole of a solid ionic compound into
gaseous constituent ions.

32.  For example:
The lattice enthalpy of NaCl is 788kj/mol. This
means that 788kj/mol of energy is required to
separate one mole of solid NaCl into one mole
of Na+ (g) and one mole of Cl-(g) to an infinite
distance.

33. BondLength
It is defined as the equilibrium distance
between the centres of the nuclei of the two
bonded atoms. It is expressed in terms of A.
Experimentally, it can be defined by X-ray
diffraction or electron diffraction method.

34.  The vander Waals radius represents the
overall size of the atom which includes its
valence shell in a nonbonded situation.

35.  Covalent and vander Waals radii in a chlorine
molecule.

36. Bond Angle
It is defined as -the angle between the lines
representing the orbitals containing the
bonding – electrons.
It helps us in determining the shape. It can be
expressed in degree. Bond angle can be
experimentally determined by spectroscopic
methods.

37. Bond Enthalpy:
It is defined as the amount of energy required
to break one mole of bonds of a particular type
to separate them into gaseous atoms.
Bond Enthalpy is also known as bond
dissociation enthalpy .
Unit of bond enthalpy = kJ mol-1
The magnitude of bond enthalpy is also related
to bond multiplicity. Greater the bond
multiplicity, more will be the bond enthalpy.

38.  H-H bond enthalpy in hydrogen molecule-
435.8kj/mol
 O=O (g)O(g)+O(g) H=498kj/mol

39.  For poly atomic molecule, the enthalpy
needed to break the two O-H bonds is not
the same.
 The difference in the H value shows that the
second O-H bond undergoes some change
because of changed chemical environment.
 Therefore in polyatomic molecules the term
mean or average bond enthalpy is used.

40.  Bond Order
According to Lewis, in a covalent bond, the
bond order is given by the number of bonds
between two atoms in a molecule. For
example,
Bond order of H2 (H —H) =1
Bond order of 02 (O = O) =2
Bond order of N2 (N N) =3
Isoelectronic molecules and ions have
identical bond orders.
 For example, F2 and O2
2- have bond order =
1.

41.  N2, CO and NO+ have bond order = 3. With
the increase in bond order, bond enthalpy
increases and bond length decreases. For
example,

42.  Resonance Structures
There are many molecules whose behaviour
cannot be explained by a single-Lewis
structure, For example, Lewis structure of
Ozone represented as follows:

43.  Thus, according to the concept of resonance,
whenever a single Lewis structure cannot
explain all the properties of the molecule, the
molecule is then supposed to have many
structures with similar energy. Positions of
nuclei, bonding and nonbonding pairs of
electrons are taken as the canonical structure
of the hybrid which describes the molecule
accurately.

44.  For 03, the two structures shown above are
canonical structures and the III structure
represents the structure of 03 more
accurately. This is also called resonance
hybrid.
Some resonating structures of some more
molecules and ions are shown as follows:

45.  Structure of CO3
2-
Resonance stabilizes the molecule as the
energy of the resonance hybrid is less than the
energy of any single cannonical structure.
Resonance averages the bond characteristics as
a whole.

46.  Polarity of Bonds
Polar and Non-Polar Covalent bonds
Non-Polar Covalent bonds: When the atoms
joined by covalent bond are the same like; H2,
02, Cl2, the shared pair of electrons is equally
attracted by two atoms and thus the shared
electron pair is equidistant to both of them.
Alternatively, we can say that it lies exactly in
the centre of the bonding atoms.
 As a result, no poles are developed and the
bond is called as non-polar covalent bond.
The corresponding molecules are known as
non-polar molecules.
For Example,

47.  Polar bond: When covalent bonds formed
between different atoms of different
electronegativity, shared electron pair between
two atoms gets displaced towards highly
electronegative atoms.
 For Example, in HCl molecule, since
electronegativity of chlorine is high as compared
to hydrogen thus, electron pair is displaced
more towards chlorine atom, thus chlorine will
acquire a partial negative charge (δ–) and
hydrogen atom have a partial positive charge
(δ+) with the magnitude of charge same as on
chlorination. Such covalent bond is called polar
covalent bond.

48. Dipole Moment:
 Due to polarity, polar molecules are also known as
dipole molecules and they possess dipole
moment. Dipole moment is defined as the product
of magnitude of the positive or negative charge
and the distance between the charges.

49. Applications of Dipole moment:
 (i) For determining the polarity of the molecules
 (ii) In finding the shapes of the molecules
 (iii) In calculating the percentage ionic character of
polar bonds.
 (iv)To distinguish between cis and trans isomers

50.  Polyatomic molecules
 As a polyatomic molecule has more than one
polar bond ,the dipole moment is equal to the
resultant dipole moment of all the individual
bonds.
 The magnitude of resultant dipole moment
not only depends upon the values of the
individual dipole moment of the bonds but
also on their arrangement in space.
 Dipole moment of water is 1.84 D which is
equal to the resultant dipole moment of two
O-H bonds.

51.  In carbon dioxide molecule, there are two
polar bonds. These polar bonds possesses
the same value of dipole moment but the
overall dipole moment of molecule is found
to be zero. Individual dipole moment in this
molecule are of equal magnitude but their
directions are opposite to each and hence
cancel out.

52.  In symmetrical molecules like Boron
trifluoride ( BF3) Methane ( CH4) and carbon
tetrachloride (CCl4), the molecular dipole
moment is found to be zero. Individual dipole
moments cancel out on account of the
symmetry of the molecule. In BF3 ,the
resultant of two bond moments being equal
and opposite to that of the third cancel out.


54.  Dipole moment of NH3 and NF3
 Both NH3 and NF3 molecules have pyramidal
shape with one lone pair of electrons on N
atom. As fluorine is highly electronegative, it
appears that N-F bond should be more polar
and the net dipole moment of NF3 should be
much greater than that of NH3.

55.  The dipole formed between the lone pair and
nitrogen atom has to be taken into
consideration which is in the direction of the
lone pair.
F is more electronegative than nitrogen
therefore direction of bond is from nitrogen
to fluorine whereas nitrogen is more
electronegative than hydrogen, the direction
of bond is from hydrogen to nitrogen.

56.  In case of NH3 the orbital dipole due to lone
pair is in the same direction as the resultant
dipole moment of the N-H bonds, Whereas in
NF3 the orbital dipole is in the direction
opposite to the resultant dipole moment of
the three N-F bonds.
 The orbital dipole because of lone pair
decreases the effect of the resultant N-F
bond moments, which results in the low
dipole moment of NF3 .

57.  When cations and anions approach each
other, the valence shell of anions are pulled
towards cation nucleus due to the coulombic
attraction and thus shape of the anion is
deformed.
 This phenomenon of deformation of anion by
a cation is known as polarization and the
ability of cation to polarize a nearby anion is
called as polarizing power of cation.

58.  Fajan points out that greater is the
polarization of anion in a molecule more is
covalent character in it. This is Fajan’s rule.
 1. The smaller the size of the cation and
larger the size of the anion, the greater the
covalent character of an ionic bond.
 2.The greater the charge on the cation, the
greater the covalent character of the ionic
bond.

59.  For cations of the same size and charge, the
one with electronic configuration (n-1)𝒅 𝒏
𝒏𝒔 𝟎
 Typical of transition metals is more polarising
than the one with noble gas configuration,
𝒏𝒔 𝟐 𝒏𝒑 𝟔.
 The cation polarises the anion, pulling the
electronic charge toward itself and thereby
increasing the electronic charge between the
two.

61.  The Valence Shell Electron Pair Repulsion
(VSEPR) Theory
Sidgwick and Powell in 1940,
proposed a simple theory
based on repulsive character
of electron pairs in the
valence shell of the atoms. It
was further developed by
Nyholm and Gillespie (1957).

62.  Main Postulates are the following:
(i) The exact shape of molecule depends upon
the number of electron pairs (bonded or non
bonded) around the central atoms.
(ii) The electron pairs have a tendency to repel
each other since they exist around the central
atom and the electron clouds are negatively
charged.
(iii) Electron pairs try to take such position which
can minimize the repulsion between them.
(iv) The valence shell is taken as a sphere with
the electron pairs placed at maximum distance.
(v) A multiple bond is treated as if it is a single
electron pair and the electron pairs which
constitute the bond as single super pairs.

63. 1)The shape of a molecule containing only two
atoms is always linear.
2)For molecules containing three or more atoms,
one of the atoms is called the central atom to
which other atoms are linked.
3.The order of repulsion between electron pairs as
follow:
 Lone Pair- lone pair > Lone Pair- bond- pair >
Bond Pair- bond pair
4.The exact shape of the molecule depend upon
the total number of electron pairs present around
the central atom.

64.  The prediction of geometrical shapes of
molecules with the help of VSPER theory, it is
convenient to divide molecules into two
categories as
 (i) Molecules in which the central atom has no
lone pair.
 (ii) Molecules in which the central atom has
one or more lone pairs.

66.  Formula to find Electron pair and lone pair.
 Electron pair=
𝟏
𝟐
𝒗𝒆 + 𝑴𝑨
 Ve=valence electron
 MA-Monovalent atom
 Lone pair= Electron pair- Bond pair
 Bond pair= no of bonds.
Check with NH3, H2O

73.  Valence Bond Theory
 This theory was put forward by Heitler and
London in 1927 and was further developed by
Pauling and others.
 In terms of energy
 When the two atoms are far apart from each
other there is no interaction between them.
 When they come closer to each other ,the new
forces come into operation.

74. These forces are of two types:
 1)The forces of repulsion between the nuclei of
these combining atoms and between the
electrons of these atoms. These forces tends
to increase the energy of the system.
 2)The forces of attraction between the nucleus
of one atom and the electrons of other atom.
These forces tend to decrease the energy of
the system.
 If in a system, these new forces can decrease
the energy ,then possibility of chemical
bonding exist and if these forces lead to
increase in energy, the chemical bonding is not
possible.

75.  If we have two hydrogen atoms A and B kept
at infinite distance from each other. There
will be no interaction between them. But as
they begin to come closer, the following new
forces will start operating:

77.  1)Force of attraction between nucleus of A
and electron of nucleus B.
 2)Force of attraction between nucleus of B
and electron of A.
 3)Force of repulsion between electrons of the
atoms.
 4)Force of repulsion between nuclei of the
two atoms.

78.  In hydrogen ,the magnitude of the attractive
forces is more than that of repulsive forces. As a
result, the potential energy of the system
decreases and a molecule of hydrogen is formed .
 As the atoms start coming closer to each other
from infinite distance, they start interacting with
each other and the system starts losing its energy
as the forces of attraction exceeds the forces of
repulsion. But at a certain equilibrium distance
,the forces of repulsion are just balanced by the
force of attraction and the energy of the system
become minimum. The two hydrogen atoms are
said to be bonded together to form a stable
system i.e. a molecule.
 The distance between the two nuclei is
called Bond length.

79.  If we want to break the bond ,i.e. to separate
the atoms, we have to supply the same
amount of energy.
 A stronger bond is that which requires
greater energy for the separation of atoms.
 In the formation of a strong bond, more
energy should be released by the system.
Lesser the amount of energy liberated,
weaker will be bond formed and larger is the
amount of energy liberated, stronger will be
the bond formed.
 The energy required to break one mole of
bonds of the same kind is known as bond
energy or bond dissociation energy.

80.  Orbital overlap concept:
A covalent bond is formed by the partial overlap
of the two half filled atomic orbitals containing
electrons with opposite spins.
Formation of hydrogen molecule:
When two hydrogen atoms having electrons with
opposite spins come close to each other, their s
orbitals overlap with each other resulting in the
union of the two atoms to form a molecules.

81.  The strength of the bond depends upon the
extent of overlapping. Greater the
overlapping ,stronger is the bond formed.
 The valence bond theory explains the shape,
the formation and directional properties of
bonds in polyatomic molecules like CH4, NH3
and H2O in terms of overlap and
hybridisation of atomic orbitals.

82.  Overlapping of Atomic orbitals:
 When orbitals of two atoms come close to
form bond, their overlap may be positive,
negative or zero depending upon the
sign(phase) and direction of orientation of
amplitude of orbital wave function in space.

85.  Types of Orbital Overlap:
Depending upon the type of overlapping, the
covalent bonds are of two types, known as
sigma (σ ) and pi (π) bonds.
(i) Sigma (σ bond): Sigma bond is formed by
the end to end (head-on) overlap of bonding
orbitals along the internuclear axis.
(ii) The axial overlap involving these orbitals is
of three types:

86.  s-s overlapping: In this case, there is overlap
of two half-filled s-orbitals along the
internuclear axis as shown below:
 s-p overlapping: This type of overlapping
occurs between half-filled s-orbitals of one
atom and half filled p-orbitals of another
atoms.

87.  p-p overlapping: This type of overlapping
takes place between half filled p-orbitals of
the two approaching atoms.

88.  (ii) pi (π bond): π bond is formed by the
atomic orbitals when they overlap in such a
way that their axes remain parallel to each
other and perpendicular to the internuclear
axis. The orbital formed is due to lateral
overlapping or side wise overlapping.

89.  Strength of Sigma and pi Bonds
Sigma bond (σ bond) is formed by the axial
overlapping of the atomic orbitals while the
π-bond is formed by side wise overlapping.
Since axial overlapping is greater as
compared to side wise. Thus, the sigma bond
is said to be stronger bond in comparison to
a π-bond.

91.  Hybridisation
Hybridisation is the process of intermixing of
the orbitals of slightly different energies so as
to redistribute their energies resulting in the
formation of new set of orbitals of equivalent
energies and shape.

92.  Salient Features of Hybridisation:
(i) Orbitals with almost equal energy take part
in the hybridisation.
(ii) Number of hybrid orbitals produced is
equal to the number of atomic orbitals
mixed,
(iii) Geometry of a covalent molecule can be
indicated by the type of hybridisation.
(iv) The hybrid orbitals are more effective in
forming stable bonds than the pure atomic
orbitals.

93.  Conditions necessary for hybridisation:
(i) Orbitals of valence shell take part in the
hybridisation.
(ii) Orbitals involved in hybridisation should
have almost equal energy.
(iii) Promotion of electron is not necessary
condition prior to hybridisation.
(iv) In some cases filled orbitals of valence
shell also take part in hybridisation.

94.  Types of Hybridisation:
(i) sp hybridisation: When one s and one p-
orbital hybridise to form two equivalent
orbitals, the orbital is known as sp hybrid
orbital, and the type of hybridisation is called
sp hybridisation.
Each of the hybrid orbitals formed has 50% s-
characer and 50%, p-character. This type of
hybridisation is also known as diagonal
hybridisation.

96. (ii) sp2 hybridisation: In this type, one s and two
p-orbitals hybridise to form three equivalent
sp2 hybridised orbitals.
All the three hybrid orbitals remain in the same
plane making an angle of 120°. Example. A few
compounds in which sp2 hybridisation takes
place are BF3, BH3, BCl3 carbon compounds
containing double bond etc.

98.  (iii) sp3 hybridisation: In this type, one s and
three p-orbitals in the valence shell of an
atom get hybridised to form four equivalent
hybrid orbitals. There is 25% s-character and
75% p-character in each sp3 hybrid orbital.
The four sp3 orbitals are directed towards
four corners of the tetrahedron.

100.  The angle between sp3 hybrid orbitals is
109.5°.
A compound in which sp3 hybridisation
occurs is, (CH4). The structures of NH2 and
H20 molecules can also be explained with the
help of sp3 hybridisation.

101.  Formation of Methane Molecule ( CH4 ):
 Step -1: Formation of the excited state of a
Carbon atom:
 The carbon atom in the ground state takes up
some energy and goes to the excited
state. In this process, a pair of electrons in
2s orbital splits up and one of the electron
from this pair is transferred to empty 2pz
orbital. Thus the excited state has four half-
filled orbitals.

103.  Step – 2: Hybridization of Orbitals:
 One orbital of 2s and three orbitals of 2p mix
up forming four hybrid orbitals of equivalent
energy. These four new equivalent orbitals
are called sp3 hybrid orbitals. They are
identical in all respect.

104.  Bonds:
 Four sp3 hybrid orbitals of carbon atom
having one unpaired electron each overlap
separately with 1s orbitals of four hydrogen
atom along the axis forming four covalent
bonds (sigma bonds). The bonds between
carbon and hydrogen are sp3– s. Thus H – C –
H bond angles are 109.5°. The molecule is
tetrahedral. All C-H bonds in methane are of
equal strength.

106. Formation of Ammonia molecule:
In NH3, the valence shell electronic
configuration of nitrogen in the ground state is
2s22px12py12pz1 having three unpaired
electrons in the Sp3 hybrid orbitals and a lone
pair of electrons is present in the fourth one.
These three hybrid orbitals overlap with 1s
orbitals of hydrogen atoms to form three N-H
sigma bonds.

107.  The force of repulsion between a lone pair
and a bond pair is more than the force of
repulsion between two bond pairs of
electrons.
 The molecule thus get distorted and the bond
angle is reduced to 1070 from 109.50. The
geometry of such molecule will be pyramidal.
 The Atomic number of Nitrogen is 7. Its
configuration in ground state is 1s2, 2s2, 2p3

108.  Formation of water molecule:
 Electron Configuration:
 The atomic number of Oxygen is 8. Its
configuration in ground state is 1s2, 2s2, 2p4

109.  Hybridization :
 In the formation of water molecule one
2s orbital and three 2p orbitals of Oxygen
mix up forming four hybrid orbitals of
equivalent energy. These four new equivalent
orbitals are called sp3 hybrid orbitals. They
are identical in all respect. The two hybrid
orbitals have paired electrons and they are
non – bonding orbitals. Other two orbitals
are half-filled and they are bonding orbitals.
The nonbonding pairs of hybridized orbitals
are called lone pairs. These hybridized
orbitals are in the directions of four corners
of a regular tetrahedron.

110. The two half-filled (containing unpaired
electron) sp3 hybrid orbitals of oxygen overlap
axially with two half-filled 1s orbitals of two
hydrogen atoms separately to form two O-H
bonds (sigma bond). The remaining two hybrid
orbitals containing a lone pair of the electron
remains non bonded. The bond angle is
reduced to 104.5 from 109.5degree and the
molecule acquires V-Shape or angular gometry.

112.  Sp3 hybridiorbital

113.  In ethane molecule
 Sp3-sp3 sigma bond
 Three hybrid orbitals of each carbon atoms are
used in forming sp3-s sigma bonds with
hydrogen atoms.
 Ethane C-C bond length is 154pm and each C-
H bond length is 109pm.

114. 

115.  Sp2 hybridisation in C2H4
 It consists of one sp2-sp2 sigma bond
 One pi bond between p orbitals
 4 C-H sigma bond sp2-s[108pm]
 H-C-H bond angle is 117.6

116.  Sp hybridisation in ethyne.
 In the formation of ethyne molecule, both the
carbon atoms undergo sp-hybridisation
having two unhybridised orbital 2py and 2px.
 1. One C-C sigma bond
 2. two C-H sigma bond sp-s
 3.Two pi bonds
 Each of two unhybridised p orbitals of both
the carbon atoms overlaps sidewise to form
pi bonds between the carbon atoms.

118.  Hybridisation of elements involving d orbitals:

119.  Formation of PCl5-sp3d hybridisation:
 In PCl5 the five sp3d orbitals of phosphorus
overlap with the singly occupied p orbitals of
chlorine to form five P-Cl sigma bonds.
 Three P-Cl bonds lie in one plane and make
an angle of 120 with each other these bonds
are called as equatorial bonds.
 The remaining two P-Cl bonds one lying
above and other lying below the equatorial
plane, make an angle of 90 degree with the
plane. These bonds are called axial bonds.

120.  As the axial bond pairs suffer more repulsive
interaction from the equatorial bond pairs,
therefore axial bonds have been found to be
slightly longer and hence slighltly weaker
than the equatorial bonds, which makes PCl5
molecule more reactive.

121.  Formation of SF6: sp3d2 hybridisation

122.  Molecular Orbital Theory:
 1)The atomic orbitals overlap to form new
orbitals called molecular orbitals.
 When two atomic orbitals overlap or combine
,they lose their identity and form new orbitals.
The new orbitals thus formed are
called molecular orbitals.
 2)Molecular orbitals are the energy states of a
molecule in which the electrons of the molecule
are filled just as atomic orbitals are the energy
states of an atom in which the electrons of the
atoms are filled.


123.  3)A molecular orbitals gives the electron
probability distribution around a group of
nuclei just as atomic orbitals give the electron
probability distribution around the single
nucleus.
 4)Only those atomic orbitals can combine to
form molecular orbitals which have
comparable energies and proper orientation.
 For example :1s can combine with 1s and not
with 2s.
 5)The number of molecular orbitals formed is
equal to number of combining atomic
orbitals.

124.  6)When two atomic orbitals combine ,they form two new
orbitals called bonding molecular orbitals and antibonding
molecular orbitals.
 7)The bonding molecular orbitals has lower energy and
hence greater stability than the corresponding antibonding
molecular orbitals.
 8)The bonding molecular orbitals are represented by σ , π,
δ whereas the corresponding antibonding molecular orbitals
are represented by σ∗ , π∗, δ∗
 9)The shape of molecular orbitals formed depend upon the
type of the combining atomic orbitals.
 10)The filling of molecular orbitals takes place according to
same rules as those of the atomic orbitals. These are:

125.  1)Aufbau principle : Molecular orbitals are
filled in order of the increasing energies.
 2)Pauli exclusion principle: Molecular orbital
can have maximum of 2 electrons and these
must have opposite spin.
 3)Hund’s rule of maximum
multiplicity :Pairing of electrons in the
degenerate molecular orbitals does not take
place until each of them has got one electron
each.

126.  Formation of molecular orbitals
 An atomic orbital is an electron wave, the
waves of the two atomic orbitals may be in
phase or out of phase. Suppose ΨA and
ΨB represent the amplitude of the electron
wave of the atomic orbitals of the two atoms
A and B.

127.  Case 1: When the two waves are in phase so
that they add up and amplitude of the wave is
 Φ= ΨA + ΨB

128.  Case 2 : when the two waves are out of
phase, the waves are subtracted from each
other so that the amplitude of new wave is
 Φ ´= ΨA – ΨB

129.  The molecular orbitals formed by the additive
effect of the atomic orbitals is called bonding
molecular orbitals and the molecular orbitals
formed by the subtractive effect of atomic is
called antibonding molecular orbitals.
 The probability of finding the electrons in the
bonding molecular orbital increases where as
it decreases in the antibonding molecular
orbital.

130.  Atomic orbitals are represented by s, p, d ,
the bonding molecular orbitals are
represented by σ , π, δ and the
corresponding antibonding molecular orbitals
are represented by σ∗ , π∗, δ∗.

131.  Crest of the electron wave are usually given
by + sign and troughs a – sign.
 Bonding molecular orbital is formed by the
combination of + and + and – with – part of
the electron waves whereas antibonding
molecular orbital are formed by the overlap
of + with – part.
 The lowering of energy of the bonding
molecular than the combining atomic orbital
is called stabilization energy.
 Increase in energy of the antibonding
molecular orbitals is called destabilization
energy.

132.  In the bonding molecular orbital ,the electron
density in the internuclear region is high. As
a result the nuclei are shielded from each
other.
 In the antibonding molecular orbital, the
electron density in the internuclear region is
very low.
 The lowering of energy of the bonding
molecular than the combining atomic orbital
is called stabilization energy.
 Increase in energy of the antibonding
molecular orbitals is called destabilization
energy.

133.  Conditions for the combination of atomic orbitals
 1)The combining atomic orbitals should have
comparable energy
 For Ex: For homonuclear diatomic molecules , 1s
atomic orbital of one atom can combine with 1s
atomic orbital of another atom. 2s can combine
with 2s , 2p with 2p and so on. 1s cannot
combine with 2s, 2s cannot combine with 2p.
 2)The combining atomic orbitals must have
proper orientation so that they are able to
overlap to a considerable extent.
 3)The extent of overlapping should be large.
Greater the overlap, greater will be the electron
density between the nuclei.

134.  Types of molecular orbitals formed:
 1)If two atomic orbitals overlap along the
internuclear axis ,the molecular orbital formed is
called σ molecular orbital.
 2)If two atomic orbitals overlap sideways, the
molecular orbital formed is called π molecular
orbital.
 3)s orbitals are spherically symmetrical ,their
wave function has the same size in all the
directions.
 4)In p- orbital, one lobe is given a + sign and the
other a – sign.


135.  5)Overlapping of + part of the electron cloud
of one atom with + part of the electron cloud
of second atom implies addition of the
atomic orbitals leading to the formation of
bonding molecular orbitals.
 6)The overlap of + part of the electron cloud
of one atom with – part of the electron cloud
of the second atom means the subtraction of
the atomic orbitals leading to the formation
of antibonding molecular orbital.

136.  1s with 1s

137.  The wave function of two 1s atomic orbitals
can combine in two different ways :
 a) When both have the same sign.
 b)When they have different signs.
 If one of the wave function is assigned a +ve
sign, the other may be either +ve or -ve sign.
The bonding molecular orbital formed is
designated as σ (1s), σ indicating that the
overlap is along the internuclear axis and 1s
indicating that 1s atomic orbitals have
combined to form the molecular orbital. The
corresponding antibonding molecular orbital
is designated as σ∗

138. 2s with 2s
 The bonding and antibonding molecular
orbitals formed are designated as σ (2s)
and σ∗(2s).
2pz with 2pz
 Taking Z-axis as the internuclear axis, the
molecular orbitals formed are:

139.  2px with 2px

140.  Energy level diagram for Molecular orbitals:
 The first ten molecular orbitals may be
arranged in order of energy as follow:
 σ(1s) <σ∗(1s) < σ(2s) <σ∗(2s) < π(2px) =
π(2py) < σ(2pz) < π∗(2px) =π∗(2py) <π∗( 2pz)

142.  Relationship between electronic configuration and
Molecular behaviour
 1) Stability of molecules in terms of bonding and
antibonding electrons
 Number of electrons present in the bonding
orbitals is represented by Nb and the number of
electrons present in antibonding orbitals by Na.

143.  1) If Nb > Na ,the molecule is stable because
greater number of bonding orbitals are
occupied than antibonding orbital, resulting
in a net force of attraction.
 2) If Nb < Na , the molecule is unstable
because the antibonding influence is greater
than the bonding influence, resulting in net
force of repulsion.
 3) If Nb = Na ,the molecule is again unstable
because influence of electrons in the
antibonding molecular orbital is greater than
the bond influence of electron in the bonding
molecular orbitals.

144.  2) Stability of molecules in terms of bond
order
 Bond order is defined as half of the difference
between the number of electrons present in
the bonding and antibonding orbitals.
 Bond Order = ½ ( Nb – Na)
 The molecule is stable if Nb > Na ie. bond
order is positive. The molecule is unstable
if Nb < Na i.e. the bond order is negative or
zero.

145.  3) Relative stability of molecule in terms of
bond order
 For diatomic molecules ,the stability is
directly proportional to the bond order.
 A molecule with the bond order of 3 is more
stable than a molecule with bond order of 2
and so on.

146.  4) Nature of bond in terms of bond order :
 Bond order 1 ,2 and 3 mean single ,double
and triple bond.
 5) Bond length in terms of bond order:
 Bond length is found to be inversely
proportional to the bond order. Greater the
bond order, shorter is the bond length.

147.  Diamagnetic and paramagnetic nature of the
molecules
 If all the electrons in the molecule are paired,
it is diamagnetic in nature.
 If the molecules has some unpaired electrons
,it is paramagnetic in nature.
 Greater the number of unpaired electrons
present in the molecular or ion, greater is its
paramagnetic nature.

148.  HYDROGEN

149.  HELIUM

150.  LITHIUM

151.  DICARBON-C2

152.  OXYGEN

153. HydrogenBonding
When highly electronegative elements like
nitrogen, oxygen, flourine are attached to
hydrogen to form covalent bond, the electrons
of the covalent bond are shifted towards the
more electronegative atom.
Thus, partial positive charge develops on
hydrogen atom which forms a bond with the
other electronegative atom.
This bond is known as hydrogen bond and it is
weaker than the covalent bond.
For example, in HF molecule, hydrogen bond
exists between hydrogen atom of one molecule
and fluorine atom of another molecule.
It can be depicted as

154.  Types of H-Bonds
(i) Intermolecular hydrogen bond
(ii) Intramolecular hydrogen bond.
(i) Intermolecular hydrogen bond: It is formed
between two different molecules of the same or
different compounds. For Example, in HF
molecules, water molecules etc.
(ii) Intramolecular hydrogen bond: In this type,
hydrogen atom is in between the two highly
electronegative F, N, O atoms present within the
same molecule. For example, in o-nitrophenol,
the hydrogen is in between the two oxygen
atoms.