Lewis and Kossel were the first to successfully explain chemical bonding in terms of electrons. They proposed that atoms bond to achieve stable electron configurations like noble gases. Lewis pictured atoms as a positively charged nucleus surrounded by 8 electrons. Atoms bond by transferring or sharing electrons to achieve stable octets. Lewis symbols were introduced to represent valence electrons. Covalent bonds form when atoms share valence electrons in order to gain full outer shells. Resonance structures represent situations where no single Lewis structure adequately describes the molecule.

1. Chemical Bonding And
Molecular Structures

3. K o s s e l L e w i s
A p p r o a c h T o
C h e m i c a l B o n d i n g
• In 1916, Kossel and Lewis were the first to
become independently successful in giving
a satisfactory explanation about the
formation chemical bond in terms of
electrons.
• They were the first to provide some logical
explanation of valence which was based on
the inertness of noble gases.

4. L e w i s a p p r o a c h
• He pictured the atom in terms of a positively charged
‘Kernel’ and the outer shell which could accommodate
a maximum of eight electrons.
• He assumed that these eight electrons occupy the
corners of a cube which surround the ‘Kernel’.
• Lewis postulated that atoms achieve the stable octet
when they are linked by chemical bonds.
• In sodium and chlorine, this can happen by transfer
of electron from sodium to chlorine.
• In case of other molecules like CL2 the bond is formed
by sharing of a pair of electrons between atoms.

5. L e w i s s y m b o l
In formation of a molecule, only the
outer shell electron take part in
chemical combination and they
are called valence electrons.
The inner shells do not get involved
in the combination process.
G.N. Lewis, an American chemist
introduced simple notations to
represent valence shell electrons
in an atom and these are known
as Lewis symbols.

7. S i g n i f i c a n c e o f
L e w i s s y m b o l
• The number of dots represent the number of
valence electrons.
• This number of valence electrons help to
calculate the common or group valence of
that element.
• The group valence of the element is generally
either equal to the number of dots in Lewis
symbol or 8 minus the number of dots.

8. • Other important facts given by Kossel are-
• In periodic table, the highly electronegative halogens and the
highly electropositive alkali metals are separated by 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 electron by the respective atoms.
• The negative and positive ions thus formed attain stable
noble gas electronic configurations.
• The negative and positive ions are stabilized by electrostatic
attraction.

9. Example
• Na  Na+ + e-
[N e ] 3s 1 [N e ]
• Cl + e-  Cl-
[N e } 3s 2 3p 5 [N e ] 3s 2 3p 6 o r [A r ]
=> Na+ + Cl-  NaCl

10. O c t e t r u l e
• According to electronic theory of chemical
bonding, atoms can combine either by
transfer of valence electrons from one atom
to another or by sharing of valence in order
to have an octet in their valence shells. this is
known as octet rule.

11. R e c a l l : E l e c t r o n s a r e
d i v i d e d b e t w e e n c o r e
a n d v a l e n c e
e l e c t r o n s .
A T O M c o r e
v a l e n c e
N a 1s 2 2s 2 2p 6 3s 1 [N e ]
3s 1
B r [A r ] 3d 10 4s 2 4p 5 [A r ] 3d 10
4s 2 4p 5
C o v a l e n t
B o n d i n g
Covalent bond is the sharing of the VALENCE
ELECTRONS of each atom in a bond
Br Br

12. C o v a l e n t b o n d
• The bond formed by atoms when they share their
valence electrons to gain octet is known as covalent
bond.
• 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 pairs
, a triple bond is formed.

13. C o v a l e n t b o n d
d i a g r a m s
H 2O
N H 3
C H 4

14. L e w i s r e p r e s e n t a t i o n
o f s i m p l e
m o l e c u l e s (t h e L e w i s
s t r u c t u r e s )
• The Lewis dot structures provide a picture of bonding in
molecules and ions in terms of shared pairs of electrons and
the octet rule.
• The Lewis dot structures can be written by adopting the
following steps.
– 1)the total number of electrons required for writing the
structures are obtained by adding the valence electrons of
the combining atoms.
– 2) 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.

15. – 3)one must know the chemical symbols of the
combining atoms and their skeletal structures.
– 4)in general, the least electronegative atom
occupies the central position.
– 5)after accounting for shared pairs of electrons
for single bonds, the remaining electron pairs are
either utilized for multiple bonding or remain as
lone pairs.

17. F o r m a l c h a r g e
• A formal charge (FC) is the charge assigned to
an atom in a molecule, assuming that electrons in
a chemical bond are shared equally between
atoms, regardless of relative electro negativity.
• 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 Lewis structure.

18. • Formal charge = (total no. of valence
electrons in the free
atom)-(total no. of non
bonding electrons)-
(1/2(total no. of bonding
electrons)

19. • Formal charge on ‘1’ =6-2-1/2(6)=+1
• Formal charge on ‘2’ =6-4-1/2(4)=0
• Formal charge on ‘3’ =6-6-1/2(2)=-1

20. • Formal charges do not indicate real charge separation
within the molecule. These charges help in keeping
track of the valence electrons in the molecule.
• Formal charges help in the selection of the lowest
energy structure from a number of possible Lewis
structures for a given species.
• Generally the lowest energy structure is the one with
the smallest formal charges on the atoms. The formal
charge is a factor based on pure covalent view of
bonding in which electron pairs are shared equally
by neighboring atoms.

21. L i m i t a t i o n s o f
t h e o c t e t r u l e• It is not universal. There are
exceptions to it.
1)the incomplete octet of the
central atom
In some compounds, the
number of electrons
surrounding the central
atom is less than eight.
Examples- LiCl, BeH2 and
BCl3

22. 2) Odd-electron molecule
in molecules with an odd number of
electrons like nitric oxide, NO and nitrogen
dioxide, the octet rule is not satisfied for all
atoms.

23. 3)The Expanded Octet
Elements in and beyond the third period of
the periodic table have, apart from 3s and 3p
orbitals, 3d orbitals also available for
bonding. In a number of compounds of these
elements there are more than eight valence
electrons around the central atom. This is
termed as expanded octet.
Examples-PF5, SF6, H2SO4

24. 4) Octet rule is based on the chemical inertness
of noble gases. However, some noble gases
also combine with oxygen and fluorine to
form a number of compounds like
XeF2, KrF2, XeOF2 etc.
5) This theory does not account for the shape of
the molecules.
6) It does not explain the relative stability of
the molecules being totally silent about the
energy of a molecule.

26. B o n d
p a r a m e t e r s
• Bond length
bond length is defined as the equilibrium
distance between the nuclei of two bonded
atoms in a molecule.
Bond lengths are measured by spectroscopic, X-
ray diffraction techniques etc.
Each atom of the bonded pair contribute to the
bond length.

27. • In case of a covalent bond, the contribution
from each atom is called the covalent radius
of that atom.
• The covalent radius is measured
approximately as the radius of an atom’s core
which is in contact with the core of an
adjacent atom in a bonded situation.
• The van der Waals radius represents the
overall size of the atom which includes its
valence shell in a non bonded situation.

28. • Bond angle
It is defined as the angle between the orbitals
containing bonding electron pairs around the
central atom in a molecule or complex ion.
Bond angle is expressed in degree.

29. • Bond enthalpy
It is defined as the amount of energy required to break one mole of bonds
of a particular type between two atoms in a gaseous state.
The unit of bond enthalpy is kJ per mol.
Example-
The H-H bond enthalpy in hydrogen molecule is 435.8 kJ/mol.
H2 -------> H(g)+H(g) ΔaH=435.8 kJ/ mol
O2 (O=O) (g) ------> O(g)+O(g) ΔaH=498 kJ/mol

30. • The larger the bond dissociation enthalpy, stronger will be
the bond in the molecule.
HCl---H + Cl
• In case of polyatomic molecules, the measurement of bond
strength in the following way-
let us take the example of H2O molecule.
H2O---H+OH; ΔaH1=502 kJ/mol
OH---H+O; ΔaH2=427 kJ/mol
In this case, the average bond enthalpy is taken.
average bond enthalpy=502+427=464.5 kJ/mol
2

31. B o n d o r d e r
• In Lewis description of covalent bond, the
bond order is given by the number of bonds
between the two atoms in a molecule.
• For example, the bond order of h2 which has a
single shared pair of electrons is 1
• The bond order of O2 with 2 shared pairs of
electrons is 2
• The bond order of N2 with 3 shared pairs of
electron is 3.

32. • For N2, the bond order is 3 and its bond
enthalpy is 946 kJ/mol and it is one of the
highest for diatomic molecule.
• Isoelectronic molecules and ions have identical
bond orders.
example- N2 and CO have bond order 3.
A general correlation useful for understanding
the stabilities of molecules is that: with increase
in bond order, bond enthalpy increases and
bond length decreases.

34. R e s o n a n c e
s t r u c t u r e s• A single Lewis structure may be inadequate for
representation of a molecule.
• According to the concept of resonance, whenever
a single Lewis structure cannot describe a
molecule accurately, a number of structures
with similar energy, position of nuclei, bonding
and non-bonding pairs of electrons are taken as
the canonical structures of the hybrid which
describes the molecule accurately.

35. • For example, O3 molecule can
be represented by two
structures as shown in the
figure.
• For O3 these two structures
constitute the canonical
structures or resonance
structures and their hybrid
shown by the third figure
represents the structure of o3
more accurately. this is called
resonance hybrid.

36. • In general it may be stated that
– Resonance stabilizes the molecule as the energy
of the resonance hybrid is less than the energy of
any single canonical structure
– Resonance averages the bond characteristics as a
whole.
Resonance Energy
It is the difference between the actual bond energy of
the molecule and that of the most stable of the
resonating structures (having least energy).

38. Polarity o f B o n d s
The existence of a hundred percent ionic or covalent bond represents an ideal situation. In
reality no bond or a compound is either completely covalent or ionic. Even in case of
covalent bond between two hydrogen atoms, there is some ionic character.
When covalent bond is formed between two similar atoms, for example in H2, O2, Cl2, N2 or F2,
the shared pair of electrons is equally attracted by the two atoms. As a result electron pair
is situated exactly between the two identical nuclei. The bond so formed is called non polar
covalent bond. Contrary to this in case of a heteronuclear molecule like HF, the shared
electron pair between the two atoms gets displaced more towards fluorine since the
electronegativity of fluorine (Unit 3) is far greater than that of hydrogen. The resultant
covalent bond is a polar covalent bond.
As a result of polarisation, the molecule possesses the dipole moment (depicted below) which can
be defined as the product of the magnitude of the charge and the distance between the
centres of positive and negative charge. It is usually designated by a Greek letter ‘μ;’.
Mathematically, it is expressed as follows :
Dipole moment (μ) = charge (Q) x distance of separation (r)
Dipole moment is usually expressed in Debye units (D).The conversion factor is
1 D = 3.33564 x 10– -30 C m
where C is coulomb and m is meter.

39. • Further dipole moment is a vector quantity and is depicted by a
small arrow with tail on the positive centre and head pointing
towards the negative centre. For example the dipole moment of HF
may be represented as :
• The shift in electron density is symbolised by crossed arrow ( )
above the Lewis structure to indicate the direction of the shift.
• In case of polyatomic molecules the dipole moment not only depend
upon the individual dipole moments of bonds known as bond
dipoles but also on the spatial arrangement of various bonds in the
molecule. In such case, the dipole moment of a molecule is the vector
sum of the dipole moments of various bonds. For example in H2O
molecule, which has a bent structure, the two O-H bonds are oriented
at an angle of 104.50. Net dipole moment of 6.17 x 10– -30 C m
(1D = 3.33564 x 10– -30 C m) is the resultant of the dipole
moments of two O-H bonds.
• Peter Debye, the Dutch chemist received Nobel prize in 1936 for his
work on X-ray diffraction and dipole moments. The magnitude of
the dipole moment is given in Debye units in order to honour him.
• Net Dipole moment, μ = 1.85 D = 1.85 x 3.33564 x 10– -30 C m =
6.17 10– 30 C m

40. • The dipole moment in case of BeF2 is zero.
This is because the two equal bond dipoles
point in opposite directions and cancel the
effect of each other.
• In tetra-atomic molecule, for example in BF3,
the dipole moment is zero although the B – F
bonds are oriented at an angle of 120° to one
another, the three bond moments give a net
sum of zero as the resultant of any two is
equal and opposite to the third.

41. • In case of NH3 and NF3 molecule. Both the molecules
have pyramidal shape with a lone pair of electrons
on nitrogen atom. Although fluorine is more
electronegative than nitrogen, the resultant dipole
moment of NH3 ( 4.90 x 10–3 the orbital dipole due to
lone pair is in the same direction as the resultant
dipole moment of the N – H bonds, whereas in NF3the
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.

44. The partial covalent character of ionic bonds was discussed by
Fajan’s in terms of the following rules:
• The smaller the size of the cation and the larger the size
of the anion, the greater the covalent character of an ionic
bond.
• The greater the charge on the cation, the greater the
covalent character of the ionic bond.
• For cations of the same size and charge, the one, with
electronic configuration (n-1)dnnso, typical of transition
metals, is more polarising than the one with a noble gas
configuration, ns2 np6, typical of alkali and alkaline
earth metal cations. The cation polarises the
anion, pulling the electronic charge toward itself and
thereby increasing the electronic charge between the two.
This is precisely what happens in a covalent
bond, i.e., build-up of electron charge density between the
nuclei. The polarising power of the cation, the
polarisability of the anion and the extent of distortion
(polarisation) of anion are the factors, which determine the
per cent covalent character of the ionic bond.

46. V a l e n c e S h e l l
E l e c t r o n P a i r
R e p u l s i o n
T h e o r y
• Sidgwick and Powell in 1940, proposed a simple theory
based on the repulsive interactions of the electron pairs in
the valence shell of the atoms. It was further developed and
redefined by Nyholm and Gillespie (1957).
• The main postulates of VSEPR theory are as follows:
• • The shape of a molecule depends upon the number of
valence shell electron pairs (bonded or non-bonded) around
the central atom.
• • Pairs of electrons in the valence shell repel one another
since their electron clouds are negatively charged.
• • These pairs of electrons tend to occupy such positions in
space that minimise repulsion and thus maximise distance
between them.
• • The valence shell is taken as a sphere with the electron pairs
localising on the spherical surface at maximum distance

47. • • A multiple bond is treated as if it is a single electron pair
and the two or three electron pairs of a multiple bond are
treated as a single super pair.
• • Where two or more resonance structures can represent a
molecule, the VSEPR model is applicable to any such
structure.
• The repulsive interaction of electron pairs decrease in the
order:
• Lone pair (lp) – Lone pair (lp) > Lone pair (lp) – Bond pair
(bp) > Bond pair (bp) – Bond pair (bp)
• Nyholm and Gillespie (1957) refined the VSEPR model by
explaining the important difference between the lone pairs
and bonding pairs of electrons. While the lone pairs are
localised on the central atom, each bonded pair is shared
between two atoms. As a result, the lone pair electrons in a
molecule occupy more space as compared to the bonding
pairs of electrons. This results in greater repulsion between
lone pairs of electrons as compared to the lone pair – bond
pair and bond pair – bond pair repulsions. These repulsion
effects result in deviations from idealised shapes and
alterations in bond angles in molecules.

48. • For the prediction of geometrical shapes of molecules
with the help of VSEPR theory, it is convenient to
divide molecules into two categories as (i) molecules
in which the central atom has no lone pair and (ii)
molecules in which the central atom has one or more
lone pairs.
• The VSEPR Theory is able to predict geometry of a
large number of molecules, especially the compounds
of p-block elements accurately. It is also quite
successful in determining the geometry quite-
accurately even when the energy difference between
possible structures is very small. The theoretical basis
of the VSEPR theory regarding the effects of electron
pair repulsions on molecular shapes is not clear and
continues to be a subject of doubt and discussion.

50. • Shape (g e o me t r y ) o f S o me
S i mp l e M o l e c u l e s /I o n s w i t h
C e n t r a l I o n s h a v i n g O n e o r
mo r e L o n e P a i r s o f
E l e c t r o n (E ).

52. • S h a p e s o f M o l e c u l e s
c o n t a i n i n g B o n d P a i r a n d
L o n e P a i r

55. O R B I T A L
O V E R L A P
C O N C E P T• If orbitals of 2 atoms are mixed with each other partially during bond
formation, then the phenomenon is called as overlapping of orbitals.
• TYPES OF ORBITAL OVERLAPS:
• 1.Positive overlap: If the symmetry of both the atomic orbitals is the same, then it
is called as positive overlap i.e. if the symmetry of the overlapping orbitals is
either positive, or negative, then it is a type of positive overlap.
• 2.Negative overlap: If the symmetry of the atomic orbitals is not the
same, i.e., one is positive and the other negative, then it is a type of negative
overlap.
• 3.Zero overlap: If overlap of orbitals present in 2 different planes takes place
, then it is called as zero overlap. E.g.. Px overlaps with Py (in real
situation, overlapping does not takes place.)

56. T Y P E S O F O V E R L A P S
R E S U L T I N G I N S I G M A
B O N D F O R M A T I O N

57. T Y P E S O F
O V E R L A P S
R E S U L T I N G
I N P I B O N D
F O R M A T I O N
Due to lateral/sideways
overlap of P-P orbitals present
in the same plane, Pi bond is
formed.

58. C O M P A R I S I O N
SIGMA BONDS
• It is a strong bond.
• Electron cloud is
symmetrical along the inter-
nuclear axis.
• There can be free rotations of
atoms around this bond.
• These are less reactive.
• Shape of the molecule is
determined by these bonds.
• Sigma bonds have
independent existence.
PI BONDS
• It is a weak bond.
• Electron cloud is
asymmetrical.
• Free rotation is not possible
around this bond.
• These are more reactive.
• These bonds do not affect the
shape of the molecule.
• Pi bond always exists with a
sigma bond.

59. H Y B R I D I Z A T I O N
The intermixing of different atomic orbitals of approximately equal energy levels to
produce hybrid orbitals before bond formation is called as HYBRIDIZATION.
Here arrangement of hybrid orbitals are such that there is minimum repulsion in
between the hybrid orbitals.
No. of orbitals mixed=No. of hybrid orbitals produced.
DIFFERENT TYPES OF HYBRIDISATIONS:
Sp Hybridization
NO. of hybrid orbitals produced=2
Structure = LINEAR
Bond angle = 180 degree …..e.g. BeF₂

60. s p
H Y B R I D I D S A T I O N

61. • SP2 Hybridization:
• No. of hybrid orbitals produced = 3
• Arrangement of these orbitals = TRIGONAL PLANAR
• Bond angle=120 degree……..ex. BF₃, etc.
• SP3 Hybridization:
• No. of hybrid orbitals = 4
• Arrangement= TETRAHEDRAL
• BOND ANGLE…..
• IN CH₄ = 109.28 degree
• IN NH₃ = 107.3 degree.
• Here, in methane, sigma bond is formed between H and C atom due to
overlapping of sp3 orbital of H atom with S orbital of H atom.
• Structure of NH₃ – TRIANGULR PYRAMIDAL.

62. • SP2
Hybridization
• SP3
Hybridization

63. • SP3d Hybridization:
• NO. of hybrid orbitals produced = 5
• Structure = TRIGONAL BIPYRAMIDAL
• BOND ANGLES:
• Equatorial = 120 degree.
• Axial = 90 degree.
• If lone pair of electron is present at central atom, its position is always
equatorial.Ex.PCl5
• Length of axial bonds is longer than that of equatorial bonds because of
minimum repulsion.
• IT may have different shapes according to no. of lone pairs it has:
• 1 lone pair – seesaw shape………e.g. SF₄
• 2 lone pairs – bent T shape………E.g. BrF₃
• 3 Lone pairs – Linear shape

65. • SP3d2 Hybridization:
• No. of hybrid orbitals produced = 6
• Arrangement = OCTAHEDRAL……e.g. SF₆ ( lone pairs = 0)
• Shape may be:
• Square pyramidal…..e.g. BrF₅ ( lone pairs = 1)
• Square planar…….e.g. XeF4₄ (lone pairs = 2)
• SP3d3 Hybridization:
• No. of hybrid orbitals = 7
• Arrangement = PENTAGONAL BIPYRAMIDAL
• BOND ANGLES:
• Axial with equatorial = 90 degree.
• Equatorial to equatorial = 72 degree.

66. • SP3d2
Hybridization
• Sp3d3
Hybridization

67. R U L E S R E G A R D I N G
H Y B R I D I Z A T I O N
• Only orbitals of approximately same energy levels can take part.
• No. of orbitals mixed = No. of hybrid orbitals produced.
• Most hybrid orbitals are similar but not always identical in shape. They
may differ from one another in their orientation in space.
• The electron waves in hybrid orbitals repel each other and this tend to the
farthest apart.
• Hybrid orbitals can form only sigma bonds.
• Depending on the number and the nature of the orbitals undergoing
hybridization, various types of hybrid orbitals directing towards the
corners of specified geometrical figures come into existence.

70. • C O N D I T I O N S F O R
C O M B I N A T I O N O F
A T O M I C O R B I T A L S
• For atomic orbitals to combine, resulting in the
formation of molecular orbitals , the main conditions
are :
• The combining atomic orbitals should have almost the
same energies. For example, in the case of diatomic
molecules, 1s-orbital of one atom can combine with 1s-
orbital of the other atom, but 1s-orbital of one atom
cannot combine with 2s-orbital of the other atom.
• The extent of overlap between the atomic orbitals of the
two atoms should be large.
• The combining atomic orbitals should have the same
symmetry about the molecular axis. For
example, 2Pxorbital of one atom can combine with
2Px orbital of the other atom but not with 2Pz orbital .
• Note : It may be noted that Z-axis is taken as the
inter-nuclear axis according to modern conventions.

71. D e s i g n a t i o n s o f
M o l e c u l a r O r b i t a l s
Just as atomic orbitals are designated as
s, p, d, f etc molecular orbitals of diatomic molecules are
named σ (sigma) ,
π (pi) , δ (delta) etc.
M O L E C U L A R O R B I T A L S
The molecular orbitals which are cylindrically symmetrical
around inter-nuclear axis are called σ - molecular orbitals.
The molecular orbital formed by the addition of 1s orbitals
is designated as σ 1s and the molecular orbital formed by
subtraction of 1s orbitals is designated as σ * 1s .
Similarly combination of 2s orbital results in the
formation of two
2 s - molecular orbitals designated as σ 2s and σ * 2s

73. 1.Determine the number of electrons in the molecule. We get the
number of electrons per atom from their atomic number on the
periodic table. (Remember to determine the total number of electrons,
not just the valence electrons.)
2.Fill the molecular orbitals from bottom to top until all the electrons
are added. Describe the electrons with arrows. Put two arrows in each
molecular orbital, with the first arrow pointing up and the second
pointing down.
3. Orbitals of equal energy are half filled with parallel spin before
they begin to pair up.

74. Stability of the molecule with bond order.
Bond order = 1/2 (#e- in bonding MO's - #e- in
antibonding MO's)
We use bond orders to predict the stability
of molecules :-
• If the bond order for a molecule is equal to zero, the
molecule is unstable.
• A bond order of greater than zero suggests a stable
molecule.
• The higher the bond order is, the more stable the bond.
We can use the molecular orbital diagram to predict whether
the molecule is paramagnetic or diamagnetic. If all the
electrons are paired, the molecule is diamagnetic. If one or
more electrons are unpaired, the molecule is paramagnetic.

75. 1. The molecular orbital diagram for a diatomic hydrogen molecule, H2, is
• The bond order is 1. Bond Order = 1/2(2 - 0) = 1
• The bond order above zero suggests that H2is stable.
• Because there are no unpaired electrons, H2 is diamagnetic.
2. The molecular orbital diagram for a diatomic helium molecule, He2, shows the following.
• The bond order is 0 for He2. Bond Order = 1/2(2 - 2) = 0
• The zero bond order for He2suggests that He2is unstable.
• If He2did form, it would be diamagnetic

76. 3. The molecular orbital diagram for a diatomic oxygen molecule, O2, is
• O2has a bond order of 2. Bond Order = 1/2(10 - 6) = 2
• The bond order of two suggests that the oxygen molecule is stable.
• The two unpaired electrons show that O2is paramagnetic

77. Diatomic molecules are molecules composed only of two atoms,
of either the same or different chemical elements. The prefix di-
is of Greek origin, meaning two. Common diatomic molecules
are hydrogen (H2), nitrogen (N2), oxygen (O2), and carbon
monoxide (CO). Seven elements exist as homonuclear diatomic
molecules at room temperature: H2, N2, O2, F2, Cl2, Br2, and I2.
Many elements and chemical compounds aside from these
form diatomic molecules when evaporated. The noble gases do
not form diatomic molecules: this can be explained using
molecular orbital theory (see molecular orbital diagram).

78. INTRODUCTION
1.Two H atoms in their ground state configuration
come together and form a single bond. The bond
formation stabilizes both atoms and, therefore, is lower
in energy than the atomic orbitals. This is also
observed in Valence Bond Theory, which implies that
each H atom in H2shares its electron with one another,
so that both can achieve the stable configuration of He.
2.On top of that, MO Theory allows one to compute the
amount of energy released from a bond formation and
a distance between two bonded atoms as well as predict
the magnetic property of a molecule (or a substance).
For H2, the bond strength is -432 kJ/Mol, and the
bond length is 74 angstrom (or 74 pm). H2 is a
diamagnetic molecule because the electrons paired up;
therefore, it is not attracted by a magnetic field.

79. B o n d i n g a n d A n t i -b o n d i n g
m o l e c u l a r o r b i t a l s i n H 2
1.Each H atom has a 1s atomic orbital. When two H atoms come to a proper proximity, their 1s
orbitals interact and produce two molecular orbitals: a bonding MO and an anti-bonding MO.
2.If the electrons are in phase, they have a constructive interference. This results in a bonding
sigma MO (σ1s). This MO has an increased probability of finding electrons in the bonding
region.
Figure 2: Schematic representation of the bonding molecular orbital σ(1s)
If the electrons are out of phase, they have a destructive interference. This results in an anti-
bonding sigma MO (σ*1s). This MO has a decreased probability of finding electrons in the
bonding region. (Valence Bond Theory does not explain this phenomenon.)
Figure 3:Schematic representation of antibonding molecular orbital σ*(1s)
Note that there is a nodal plane in the anti-bonding

80. B o n d o r d e r i n H 2
Bond order = 1/2 (#e- in bonding MO - #e- in antibonding MO)
For H2, bond order = 1/2 (2-0) = 1, which means H2 has only one
bond. The antibonding orbital is empty. Thus, H2 is a stable
molecule.
Again, in the MO, there is no unpaired electron, so H2 is
diamagnetic

82. H Y D R O G E N B O N D
In compounds of hydrogen with strongly electronegative
elements, such as fluorine, oxygen and
nitrogen, electron pair shared between the two atoms
lie far away from the hydrogen atom. As a result, the
hydrogen atom becomes highly electropositive with
respect to the other atom. This phenomenon of charge
separation in the case of hydrogen fluoride is
represented as . Such a molecule is said to be polar .
The molecule behaves as a dipole because one end
carries a positive charge and the other end a negative
charge. The electrostatic force of attraction between
such molecules should be very strong. This is
because the positive end of one molecule is attracted by
the negative end of the other molecule . Thus, two or
molecules may associate together to form larger
cluster of molecules. This is illustrated below for the
association of several molecules of hydrogen fluoride.

83. • The cluster of HF molecules may be described as
(HF)n.
• It may be noted that hydrogen atom is bonded to
fluorine atom by a covalent bond in one molecule and
by electrostatic force or by hydrogen bond to the
fluorine atom in the adjacent molecule . Hydrogen
atom is thus seen to act as a bridge between the two
fluorine atoms.

84. • The hydrogen bond is represented by a dotted line. The
solid lines represent the original(covalent ) bond
present in the molecule.
• Chlorine, bromine and iodine are not as highly
electronegative as fluorine and therefore, the shared
pair of electrons in the case of HCl , HBr and HI do
not lie as far away from hydrogen as in the case of
HF. The tendency to form hydrogen bond in these
cases is therefore less.
• Water molecule, because of its bent structure, is also a
dipole, oxygen end carrying a negative charge and
hydrogen end carrying a positive charge. Hydrogen
bond taking place in this case as well, as represented
below:
• The cluster of water molecules may be described as
(H2O)n

85. • The nature of hydrogen bond
• The hydrogen bond is a class in itself. It arises from
electrostatic forces between positive end (pole) of one
molecule and the negative end(pole) of the other
molecule generally of the same substance. The
strength of hydrogen bond has been has been found to
vary between 10 - 40 kJ mol−1 (i.e., 6.02 x
1023 bonds) while that of a covalent bond has been
found to be of the order of 400 kJ mol−1 . Thus a
hydrogen bond is very much weaker than a covalent
bond. Consequently, the length of hydrogen bond is
bigger than the length of a covalent bond.
• In the case of hydrogen fluoride, for instance, while
the length of the covalent bond between F and H atoms
is 100 pm, the length of hydrogen bond between F and
H atoms of neighbouring molecules is 155 pm.

86. • T y p e s o f h y d r o g e n
b o n d i n g
• Hydrogen bonding may be classified into
two types :
• I n t e r m o l e c u l a r
h y d r o g e n b o n d i n g
This type of hydrogen bonding involves electrostatic
forces of attraction between hydrogen and
electronegative element of two different molecules of
the substance. Hydrogen bonding in molecules of
HF, NH3 , H2O etc. are examples of intermolecular
hydrogen bonding.

87. • I n t r a m o l e c u l a r
h y d r o g e n b o n d i n g
This type of bonding involves electrostatic forces of
attraction between hydrogen and electronegative
element both present in the same molecule of the
substance. Examples o-nitrophenol and
salicylaldehyde.
• p-Nitrophenol , on account of large distance between
two groups , does not show any intramolecular
hydrogen bonding. On the other hand, it shows the
usual inter molecular hydrogen bonding , as
illustrated below:

88. • As a result of intermolecular hydrogen bonding, the
para derivative undergoes association, resulting in an
increase in molar mass and hence an increase in
boiling point. In ortho derivative, on account
of intramolecular hydrogen bonding , no such
association is possible. Consequently, the ortho
derivative is more volatile than the para derivative.
Thus, while ortho nitrophenol is readily volatile in
steam , para nitrophenol is completely non-volatile.
The two derivatives can thus be separated from each
other by steam distillation.

89. • Density in solid state(ice) is less than that in liquid
state . This is some what unusual because in most
substances density in solid is more than that in
liquid state.
• Water contracts when heated between 0°C and 4°C .
This is again unusual because most substances
expand when heated in all temperature ranges.
• Both these peculiar features are due to hydrogen
bonding, as discussed below :

90. • In ice, hydrogen bonding between H2O
molecules is more extensive than in liquid
water. A substance in solid state has a
definite structure and the molecules are
more rigidly fixed relative to one another
than in the liquid state. In ice, the H2O
molecules are tetrahedrally oriented with
respect to one another.At the same time
, each oxygen atom is surrounded
tetrahedrally by four hydrogen atoms, two
of these are bonded covalently and the
other two by hydrogen bonds.The
tetrahedral open cage-like crystal structure
of ice. The central oxygen atom A is
surrounded tetrahedrally by the oxygen
atoms marked
1,2, 3 and 4.The hydrogen bonds are
weaker and therefore, longer than covalent
bonds. This arrangement gives rise to an
open cage-like structure , as shown in the
Fig. There are evidently a number of ‘holes'
or open spaces.

91. • These holes are formed because the hydrogen bonds
holding the H2O molecules in ice are directed in
certain definite angles . In liquid water such
hydrogen bonds are fewer in number. Therefore, as ice
melts, a large number of hydrogen bonds are broken.
The molecules, therefore, move into the ‘hole' or open
spaces and come closer to one another than they were
in the solid state. This results in a sharp increase in
density . The density of liquid water is, therefore
higher than that of ice.
• As liquid water is heated from 0°C to 4°C, hydrogen
bonds continue to be broken and the molecules come
closer and closer together. This leads
to contraction. However, there is some expansion of
water also due to rise in temperature as in other
liquids. It appears that up to 4°C, the former effect
predominates and hence the volume increases as the
temperature rises.

92. • It can be easily realised that without
hydrogen bonding , water would have existed
as a gas like hydrogen sulphide. In that case
no life would have been possible on this globe.
• Hydrogen bonding also exists in all living
organisms, whether of animal or of vegetable
kingdom. Thus, it exists in various
tissues, organs, blood, skin and bones in
animal life. It plays an important role in
determining structure of proteins which are
so essential for life.

93. • Hydrogen bonding plays an important role in
making wood fibres more rigid and thus makes it an
article of great utility. The cotton, silk or synthetic
fibres owe their rigidity and tensile strength to
hydrogen bonding. Thus hydrogen bonding is of
vital importance for our clothing as well. Most of our
food materials also consists of hydrogen bonded
molecules. Sugars and carbohydrates , for example,
have many -OH groups. The oxygen of one such group
in one molecule is bonded with -OH group of another
molecule through hydrogen bonding. Hydrogen
bonding is thus a phenomenon of great importance in
every day life.