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3
...

• Chemical bonds are of many types
a) Ionic bond
b) Covalent bond
c) Co-ordinate covalent bond
d) Metallic bond, etc
...

• Bond formation is exothermic and bond breaking is endothermic
...
cal ;
H - H → H + H - 104 k
...

• In the bond formation, some energy is released and potential energy of system decreases
...
So that the attractive and repulsive
forces are balanced
...

• In exothermic reaction, the number of bonds formed in the products is greater than number of
bonds broken in the reactants (or)
• Strong bonds are formed in the products and weak bonds are broken in the reactants
...

Electronic Theory of Valency:• This was proposed by Kossel and Lewis
...

• Valence electrons are responsible for bonding process
...
Thus, all inert gases have octet and
helium has duplet configuration
...

• Atoms of all other elements contain less than 8 electrons in valence shell
...

• Attaining octet configuration in the valence shell is called octet rule or octet theory
...
g
...

• Octet configuration can be achieved by loosing or gaining or mutual sharing of electrons
...

Atom - Valence = Core
VALENCE or VALENCY:
It is the combining capacity of an element i
...
, number of bonds formed by the element
...

• The strong electrostatic force of attraction between oppositely charged ions which are formed by
the transfer of the electrons is called Ionic bond
...
e atoms of different electronegativities
...

• It cannot be formed between same or similar atoms
...
Most ionic compound is CsF (Cesium fluoride)
• To form an ionic bond, the electronegatives between combining atoms should be greater than 1
...

• Ionic bond is generally formed between electropositive and electronegative element or less
electronegative and more electronegative elements
...


•

FACTORS FAVOURABLE FOR IONIC BOND FORMATION
The ease of formation of ionic bond depends on the case of formation of cation and anion
...
: F > Cl > Br > I
Eg
...
Ps Electron affinity: Atoms with high electron
will form cations readily
...

Eg
...
: Cl > Br > I
Electron affinity decreases
IP → increases
Ease of formation decreases
...

3) Charge: Cation with less positive charge Charge: Anion with less negative charge is
readily formed
...
: F– > O–2 > N–3
Eg
...

charge
...
8
2
...
18
Inert gas configuration
Pseudolnert gas configuration
Higher lattice energy also favours ionic bond formation
b)

LATTICE ENERGY:- (ν)
The amount of energy released when the oppositely charged gaseous ions combine to form one mole of
solid ionic crystal (or)
The amount of energy absorbed to separate one mole of solid ionic crystal into oppositely charged
gaseous ions is called lattice energy
...
2 kcal
Na+(g) + Cl-(g)
→ NaCl(s) + 782 KJ/mole
+
NaCl(s) → Na (g) + Cl⎯(g) – 782 KJ/mole
•
•

In a given ionic crystal, there are attractions between opposite charges and repulsions between
electron clouds of cation and anion
...

PE att n = −

NAZ + Z − e 2
r

PE rep n = +

NBe 2
rn

Lattice energy (u) = −

NAZ + Z − e 2 NBe 2
+
r
rn

N → Avagadro's number
A → Madelung's constant
Z+ → Positive charge
Z– → Negative charge
e → Charge of e⎯
B → Repulsive co-efficient
n → Born exponent
Lattice energy is inversely proportional to the sum of radii of cation and anion
...

Born-Haber's cycle:
The basis for Born-Haber's cycle is Hess's law
...

3

Chemical Bonding and molecular structure
Lattice energy cannot be determined by direct experimental methods
...
Eg
...
5 KJ / mole)
One mole sodium reacts with half mole chlorine gas to form solid NaCl crystal
...

Solid sodium on heating directly changes to vapour state and the heat energy is called sublimation
energy
...
7 KJ/mole)
2) Dissociation of Cl2
Cl2 molecule dissociate into Cl - atoms
...

1
D
Cl2 (g) → Cl(g) +
...
55KJ / mol )
2
2

3) Ionisation of Na
Electron is removed from Na to form sodium cation
...
(ΔH = 492
...

e−

Cl (g) ⎯⎯ → Cl − (g) − EA (ΔH = −361
...
Energy released in
the process is called lattice energy
...
to Hess' law,
–Q= +S+

D
+I−E−U
2

– 410
...
7 + 119
...
82 –316
...

+I
Na (g)

Cl(g)

−E

⎯⎯ →
⎯

Cl−g ) + Na +
(

(g )

↑ S ↑ D/2 ↓ - U
1
Crystal structures: ⎯⎯ → NaCl
⎯
Cl2
Na (s ) +
(g )
−Q
2

4

Chemical Bonding and molecular structure
The three dimensional network in which the cations and anions are arranged at optimum distances is
called crystal lattice
...

Generally, a cation is surrounded by specific number of anions and anion is surrounded by a specific
number of cations
...

Each Na+ ions is surrounded by 6Cl- ions and each Cl⎯ ion by 6 Na+ ions
...

• In some ionic crystals like CaF2 and Na2O, co-ordination numbers are different for cation and
anion
...

F⎯ is 4
...

O2– is 8
...

rc
= limiting radius
...
e
...

ra

rc
ra

upto

co-ordination
Examples
number

Shape

0
...
155 - 0
...
225 - 0
...
414 – 0
...
732 - 0
...
C
...
C
...


Most common co-ordination numbers are 6 and 8
...

These unit cells in repetitions in 3 dimensions will give entire crystal lattice
CRYSTAL STRUCTURE OF NaCl: NaCl has face centered cubic lattice structure (FCC)
rc
is 0
...
Each Na+ ion is surrounded by six Cl- and each Clra

•

Co-ordination number is 6 because

•

ion is surrounded by 6 Na+ ions
...

Contribution of body central Na+ ion towards 1 unit cell = 1 × 1 = 1
...

4

1
2

Contribution of face central ions towards 1 unit cell = 6 × = 3
...

8

CRYSTAL STRUCTURE OF CsCl : CsCl has body centered cubic lattice [BCC]
Its co-ordination number is 8 ∵

rc
= 0
...

Number of ion pairs or formula units or molecules per unit cell = 1
...

2) Melting and boiling points:
Ionic compounds have high MPs and BPs due to strong interionic attractions
...

4) Ionic bond is non-directional in nature :
As the ionic bond is non directional in nature
...

5) Reactions of Ionic compounds :
Reactions in between Ionic compounds are very fast in aqueous solution because they does not
involve any reshuffling of bonds
...

6) Solubility:
Ionic compounds dissolve in polar solvents like H2O due to ion-dipole interactions
...

Covalent bond:
It was proposed by Lewis
...

In covalent bonding, both atoms will contribute and both will share
...

• Maximum number of bonds (covalent) formed between 2 atoms is 3
...

• Pure or 100% covalent bond is the bond formed between same atoms
...

Favourable conditions for formation of covalent bond: [Fazan's rule]
• Cation should be smaller and anion should be larger in size
...

• The electronegativity difference should be less than 1
...

Covalent bond is denoted by "⎯"
Eg
...

Covalency:
It is the number of electrons contributed by an atom or the number of covalent bonds formed by an
atom
Covalency:
covalency of hydrogen is
1
H2
covalency of oxygen is
2
O2
covalency of nitrogen is
3
N2
2
H2O covalency of oxygen is
1
H2O covalency of hydrogen is
3
NH3 covalency of nitrogen is
4
CO2 covalency of carbon is
PCl5 covalency of phosphorous is 5
covalency of sulphur is
6
SF6
Properties of covalent compounds:
• They exist as either gases or liquids due to weak Vanderwaal's forces
...

• Electrical conductance :
• They are bad conductors as they donot contain ions
...

Solubility:
They are soluble in non-polar solvents like CCl4, chloroform, C6H6 and insoluble in polar solvents like
water
...
e
...

• Some covalent compounds like HCl, HF, HI, etc
...

• Some covalent compounds like sugar, urea, glucose, alcohol, HF are soluble in polar solvents like
water due to H  bonding
...

Best solvent for Ionic and polor solvents is water
...
e
...

Ex: BeCl2 (4 e- s); BCl3 (6 e⎯ s ); PCl5 (10 e- s) SF6 (12 e- s)

•

It fails to explain single electron or odd electron bond
...
Odd e⎯ bond: O2–, NO, NO2, ClO2

•

It could not explain the shapes and bond angles of various molecules
...


Valence bond theory:
The basis of VBT is Schrodinger's wave equation i
...
wave mechanics
...

•

This theory was proposed by Hietler and London and developed by pauling and slater
...

•

The e⎯s in the overlapping orbitals must be with opposite spins
...
e
...


•

The direction in which overlapping orbitals are concentrated, the bond is formed in that direction
...


•

The molecule will be stable because the bonding electron density is in consideration along the inter
nuclear axis and that electron density keeps the two atoms attracted to each other
...

p–p>s–p>s–s

•

Smaller atoms involve in greater overlapping
...

It involves head on (or) end – on – end
overlapping
...

Pi electron density lies above and below the axes
...

8

Chemical Bonding and molecular structure
(s - s, s - p, p - p)
5) It can be formed by pure valence
orbitals or hybrid orbitals
...

7) The first formed bond between two
atoms is always sigma bond
...

If p-orbital is involved in overlapping, it may be
sigma or pi
It is formed only after the formation of sigma
bond
...


9) One of the two lobes is involved in over Both the lobes of p - orbital are involved in bond
lapping
formation
10) Free rotation of orbitals is possible Free rotation of orbitals is restricted
around sigma bond
...

• In triple bond, one σ and 2 π bonds are present
...

Postulates:
The electron pairs present in valence shell of central atom will be situated around it so that repulsions
are minimum
...

Order of repulsions in between various Electron pairs:
Lone pair - lone pair > lone pair - bond pair > bond pair - bond pair
• Lone pair is attracted by one nucleus where as bond pair by two nuclei
...
If one or
more lone pairs are present, it will have irregular geometry
Thus, shape of molecule depends on extent of mutual repulsions between various electron pairs
...

∴ Its shape is regular tetrahedral
In ammonia, there are three bond pairs and one lone pair
...
e
...

In H2O, these are two bond pairs and two lone pairs
∴ shape is irregular i
...
, angular
...

VSEPR theory is useful to predict the shapes of molecules, and type of hybridisation, based on the no
of electron pairs present in valence shell of central atom
...
of
electro
n pairs
2

Bon
d
pairs
2
3
2
4

lone
pairs

2+2
=2
2
3+3
= 3
2
6+0
=3
2
6+0
=3
2

Group number of central atom + number of bonded pairs
2

sp

linear

sp2

Trigonal planar

sp2 (2b
...
p)

Angular

sp2 (3b
...
p)

Trigonal planar

5+3
=4
2

sp3 (3b
...
p)

Pyramidal

6 + 3 −1
2

sp3

Pyramidal

=4

Hybridisation

Shape

Angle

6

180°
120°
109°

3

1

sp3

Pyramidal

107°

2

sp3

-

sp3d

4

1

sp3d

2
3
-

sp3 d
sp3 d
sp3d2

5

1

sp3 d2

4

2

sp3 d2

Angular
Trigonal
bipyramidal
Distorted
tetrahedral
T
Linear
Octahedral
Distorted
octahedral
Square planar

90°, 120°
180°

3
2
6

5

Linear
Trigonal planar
Angular
Tetrahedral

5

4

sp
sp2
sp2
sp3

2

3

1
-

10

Examples
BeCl2, CO2, HCN
BCl3,BF3, SO2
SO2, SnCl2
CH4, CCl4,CF4
NH3, H3O+ (Hydronium
ion)
H2O, H2S, Cl2O, OF2
PCl5, PF5

-

SCl4, SF4

90°, 180°
180°
90°

ClF3, BrF3,ICl3
XeF2, ICl2
SF6

-

ClF5, IFs

90°

XeF4

Chemical Bonding and molecular structure
7

-

sp3 d3

6

1

sp3 d3

7

CCl4

number of e⎯ pairs =

PCl5

number of e⎯ pairs =

SF6

number of e⎯ pairs =

IF7

number of e⎯ pairs =

H 2O

number of e⎯ pairs =

NH3

number of e⎯ pairs =

NF3

number of e⎯ pairs =

PCl3

number of e⎯ pairs =

POCl3 number of e⎯ pairs =
SOCl2 number of e⎯ pairs =
XeF2 number of e⎯ pairs =
XeF4 number of e⎯ pairs =
XeF6 number of e⎯ pairs =
ClF3

number of e⎯ pairs =

4+4
=4
2
5+5
=5
2
6+6
=6
2
7+7
=7
2
6+2
=4
2
5+3
=4
2
5+3
=4
4
5+3
=4
2
5+3
=4
2
6+2
=4
2
8+2
=5
2
8+4
=6
2
8+6
=7
2
7+3
=5
2

Pertagonal
bipyramidal
Distorted
octahedral

72°, 90°
-

IF7
XeF6
tetrahedral

sp3d

trigonal bipyramidal

sp3 d2

octahedral

sp3d3

pentagonal bipyramidal

(2 b
...
p) sp3

Angular

(3 b
...
p) sp3

Pyramidal

sp3(3b
...
p)

Pyramidal

sp3(3b
...
p)

Pyramidal

sp3(4b
...
p)

Tetrahedral

sp3(3b
...
p)

Pyramidal

sp3d (2b
...
p)

Linear

sp3d2 (4b
...
p)

Square planar

sp3d3(6b
...
p)

distorted octahedral

sp3d (3b
...
p)

T shape

Distorted tetrahedral
sp3d (4b
...
p)
The above formula to calculate the number of electron pairs is applicable only for simple molecules or
ions mentioned above
...

Co-ordinate covalent bond (dative bond):
It is proposed by Sidgewick
...

• Thus, in covalent bond both atoms will contribute and share, but in co-ordinate bond, one
contributes and both will share
• To form co-ordinate covalent bond, there must be an electron pair donor and electron pair
acceptor
...

Co-ordinate bond is semi polar bond
Formation of co-ordinate bond involves over-lapping between completely filled orbital of donor
with vacant orbital of acceptor ∴Co-ordinate bond is rigid and directional like covalent bond
...


Eg : (i) H3 N + BF3 → {H3N→BF3}

...
O
...
Some of the properties are in
between to those of Ionic and covalent compounds due to semi polar nature of the bond
...

2) Their melting and boiling points are low due to weak intermolecular forces
...

4) They are soluble in non-polar solvents and insoluble in polar solvents like water
...

6) The reactions in between co-ordinate compounds are very slow as they involve shuffling of bonds
...

• Due to difference in electronegativity more electronegative atom develops partial negative charge
and less electronegative atom develops partial positive charge
...

Hydrogen bond
H δ+ − O δ− − − − − − H δ+ − O δ−
|
Hδ +

|
Hδ

+

H - bond is denoted by a broken line (
...

12

Chemical Bonding and molecular structure
•
•
•

In H - bond, H is sandwiched between two electronegative atoms
...

H - bond is formed by molecules or ions where H is covalently bonded to more electronegative and
smaller atoms like F, O, N
...

• Among HF, H2O, HI; the strongest H - bonds are formed by HF molecules
...
H – F
2) H – OH …… H – OH
3) H – NH – H …
...
F > H
...
N
• When compared to covalent bond hydrogen bond is weaker, but longer i
...
, higher bond length
...
76 Å O (Hydrogen bond)
...

Types of H bonds :
Intermolecular H – bonds :
H - bond is formed between two same molecules or different molecules i
...
, H-bond between H of
one molecule and more electronegative atom of another molecules
Eg :

H–F –……H –F
NH3, RNH2, R2NH, ROH
 +
 –
 –
H –O H…O H
Carboxylic acids, Glucose, fructose, Para-nitrophenol, para chlorophenol
Parahydroxy benzal dehyde
OH
...
HO
Intramolecular H - bond :
H- bond is formed with in the same molecule i
...
, H-bond between + Hydrogen and  atom both
belonging to same molecule
Eg
...

Salicyaldehyde
Orthonitrophenol
H
|
C = Os−
O-H

s+

O
13

O

Chemical Bonding and molecular structure

H
N
O

CHO
C6H 4
ΟΗ

OH
C6H 4
ΝΟ2

Effect of H - bonding: (Intermolecular)
Due to H - bonding,
1) Molecular association increases
2) Melting and boiling points increase
3) Voltaile nature decreases
4) Solubility in water increase
5) Physical state may change
The above effects are observed in case of intermolecular H-bonding but not in the case of
intramolecular H-bonding
...
V A group Hydrides
NH3
CH4
SiH4
PH3
GeH4
AsH3
SnH4
SbH3
BiH3
PbH4
Though, molecular weight of NH3 is less, its BP is much higher than those of PH3 and ASH3 because of
H - bonding
...

• H2O is liquid while H2S is gas due to
H- bonding in H2O
...

• Certain covalent substances like glucose, fructose, sugar, urea, alcohol, amines, carboxylic acids
are soluble in water due to H – bonding
...

• Paranitrophenol is less volatile because it forms intermolecular H- bonds
...

14

Chemical Bonding and molecular structure
•

Certain substances like acetic acid, benzoic acid hydroflouric acid will exist in dimeric form due to
H-bonding
...
But, HF can form H-bonds even in vapour
state
...
It is due to
1) H2O forms double the number of H - bonds than HF
...

Thus, it is not necessary to break all H - bonds in HF to vapourise it
...

In such molecules, positive or negative charges are not developed on any atom
...

Eg
...

•
If covalent bonds are present between dif
...
More electronegative atom shares more and less electron negative atom shares less As a
result, more electronegative develops negative charge; and less electronegative atom develops
positive charge
...

The above polar molecules are called dipoles
...

HF > HCl > HBr > HI
O - H > S - H > Se - H > Te - H
O-H>N-H>S-H
(EN difference decrease s, polarity decreases)
N - Cl < P - Cl < As - Cl < Bi - Cl
(EN difference increases, polarity increases)
I - F > Cl - F > I - Cl > Br - Cl
(EN difference decreases, polarity decreases)
Dipole moment :
The magnitude of polarity in the molecule is expressed in terms of dipole moment value
...

μ=e×d
μ=δ×l
Where μ → dipole moment
e, δ → charge
d, l → distance (bond length)
Units of dipole moment : Debyes
1 Debye = 10–18 e
...
u cm
1 Debye = 3
...

•
Dipole moment is a vector quantity
...

•
If only bond - pairs are present the molecule has regular shape and its dipole moment will be zero
due to mutual cancellation of the bond moments
...

Eg : CCl4, CF4, SiCl4, SiF4, CH4, SnCl4
•
The molecule in which central atom contains one or more lone pairs will have irregular geometry
and such molecules are polar and they have net dipole moment value
...

•
Thus, the molecule with polar bonds may be polar or non-polar as discussed above
...

APPLICATIONS OF DIPOLE MOMENT:
1) The shape of the molecule and hybridisation of central atom can be predicted
...

μcis > μtrans
3) Ortho, meta, para Isomeras of a compound can be distinguished
μortho > μ meta > μpara
4) % Ionic character can be calculated
...

% Ionic Character =

μ obs
× 100
μ cal

16

Chemical Bonding and molecular structure
Examples:
1
...
03 debyes
...
28 Å
Calculate % Ionic character
...
8 × 10–10 × 1
...
8 × 1
...
8 ×1
...


1
...
8 %
4
...
285

Dipole moment of
...
92 Debyes If bond length of HF is 0
...
Calculate its Ionic character
...
8 × 10–18× 0
...
8 × 0
...
92
10 −18 × 100
×
4
...
9
10 −18

192
× 100 = 44%
48 × 9
θ
μobs = 2 × bond moment × cos
2

=

3
...
95 debyes
...
Calculate the bond
moment of S - H bond
(cos 460 = 0
...
95 = 2 × x × cos 46
0
...
65
2x =

95 19
=
65 13

2x = 1
...
73
Though EN difference between N and F is greater than that of between N and H Even though, both
NH3 and NF3 are pyramidal μ NH3 > μ NF3
In NH3, the lone pair contributes in the same direction as those of bond dipoles where as in NF3 lone
pair contributes in opposite direction as those of bond – dipoles

...

N

HHH

FFF
In case of AB2 type tri - atomic molecules μ value increases with decrease in the bond angle
METALLIC BOND:
The force of attraction that binds the metal atoms in metallic crystal is called metallic bond
...

17

Chemical Bonding and molecular structure
1) Free e- theory
2) Valence bond theory
3) Molecular orbital theory
...

•
All metal atoms loose their valence e⎯ in metallic crystal
...

•
The force of attraction between positively charged metal ions and negative electron pool is called
metallic bond
...
Thus, metal
is imagined to be positively charged ions immersed in a sea of mobile electrons
...

•
In case of stronger metallic bonds, metals are hard with high melting and boiling points
...

•
Though this theory could explain conductivity, metallic luster and some other properties, it fails to
explain the differences in properties between various metals
...

•
Acc
...

•
A metal atom is bonded to its neighbouring atoms by the sharing of e- pairs
...

•
In metallic crystal, each metal atom is surrounded by numerous metal atoms
...

•
Because of resonance, the metallic crystal is stable and metallic bonds are stronger (metal atom)
...

Bond length depends on
i) Size of atom: with increase in size of bonded atom, bond-length increases
...
With increase in bondorder, bond length value decreases
1
...
34 A 0

C=C

C≡C

1

2

3

C−C

B
...
:

1
...

|
|
H−C > H−C = > H−C ≡
|
S − sp
S − sp 2
3
S − sp

iv) With increase in polarity, bond length decreases
...

Generally, the bond length will be in between that of single bond length value and double bond
length value
...
48 A 0

1
...
: In O3, bond length between two oxygens is 1
...

BOND ANGLE:It is the angle between the two adjacent bonded atoms
...

ii) S-character: With increase in S-character bond angle increases
...

iii) Repulsions in between the electron pairs : Due to repulsions in between lone pairs, bond angle
decreases
...


...

EN ↓
H 2O
104°
NH3
107°
H 2S
92°
PH3
93°
BA ↓s
H2Se 91°
ASH3 91° 30’
H2Te 90°
SbH3 91°
If the EN of bonded atoms decreases, bond angle increases
OF2 → 104°
Cl2O → 111°
BOND ENERGY:The amount of energy released when one mole of bonds are formed (or) the amount of energy absorbed
to break one mole of bond
...

H + H → H – H ; 104 kcal
H – H → H + H ; –104 kcal
In case of polyatomic molecule, the bond energy of particular bond is the average of sum of all bond
energies
...
Bond energy of CH4 is 360 k
...
cal/mole
...
cal
In Ethane there are six C - H bonds and one C - C bonds
...
cal
Factors influencing bond energy :
1) Size of bonded atom:
with increase in size of bonded atom, bond energy decreases
...

3) Presence of lone pairs: With increase in the number of lone pairs, bond energy decreases
...


− C− C− > − N− N− >
|

|

|

|


...



...
O
...

i) sp 3 − sp 3 > sp 2 − sp 2 > sp − sp
p−p > s−p > s−s
5) Polarity: With increase in polarity, bond energy increases
...

7) Type of bond fission: Bond can be fissioned by homolytic or heterolytic way
...

hom olytic fission
A − B ⎯⎯ ⎯ ⎯ ⎯ ⎯ → A • + B •
⎯
heterolyti fission
A − B ⎯⎯ ⎯ ⎯c⎯ ⎯⎯→ A + + B −

If the bond energy is more, the molecule is more stable and reactivity is less
...

Predicting the type of bonds:•
The bond between two electronegative atoms is covalent bond
...

•
The bond between electropositive and electronegative element is ionic bond
...
5H2O
20

Chemical Bonding and molecular structure
Ni(CO)4
Fe(CO)5
K4[Fe(CN)6]

Number of sigma bonds = Atomicity - 1
Atomicity → number of atoms in a compound
σ=3–1
=2
π=2
Eg
...
The number of molecular
orbitals formed is equal to the number of atomic orbitals involved and they belong
to the molecule
...
e
...

The order of energies of molecular orbitals is bonding orbitals < Non-bonding
orbitals < Anti-bonding orbitals
...

21

Chemical Bonding and molecular structure

•
•
•
•

Molecular orbital of lower energy is known as bonding molecular orbital and of
higher energy is known as antibonding molecular orbital
...

Aufbau rule, Pauli’s exclusion principle and Hund’s rule are applicable to molecular
orbitals, during the filling electrons
Their shape is governed by the shape of atomic orbitals
The increasing order of relative energies of M
...


σ 1s < σ *1s < σ 2 s < σ * 2 s < π 2 pz =
π 2 py < σ 2 px < π * 2 pz = π * 2 p y < σ * 2 px
for more than 14 electrons

σ 1s < σ *1s < σ 2s < σ * 2 s < σ 2 px
< ⎡π 2 pz = π 2 p y ⎤
⎣
⎦

< ⎡π * 2 pz = π * 2 p y ⎤ < σ * 2 px
⎣
⎦

Atomic and Molecular Orbitals Main differences
Atomic Orbitals

Molecular orbitals

1) They belong to one specific atom only

1) They belong to all the atoms in a molecule

2) They are the internal characteristic of an atom
...


3) They have simple shapes of geometries
...
etc

4) The molecular orbitals are named as

σ ,π ,δ


...


5) The stabilities of these orbitals are less than 5) The stabilities of these orbitals are either
bonding and more than the antibonding
σ - molecular orbital
1)

more or less than the atomic orbitals orbitals

Difference between σ and π MO’s
π- molecular orbital

Formed by the end on overlap along the

1) Formed by the sidewise overlap

internuclear axis

erpendicular to inter nuclear axis

2)

Overlapped region is very large

2) Over lapped region is small

3)

Rotation about the internuclear axis is

3) Rotation about the inter nuclear axis is
...


•

Note : KK = σ 1s 2
...

BOND ORDER: The relative stability of a molecule can be determined on the
basis of bond order
...
It
is equal to one half of the difference between the number of electrons in the
bonding and antibonding molecular orbitals
...
But bond
order may be fractional in some cases
...
5
2
It is paramagnetic
−
•
O2 (Super oxide ion):
Total number of electrons (16 +1) = 17
...
5
Bond order =
2
2
It is paramagnetic
•
Peroxide ion ( O 2- ) - Total number of electrons (16 + 2) =18
...
Hence paramagnetic
FORMAL CHARGE
...
Formal charge may be regarded as the charge that an atom in a
molecule would have if all the atoms had the same electronegativity
...
In case of a polyatomic ions, the net charge is possessed the
real ion as a whole and not by an particular atom
...

Q f = [ N A − N M ] = [ N A − N LP − 1/ 2 N BP ]
Where
NA= number of electrons in the valence shell in the free atom
NM= number of electrons belonging to the atom in the molecule

24

Chemical Bonding and molecular structure

NLP = number of electrons in unshared pairs, i
...
number of electrons in lone pairs
NBP = number of electrons in bond pairs, respectively

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