1. Chemical bonding
Valency?
Definition
No. of unpaired electrons in an atom is called valency of that atom
e.g.
Valency= 01
Valency= 01
Assignment:- write valencies of Mg, Ca, Iodine, Xe

2. Stability of atoms
 Stability= surety of existence.
 Stability is need of everybody; students study to get a secure career.
 Everybody utilize energy to get stability
 Atoms having unpaired electrons are unstable
 Like everything in the universe atoms want to be stable
 Atoms get stable when they have all paired electrons.
unstable

3. Octate rule:
Atoms are stable when they have 08 electrons in their last shell.
Just count electrons in the last shell.
So, to know atom is stable of unstable; we must write electronic configuration.
But it is very difficult work to do??????
What is the easier solution????
Electrons in last shell= 04

4. How do atoms complete 08 electrons in the last
shell?
unstable
Need 07 electrons
to become stable
Very difficult..?
There is an
easy way
Loose one electron
in the last shell
unstable
Need 01 electrons
to become stable
Very easy..?
stable
stable
Assignment:-
find
out
what
Br
and
Li
can
do
to
become
stable?
Get 08 electrons by
loosing and gaining

6. unstable
Need 03 electrons
to become stable
Loose 05 e or
get 03 e?
Need 04 electrons
to become stable
Loose 04 e or
get 04 e?
Get 08 electrons
by mutual sharing
unstable
Lesson learnt! Atoms help each other to become stable.
Atoms help each other by two ways: 1) by sharing electrons 2) by gaining/loosing es
This help of atoms for each other is called chemical bond.

7. Chemical Bond
• Bond =attraction
• Attraction between atoms that holds them together
• All atoms trying to achieve a stable octet
Types of chemical bonds
Covalent bond Ionic bond
Chemical bond formed
by mutual sharing of
electrons
Chemical bond formed
by loosing or gaining
electrons

8. Which bond will be formed between two atoms?
Covalent or ionic bond?
Wants to lose e Wants to gain e
Ionic bond will be formed
between these two atoms
Wants to gain e
Wants to gain e
covalent bond will be formed
between these two atoms

9. Covalent Bond is formed;
 When both atoms want to gain electrons, e.g., Cl-Cl, N-N
 When both atoms want to share electrons e.g., C-C
 When one atom wants to share electron and other atom wants to gain electron e.g., C-Cl, C-
O
Ionic Bond is formed;
 When one of the atoms wants to lose electrons, no matter other atom wants to share or gain
electrons e.g., Na-Cl, Na-C etc.

10. Chemical bond formed by mutual sharing of electrons
But how does atoms mutually share electrons…???? Answer is given by two
theories 01) Valence Bond Theory
02) Molecular Orbital Theory
Covalent Bond

11. 01) Valence Bond Theory
 It was proposed by Heitler and London in 1927
 According to this theory, covalent bond is formed
when two half filled atomic orbitals overlap with
each other
 Electron of each orbital now lives as pair of electron
 This pair of electron is shared by both atoms
 But atomic orbital of each atom is still separate, so
no new orbital is formed
Overlapping = the penetration of orbitals in
each other is called overlapping of orbitals.

12. Types of orbital overlapping
Head-to-head overlapping
=sigma bond
(a covalent bond formed by
Head-to-Head overlap of
Orbitals Is called sigma bond)
along inter nuclear axis
Side ways overlapping
=pi bond
(a covalent bond formed by Side
ways overlap of orbitals Is called
pi bond)
parallel to internuclear axis
Head
Head
Side
way
Side
way
inter nuclear axis
inter
nuclear
axis

13. Head-to-head Overlapping (sigma bond)
Overlapping along inter nuclear axis
H H
1S orbital 1S orbital S-S orbital overlap
H
1S orbital
F
2p orbital S-p orbital overlap
F
2p orbital
F
2p orbital p-p orbital overlap
Types of Head-to-head overlap
01
03
02
all overlaps are equally strong…?
NO

14. Which overlap is strong overlap..? Charge Distribution
S-orbital
Uniform Charge
Distribution
p-orbital
Directional Charge
Distribution
 For a strong bond formation; electron must remain in the region where orbitals are
overlapping
 Due to uniform charge distribution in s- orbital, electron can not be available in the region
where orbitals are overlapping
 P-orbital has directional distribution and electrons are found in the head of p-orbital.
Therefore, electron will be available in the region where orbitals are overlapping
 Therefore;
 p-p>s-p>s-s

15. Side ways overlapping
=pi (π) bond
Weak bond

16. Valence Shell electron pair repulsion theory
 This theory was proposed by; Sigwick and Powell in 1940
 There are two types of electron pairs in the valence shell of central
atom I) lone pairs II) Bond pairs
 Being –vely charged, electron pairs repel each other
 Strength of repulsion is lone pair-lone pair>lone-pair-bond pair>bond
pair-bond pair repulsion
 Due to this repulsion, electron pairs in the central atom arrange
themselves in such a way that they experience minimum repulsion
 The pattern of arrangement gives geometrical shape to molecule

17. Shapes of molecules according to VSPERT

18. What is the Shape of Ammonia (NH3) molecule
according to VSPERT..?
Steps:
1. write the molecular formula
2. write the central atom
3. Write the surrounding atom
4. Structure of molecule
5. Count total e- pairs in valence shell of central atom
6. Count bonding e- pairs in valence shell of central
atom
7. Count lone e- pairs in valence shell of central atom
NH3
N
H
04
03
01

19. Hybridization
 This concept was given by Linus Pauling
 During bond formation Atomic orbitals of central atom mix with
each other to make a new type of orbital called a hybrid molecular
orbital.
 Only those atomic orbitals mix which have electrons
 The number of hybrid orbitals formed is the same as the number of
atomic orbitals that were mixed.
 The type of hybrid orbital formed depends on the types of atomic
orbitals mixed.
 Hybridized orbitals are lower in energy compared to nonhybridized
orbitals.
CH4
01 S orbital + 03 p orbitals=04 sp3

20. Water-H2O
Central Atom (O)
01 S orbital + 03 p orbitals=04 sp3

21. Ammonia-
NH3
Central Atom (N)
01 S orbital + 03 p orbitals=04 sp3

22. Boron trifrlouride-
BF3
Central Atom (B)
01 S orbital + 02 p orbitals=03 sp2

23. Beryllium Chlouride-BeCl2
Central Atom (Be)
01 S orbital + 01 p orbitals=02 sp

24. Types of
hybridization
Sp3
Sp2
Sp
Mixing orbitals Hybridization Shape of molecule Bond angle
01 S + 03 p Sp3 tetrahedral 109.5°
01 S + 02 p Sp2 trigonal 120°
01 S + 01 p Sp linear 180°

25. Limitations of Valence Bond Theory
1) It gives No information about bond order of a molecule
2) It Can not explains magnetic properties of a molecule
3) It can not explain 04 valency of Carbon, 03 Valency of
boron, and 02 valency of Beryllium

26. Molecular Orbital Theory
 It was proposed by Mullikan, Hund and Huckle in 1927
 According to this theory, covalent bond is formed when two half
filled atomic orbitals of similar energy combine with each other and
generate new orbitals called as molecular orbitals
 In atomic orbital electron is influenced of parent nucleus; while in
molecular orbital, electron is influenced by both nuclei.
 The number of molecular orbitals formed is the same as the number
of atomic orbitals that were combined.
 Two types of MO are formed; one having higher energy, and other
having lower energy
 higher energy MO is called Anti bonding molecular orbital
 Lower energy MO is called bonding molecular orbital
 Electrons are filled in molecular orbitals as per basic electronic
configuration rules; Afbau, hund rule etc.

27. Advantages of Molecular Orbital Theory
1) It gives information about bond order of a molecule
2) It explains magnetic properties of a molecule
bond order of a molecule= how many bonds will be
formed.
If the molecule has any unpaired electron, it will be
paramagnetic (attracted to external magnetic field)
If no unpaired electron, it will be diamagnetic (repelled
by external magnetic field)

28. Valence Bond Theory Molecular Orbital Theory
Only unpaired electrons of valence shell are
involved in bonding
All electrons of valence shell are involved in
bonding
Atomic orbitals do not loose their identity Atomic orbitals convert into molecular orbitals
VBT does not talk about bond order MOT helps determine bond order of the molecule
Magnetic properties of molecules can not be
explained
MOT explains the magnetic properties
Valence Bond Theory Vs Molecular Orbital Theory

29. Hydrogen Molecule-H2
02 00
= 01
Anti-Bonding orbital
Bonding orbital
There is no
unpaired electrons
in the molecule;
therefore, it is
diamagnetic

30. Oxygen Molecule-O2
08 04
= 02
There are two
unpaired electrons
in the molecule;
therefore, it is
paramagnetic

31. Ionic Character in covalent bond
Covalent bond= mutual sharing of electrons.
Mutual sharing between two same atoms
Homoatomic molecule
Mutual sharing between two different atoms
Heteroatomic molecule
Is electron pair equally shared between two atoms..?
Each atom of molecule try to attract shared pair of electron. The stronger atom is successful in
keeping shared pair of electron slightly towards itself.
In homoatomic molecules both atoms are of same strength. Therefore; e- pair is equally shared.
In heteroatomic molecule both atoms have different strength, therefore; e- pair is slightly towards
stronger atom
Atom to which e- pair is slightly closer gets slightly –ve charge on it
While the other atom becomes slightly +ve.
(Polar Molecule)
(Non-Polar Molecule)

32. Dipole Moment
Di=two pole=charges
Molecule which has ionic character is called a dipole/polar molecule
Dipole Moment µ
It is measurement of polarity of molecule
Dipole Moment Formula ... A dipole moment is the product of the magnitude of the
charge (q) and the distance between the centers of the positive and negative charges
(d)
µ=q.d
SI unit= coulomb.meter
Common unit=debye
1 debye =3.335× 10-30 C m
For Nonpolar molecules
µ=0
For polar molecules
µ>0

33. Bond Energy
Amount of energy required to break the bonds of 1 mole of substance is called Bond Energy
Factors Affecting bond energy
B.E ∝ Ionic character in covalent bond
B.E 1/∝ Bond Length
Bond Length=distance between centers of
covalently bonded atoms
1.2 Ao
1.34 Ao

35. States of Matter
H2O H2O
H2O
H2O

36. Difference in
arrangement of
molecules
Tightly
packed
Each molecule is
free to move
Molecules are far
apart
Can not
move
Molecules are
close together
but can move
freely
How molecules
stay together?
Solid Liquid Gas
States of matter

37. Why there are different states
of matter??
Van der walls force
weak electric forces that
attract molecules to one
another.

39. How to decide which will be the
normal state of a given
substance….?
Normal
temperature=room
temperature= 25 C
Normal pressure= 1
atm
State of matter depend
on distance between
molecules.
State of matter depend
on temperature and
pressure.
Boiling Point:-the temperature at
which the material transforms into
the gas phase
Melting point:-The temperature at
which solid changes its state to
liquid at atmospheric pressure

40. States of Matter-
GASES
A gas is one of the four fundamental states of matter
consisting of particles that have neither a defined volume nor
shape.

41. Collisions of Gas Particles

42. Collisions of Gas Particles

43. Kinetic Theory

44. Kinetic Molecular Theory
Postulates of the Kinetic Molecular Theory of Gases
1. Gases consist of tiny particles (atoms or molecules)
2. These particles are so small, compared with the distances between
them, that the volume (size) of the individual particles can be assumed
to be negligible (zero).
3. The particles are in constant random motion, colliding with the walls of
the container. These collisions with the walls cause the pressure exerted
by the gas.
4. The particles are assumed not to attract or to repel each other.
5. The average kinetic energy of the gas particles is directly proportional
to the Kelvin temperature of the gas

45. Kinetic Molecular Theory (KMT)
1. …are so small that they are assumed to have zero volume
2. …are in constant, straight-line motion
3. …experience elastic collisions in which no energy is lost
4. …have no attractive or repulsive forces toward each other
5. …have an average kinetic energy (KE)that is proportional
to theabsolute temp. of gas (i.e., Kelvin temp.)
AS TEMP. , KE
 explains why gases behave as they do
 deals w/“ideal” gas particles…