The document discusses valence bond theory and molecular orbital theory. It explains that valence bond theory describes covalent bond formation through the overlapping of atomic orbitals between atoms. There are two types of covalent bonds: sigma bonds formed by head-on overlapping and pi bonds formed by side-by-side overlapping. Hybridization involves mixing atomic orbitals to form new hybrid orbitals that can accommodate molecular geometries. Common hybridizations include sp, sp2, and sp3 which result in linear, trigonal planar, and tetrahedral geometries respectively.

1. MOLECULARMODEL
JONALYN M. SHENTON

2. Theory
Linus Pauling became the champion of the valence bond model,
which is easier to visualize and use.

3. Valence Bond Theory
The valence bond theory was proposed by Heitler and
London to explain the formation of covalent bond
quantitatively using quantum mechanics.
Later on, Linus Pauling improved this theory by introducing
the concept of hybridization.
Valence bond (VB) theory assumes that all bonds are localized
bonds formed between two atoms by the donation of an
electron from each atom.

4. Valence Bond Theory
Valence Bond theory describes covalent bond formation as
well as the electronic structure of molecules.
The theory assumes that electrons occupy atomic orbital's of
individual atoms within a molecule, and that the electrons of
one atom are attracted to the nucleus of another atom.

5. Valence Bond Theory
A covalent bond is formed by the overlapping of two half
filled valence atomic orbital's of two different atoms.
The electrons in the overlapping orbital's get paired and
confined between the nuclei of two atoms.
The electron density between two bonded atoms increases
due to overlapping. This confers stability to the molecule.

6. Valence Bond Theory
Greater the extent of overlapping, stronger is the bond
formed.
The direction of the covalent bond is along the region of
overlapping of the atomic orbital's i.e., covalent bond is
directional.

7. Valence Bond
Theory
BY: LINUS PAULING
JONALYN M. SHENTON
Overlap of Atomic Orbitals
The sharing of electrons between atoms is viewed
as an overlap of atomic orbitals of the bonding
atoms.
valence bond model says that an s-orbital on one atom overlaps
with an s-orbital on the other to form a bond.

8. Valence Bond Theory
σ-bond
A sigma bond (symbol: σ) is a covalent bond
formed via linear overlap of two orbital's.
π-bond
A pi bond (symbol: π) is a covalent bond
formed via parallel overlap of two orbital's.
There are two types of covalent bonds based on the pattern of
overlapping as follows:
Π bond

9. σ-bond
The covalent bond formed due to overlapping of atomic
orbital along the inter nucleus axis is called σ-bond. It is a
stronger bond and cylindrically symmetrical.
Depending on the types of orbital's overlapping, the σ-bond is
divided into following types:
(i): σs-s bond, (ii): σp-p bond, (iii): σs-p bond:

10. Sigma and Pi Bonds
Sigma bond: A covalent bond resulting from the formation of a
molecular orbital by the end-to-end overlap of atomic orbitals,
denoted by the symbol σ.
For example, the C-C sigma bond in ethane is formed by the head-on overlap of
two sp3 orbitals. The C-H sigma bonds are formed by s-sp3 overlap.
A double bond contains one sigma bond and one pi bond (side-on overlap of p
orbitals).

11. σ-bond
σs-s bond:

12. σ-bond
σs-p bond:

13. σ-bond
σp-p bond:

14. Pi bond: A covalent bond resulting from the formation of a
molecular orbital by side-to-side overlap of atomic orbitals along a
plane perpendicular to a line connecting the nuclei of the atoms,
denoted by the symbol π.

15. π-bond
The covalent bond formed by sidewise
overlapping of atomic orbital's is called π-
bond. In this bond, the electron density is
present above and below the inter
nuclear axis. It is relatively a weaker bond
since the electrons are not strongly
attracted by the nuclei of bonding atoms.
Note: The 's' orbital's can only form σ-bonds, whereas the p, d & f orbital's can form both σ and π-bonds.

17. In ethylene, the C=C double bond consists of a sigma bond and a pi bond. The sigma bond is formed by
the head-on overlap of two sp2 orbitals. The pi bond is formed by the side-on overlap of two 2p orbitals.

18. Illustration

19. H2 molecule
The electronic configuration of hydrogen atom in the
ground state is 1s1.
In the formation of hydrogen molecule, two half filled 1s
orbital's of hydrogen atoms overlap along the inter-
nuclear axis and thus by forming a σs-s bond.

20. Cl2 molecule
The electronic configuration of Cl atom in the ground
state is [Ne]3s2 3px
2 3py
2 3pz
1.
The two half filled 3pz atomic orbital's of two chlorine
atoms overlap along the inter-nuclear axis and thus by
forming a σp-p bond.

21. HCl molecule
In the ground state, the electronic configuration of
hydrogen atom is 1s1.
And the ground state electronic configuration of Cl atom is
[Ne]3s2 3px
2 3py
2 3pz
1.
The half filled 1s orbital of hydrogen overlap with the half
filled 3pz atomic orbital of chlorine atom along the inter-
nuclear axis to form a σs-p bond.

22. O2 molecule
The electronic configuration of O in the ground state is [He]
2s2 2px
2 2py
1 2pz
1.
The half filled 2py orbital's of two oxygen atoms overlap
along the inter-nuclear axis and form σp-p bond.
The remaining half filled 2pz orbital's overlap laterally to
form a πp-p bond.

23. O2 molecule
Thus a double bond (one σp-p and one πp-p) is formed
between two oxygen atoms.

24. N2 molecule
The ground state electronic configuration of N is [He]
2s2 2px
1 2py
1 2pz
1.
A σp-p bond is formed between two nitrogen atoms due
to overlapping of half filled 2px atomic orbital's along the
inter-nuclear axis.

25. N2 molecule
The remaining half filled 2py and 2pz orbital's form
two πp-p bonds due to lateral overlapping. Thus a triple
bond (one and two) is formed between two nitrogen
atoms.

27. Atomic orbital - An orbital that is associated with only
one particular atom. This is in contrast to molecular
orbitals, which are spread across a collection of atoms.

28. Similarly a 2s orbital is the region where there is
the greatest chance of finding the electron that’s
further from the nucleus - this is an orbital at the
second energy level. So the nearer an electron is to
the nucleus, the lower is its energy.
s orbitals.
The orbital occupied by the hydrogen electron is called
a 1s orbital. The "1"represents the fact that the orbital is
in the energy level closest to the nucleus. The"s”
gives the shape of the orbital. s orbitals are spherically
symmetric around the nucleus.

29. The p orbitals
A p orbital is rather like 2 identical balloons tied together at the
nucleus.
Since p orbitals occupy space in this manner, they can be said to have certain directions as well. The
two lobes can be pointing in any possible direction opposite to each other as the atom tumbles around in
space. This means that at any one energy level it is also possible to have three absolutely equivalent p
orbitals pointing at right angles to each other which can be in the x,y and z axis. Therefore p orbitals can
have directions px, py and pz as show in the illustration below.

30. SIGMA AND PI BOND

31. Hybridization
The intermixing of two or more pure atomic orbital's of an
atom with almost same energy to give same number of
identical and degenerate new type of orbital's is known as
hybridization.
The new orbital's formed are also known as hybrid orbital's.
During hybridization, the atomic orbital's with different
characteristics are mixed with each other.

32. Hybridization - The linear combination of atomic orbitals into
hybrid orbitals that accommodate particular molecular
geometries.
Hybrid orbital - Orbitals formed from the combination of
atomic orbitals that accommodate particular geometries.
Molecular orbital - Orbitals that, in contrast to atomic orbitals,
are distributed over an entire molecule instead of being
localized to specific atoms.

34. Hybridization
In chemistry, hybridization is the concept of mixing atomic orbitals to form new
hybrid orbitals suitable for the qualitative description of atomic bonding
properties. Hybridized orbitals are very useful in the explanation of the shape
of molecules. It is an integral part of valence bond theory.

35. Types of Hybridization
sp
sp2
sp3
sp3d
sp3d2
sp3d3

36. sp Hybridization
Intermixing of one 's' and one 'p' orbital's of
almost equal energy to give two identical
and degenerate hybrid orbital's is called 'sp'
hybridization.
These sp-hybrid orbital's are arranged
linearly at by making 180 ⁰ of angle.
They possess 50% 's' and 50% 'p' character.

37. sp2 hybridization
Intermixing of one 's' and two 'p' orbital's
of almost equal energy to give three
identical and degenerate hybrid orbital's
is known as sp2 hybridization.
The three sp2 hybrid orbital's are oriented
in trigonal planar symmetry at angles of
120 ⁰ to each other.
The sp2 hybrid orbital's have 33.3% 's'
character and 66.6% 'p' character.

38. sp3 hybridization
In sp3 hybridization, one 's' and three 'p'
orbital's of almost equal energy intermix
to give four identical and degenerate
hybrid orbital's.
These four sp3 hybrid orbital's are
oriented in tetrahedral symmetry with
109 ⁰ 28' angle with each other.
The sp3 hybrid orbital's have 25% ‘s’
character and 75% 'p' character.

39. sp3d hybridization
In sp3d hybridization, one 's', three 'p' and one 'd'
orbital's of almost equal energy intermix to give five
identical and degenerate hybrid orbital's, which are
arranged in trigonal bipyramidal symmetry.
Among them, three are arranged in trigonal plane and
the remaining two orbital's are present above and
below the trigonal plane at right angles.
The sp3d hybrid orbital's have 20% 's', 60% 'p' and
20% 'd' characters.

40. sp3d2 hybridization
Intermixing of one 's', three 'p' and two 'd' orbital's of
almost same energy by giving six identical and
degenerate hybrid orbital's is called sp3d2
hybridization.
These six sp3d2 orbital's are arranged in octahedral
symmetry by making 90 ⁰ angles to each other. This
arrangement can be visualized as four orbital's
arranged in a square plane and the remaining two are
oriented above and below this plane perpendicularly.

41. sp3d3 hybridization
In sp3d3 hybridization, one 's', three 'p' and three 'd'
orbital's of almost same energy intermix to give seven
sp3d3 hybrid orbital's, which are oriented in
pentagonal bipyramidal symmetry.
Five among the sp3d3 orbital's are arranged in a
pentagonal plane by making 72⁰ of angles. The
remaining are arranged perpendicularly above and
below this pentagonal plane.

42. ē Pair Hybridization Shape
2 sp linear
3 sp2 trigonal planar
4 sp3 tetrahedral, pyramidal, or bent
5 sp3d
trigonal bipyramidal, trigonal planar,
or linear
6 sp3d2 octahedral, square planar, or linear

44. Resonance
Resonance is a mental exercise and method within the Valence Bond Theory of bonding that
describes the delocalization of electrons within molecules. It compares and contrasts two or more
possible Lewis structures that can represent a particular molecule. Resonance structures are
used when one Lewis structure for a single molecule cannot fully describe the bonding that takes
place between neighboring atoms relative to the empirical data for the actual bond lengths between
those atoms.
Resonance forms illustrate areas of higher probabilities (electron
densities).
Resonance structures do not change the relative positions of the
atoms like your arms in the metaphor. The skeleton of the Lewis
Structure remains the same, only the electron locations change.

45. "Pick the Correct Arrow for the Job"
Most arrows in chemistry cannot be used interchangeably and care must be given to
selecting the correct arrow for the job.
 ↔↔: A double headed arrow on both ends of the arrow between Lewis structures is
used to show their inter-connectivity
 ⇌⇌: Double harpoons are used to designate equilibria
 ⇀⇀: A single harpoon on one end indicate the movement of one electron
 →→: A double headed arrow on one end is used to indicate the movement
of two electrons

46. Example 1: Ozone
Consider ozone (O3)

47. Delocalization and Resonance Structures Rules
In resonance structures, the electrons are able to move to help stabilize the molecule. This movement
of the electrons is called delocalization.
1. Resonance structures should have the same number of electrons, do not add or subtract any
electrons. (check the number of electrons by simply counting them).
2. All resonance structures must follow the rules of writing Lewis Structures.
3. The hybridization of the structure must stay the same.
4. The skeleton of the structure can not be changed (only the electrons move).
5. Resonance structures must also have the same amount of lone pairs.

48. Example 2: Thiocyanate Ion
Consider the thiocyanate (CNS−CNS−) ion.
SOLUTION
1. Find the Lewis Structure of the molecule. (Remember the
Lewis Structure rules.)

49. 2. Resonance: All elements want an octet, and we can do that in multiple ways
by moving the terminal atom's electrons around (bonds too).