Hybridization involves the mixing of atomic orbitals to form new hybrid orbitals that give molecules their shape and bonding properties. There are several types of hybridization depending on the orbitals involved, including sp3, sp2, and sp hybridization. Sp3 hybridization involves one s and three p orbitals mixing to form four sp3 hybrid orbitals arranged tetrahedrally as seen in methane. Sp2 hybridization is the mixing of one s and two p orbitals to yield three sp2 hybrid orbitals in a trigonal planar arrangement as in ethylene. Sp hybridization mixes one s and one p orbital to produce two linear sp hybrid orbitals as observed in acetylene. Hybridization explains molecular geometry and

1. Hybridization

2. Hybridization
• Hybridization: is the mixing of atomic orbitals to give new hybrid
atomic orbitals which have new shape and directional properties.
• Redistribution of the energy of orbitals of individual atoms to give
orbitals of equivalent energy occurs when two atomic orbitals
combine to form a hybrid orbital in a molecule. This process is called
hybridization.
• These hybrid atomic orbitals then combine with other atomic orbitals
to form the bonds in molecules.

3. • During the process of hybridization, the atomic
orbitals of comparable energies are mixed
together and mostly involves the merging of two
‘s’ orbitals or two ‘p’ orbitals or the mixing
of an ‘s’ orbital with a ‘p’ orbital, as well
as ‘s’ orbital with a ‘d’ orbital. The new
orbitals, thus formed, are known as hybrid
orbitals.
• More significantly, hybrid orbitals are quite
useful in explaining atomic bonding properties
and molecular geometry.

4. • Examples: Carbon atom. This atom forms 4 single
bonds wherein the valence-shell s orbital mixes
with 3 valence-shell p orbitals. This
combination leads to the formation of 4
equivalent sp3 mixtures. They will have a
tetrahedral arrangement around the carbon,
which is bonded to 4 different atoms.
• The atomic orbitals of the same energy level
mainly take part in hybridization. However,
both fully-filled and half-filled orbitals can
also take part in this process, provided they
have equal energy.

5. • On the other hand, we can say that the concept
of hybridization is an extension of the valence
bond theory, and it helps us to understand the
formation of bonds, bond energies and bond
lengths.

6. Types of Hybridization
• Based on the types of orbitals involved in
mixing, the hybridization can be classified as
sp3, sp2, sp, sp3d, sp3d2 and sp3d3.
• Let us now discuss the various types of
hybridization, along with their examples.

7. • sp3 Hybridization
• When one ‘s’ orbital and 3 ‘p’ orbitals belonging to
the same shell of an atom mix together to form four new
equivalent orbitals, the type of hybridization is
called a tetrahedral hybridization or sp3. The new
orbitals formed are called sp3 hybrid orbitals.
• These are directed towards the four corners of a
regular tetrahedron and make an angle of 109°28’ with
one another.
• The angle between the sp3 hybrid orbitals is 109.280
• Each sp3 hybrid orbital has 25% s character and 75% p
character.
• Examples of sp3 hybridization are ethane (C2H6) and
methane.

8. • So, in order to predict the valency and geometry of the carbon
atom, we are going to look at its electron configuration and the
orbitals.
• C – 1s22s22p2
• The valence electrons are the ones in the 2s and 2p orbitals and
these are the ones that participate in bonding and chemical
reactions.

11. • You can see from the electron configuration that it is impossible
to make four, identical in bond length, energy, and everything
else (degenerate) bonds because one of the orbitals is a
spherical s, and the other three are p orbitals. And this is where
we get into the need for a theory that can help us explain the
known geometry and valency of the carbon atom in many
organic molecules. So, Hybridization is a theory that is used
to explain certain molecular geometries that would have
not been possible otherwise.

12. • The sp3 hybridization
• Now, let’s see how that happens by looking at methane as an
example. In the first step, one electron jumps from the 2s to the
2p orbital. This leads to the excited state of the carbon:

13. • Pay attention that the electron goes uphill as the p subshell is
higher in energy than the s subshell and this is not energetically
favorable, but we will see how it is compensated in the next step
when orbitals are mixed (hybridized).
• So, in the next step, the s and p orbitals of the excited state
carbon are hybridized to form four identical in size, shape, and
energy orbitals.

15. • The number of the hybrid orbitals is always the same as the
number of orbitals that are mixed. So, four orbitals (one 2s +
three 2p) are mixed and the result is four sp3 orbitals. These are
hybrid orbitals and look somewhat like the s and p orbitals. And
again, we call them sp3 because they are formed from one s
orbital and three p orbitals.
• The formation of these degenerate hybrid orbitals compensates
for the energy uphill of the s-p transition as they have lower
energy than the p orbitals.

16. • The four sp3-hybridized orbitals arrange in a tetrahedral
geometry and make bonds by overlapping with the s orbitals of
four hydrogens: This explains the symmetrical geometry of
methane (CH4) where all the bonds have the same length and
bond angle.

17. • All four C – H bonds in methane are single bonds that are
formed by head-on (or end on) overlapping of sp3 orbitals of
the carbon and s orbital of each hydrogen.
• The bonds that form by the head-on overlap of orbitals are
called σ (sigma) bonds because the electron density is
concentrated on the axis connecting the C and H atoms.
• Ethane – CH3-CH3 and other alkanes
If instead of one hydrogen, we connect another sp3-hybridized
carbon, we will get ethane:

18. And consequently, in all the alkanes, there is a sigma bond between
the carbon atoms and the carbon-hydrogen atoms and the carbons
are sp3 hybridized with tetrahedral geometry:

19. To generalize this, any atom with four groups (either an atom or a lone pair)
is sp3 hybridized. And the way to look at this is, in order for the four groups to be as
far away from each other as possible like we learned in the VSEPR theory, the
groups need to be in identical four orbitals which is only possible in
the sp3 hybridization.

20. VSEPR theory
• To generalize this, any atom with four groups (either an atom or
a lone pair) is sp3 hybridized. And the way to look at this is, in
order for the four groups to be as far away from each other as
possible like we learned in the VSEPR theory, the groups need to
be in identical four orbitals which is only possible in
the sp3 hybridization.

23. • sp2 Hybridization
• sp2 hybridization is observed when one s and two p
orbitals of the same shell of an atom mix to form 3
equivalent orbitals. The new orbitals formed are
called sp2 hybrid orbitals.
• sp2 hybridization is also called trigonal
hybridization.
• It involves the mixing of one ‘s’ orbital and two ‘p’
orbitals of equal energy to give a new hybrid orbital
known as sp2.
• A mixture of s and p orbital formed in trigonal
symmetry and is maintained at 1200.
• All three hybrid orbitals remain in one plane and make
an angle of 120° with one another. Each of the hybrid
orbitals formed has a 33.33% ‘s’ character and 66.66%
‘p’ character.
• The molecules in which the central atom is linked to 3
atoms and is sp2 hybridized have a triangular planar

24. Examples of sp2 Hybridization
•All the compounds of Boron, i.e., BF3 and BH3
•All the compounds of carbon, containing a carbon-carbon
double bond, Ethylene (C2H4)

26. sp2 Hybridization cont…
• sp2 Hybridization
• Boron trifluoride (BF3) is predicted to have a trigonal planar
geometry by VSEPR. First a paired 2s electron is promoted to
the empty 2py orbital

27. • This is followed by hybridization of the three occupied orbitals to
form a set of three sp2 hybrids, leaving the 2pz orbital
unhybridized

28. • The geometry of the sp2 hybrid orbitals is trigonal planar,
• The angle between any two of the hybrid orbital lobes is 120°.
• Each can bond with a 2p orbital from a fluorine atom to form the
trigonal planar BF3 molecule.

30. • The process of sp2 hybridization is the mixing of an s orbital
with a set of two p orbitals (px and py) to form a set of three
sp2 hybrid orbitals. Each large lobe of the hybrid orbitals points
to one corner of a planar triangle.

31. • Other molecules with a trigonal planar electron domain
geometry form sp2 hybrid orbitals. Ozone (O3) is an example of
a molecule whose electron domain geometry is trigonal planar,
though the presence of a lone pair on the central oxygen makes
the molecular geometry bent. The hybridization of the central O
atom of ozone is sp2.
•

32. • During the formation of CH2=CH2, the electronic
configuration of carbon in its ground state
(1s2 2s2 2p1 2p1) will change to an excited state
and change to 1s2 2s1 2px12py1 2pz1. In the
excited state, since carbon needs electrons to
form bonds one of the electrons from
2s2 orbital will be shifted to the empty 2pz
orbital to give 4 unpaired electrons.

33. • To minimize the repulsion between electrons, the three sp2-hybridized orbitals are
arranged with a trigonal planar geometry.
• The shape of the sp2-hybridized orbital has be mathematically shown to to be
roughly the same as that of the sp3-hybridized orbital. To minimize the repulsion
between electrons, the three sp2-hybridized orbitals are arranged with a trigonal
planar geometry. Each orbital lobe is pointing to the three corners of an equilateral
triangle, with angles of 120° between them. Again, geometry and hybrization can be
tied together. Atoms surrounded by three electron groups can be said to have a
trigonal planar geometry and sp2 hybridization.

34. The unhybridized 2pz orbital is perpendicular to the plane
of the trigonal planar sp2 hybrid orbtals.

35. • In the ethylene molecule, each carbon atom is bonded to two
hydrogen atoms. Thus, overlap two sp2-hybridized orbitals with
the 1s orbitals of two hydrogen atoms for the C-H sigma bonds
in ethylene (sp2(C)-1s(H). Consequently, consistent with the
observations, the four carbon-hydrogen bonds in ethylene are
identical.

36. • he C-C sigma bond in ethylene is formed by the overlap of an
sp2 hybrid orbital from each carbon.

37. • The overlap of hybrid orbitals or a hybrid orbital and a 1s orbtial
from hydrogen creates the sigma bond framework of the
ethylene molecule. However the unhybridized pz orbital on each
carbon remains.

38. • The unhybridized pz orbitals on each carbon overlap to a π
bond (pi). The orbital overlap is commonly written as pz(C)-
1pz(C). In general multiple bonds in molecular compound are
formed by the overlap of unhybridized p orbitals. It should be
noted that the carbon-carbon double bond in ethlene is made
up of two different types of bond, a sigma and a pi.

39. Overall, ethylene is said to contain five sigma bonds and one pi bond. Pi
bonds tend to be weaker than sigma bonds because the side-by-side
overlap the p orbitals give a less effective orbital overlap when compared
to the end-to-end orbital overlap of a sigma bond. This makes the pi much
easier to break which is one of the most important ideas in organic
chemistry reactions

40. • An ethylene molecule is said to be made up of five sigma bonds and one
pi bond. The three sp2 hybrid orbitals on each carbon orient to create the
basic trigonal planer geometry. The H-C-C bond angle in ethylene is
121.3o which is very close to the 120o predicted by VSEPR.
• The four C-H sigma bonds in ethylene . The carbon-carbon double bond
in ethylene is both shorter (133.9 pm) and almost twice as strong (728
kJ/mol) than the carbon- carbon single bond in ethylene (154 pm & 377
kJ/mol). Each of the four carbon-hydrogen bond in ethylene are
equivalent has have a length of 108.7 pm

42. • Rigidity in Ethene
• Because they are the result of side-by-side overlap (rather then end-to-end
overlap like a sigma bond), pi bonds are not free to rotate. If rotation about
this bond were to occur, it would involve disrupting the side-by-side
overlap between the two 2pz orbitals that make up the pi bond.
• If free rotation were to occur the p-orbitals would have to go through a
phase where they are 90° from each other, which would break the pi bond
because there would be no overlap. Since the pi bond is essential to the
structure of ethene it must not break, so there can be not free rotation
about the carbon-carbon sigma bond. The presence of the pi bond thus
‘locks’ the six atoms of ethene into the same plane.

45. Summary of H-C=C-H
• Meanwhile, out of 2s, 2px, 2py, and 2pz orbitals
in carbon, only 2px, 2py, and 2s take part in
hybridization. One 2pz orbital remains unchanged.
This leads to the formation of three sp2 hybridized
orbitals. The hybrid orbitals look like
sp3 orbitals, but they are sp2 orbitals as they are
fatter and shorter.
• The molecular orbitals after hybridization now
form different bonds between the electrons. One
carbon atom overlaps the sp2 orbital of another
carbon atom to form sp2 – sp2 sigma bond. The two
sp2 hybrid orbitals get overlapped by two hydrogen
atoms containing unpaired electrons. A pi bond is
formed by the unhybridized 2pz orbitals of each
carbon atom.

46. • sp Hybridization
• sp hybridization is observed when one s and one
p orbital in the same main shell of an atom mix
to form two new equivalent orbitals. The new
orbitals formed are called sp hybridized
orbitals. It forms linear molecules with an
angle of 180°.

47. • This type of hybridization involves the mixing
of one ‘s’ orbital and one ‘p’ orbital of equal
energy to give a new hybrid orbital known as an
sp hybridized orbital.
• The sp hybridization is also called diagonal
hybridization.
• Each sp hybridized orbital has an equal amount
of s and p characters – 50% s and 50% p
characters.

49. P ORBITAL

50. • Examples of sp Hybridization:
• All compounds of beryllium, like BeF2,
BeH2, BeCl2
• All compounds of a carbon-containing triple
bond, like C2H2.
• Formation of beryllium fluoride (BeF2). Beryllium (4Be) atom has a
ground state configuration as 1s2, 2s2. In the excited state one of the
2s-electron is promoted to 2p-orbitals. One 2s-orbital and one 2p-
orbitals of excited beryllium atom undergo hybridisation to form two
sp-hybridised orbitals as described in figure below.

51. Due to the sp-
hybridised state of
beryllium, BeF2 molecul
e has linear shape.

53. • The two sp-hybrid orbitals of carbon atom are linear and are
directed at an angle of 180° whereas the unhybridised p-orbitals
are perpendicular to sp-hybrid orbitals and also perpendicular
to each other.

54. • In the formation of acetylene, carbon atom uses its one of
the sp-hybrid orbital for overlapping with similar orbital of the
other carbon to form C—C sigma bond. The other sp-hybrid
orbital of each C atom overlaps axially with 1s-orbital of H atom
to form C—H sigma bond. Each of the two unhybridised orbitals
of both the carbon atoms overlap sidewise to form two bonds.
The electron clouds of one bond lie above and below the
internuclear axis whereas those of the other bond lie in front
and back of the inter-nuclear axis.

55. The four clouds so formed further merge into one another
to form a single cylindrical electron cloud around the
internuclear axis representing C—C sigma bond.

56. •

57. PCl5

58. • Now the five orbitals (i.e., one s, three p and one d orbitals) are
available for hybridization to yield a set of five sp3d hybrid orbitals
which are directed towards the five corners of a trigonal bipyramidal
as depicted in the below.

59. • substituents in the plane of the paper with normal bond lines. The other two
substituents are drawn going into the paper with a dashed wedged bond, or
coming out of the paper with a solid wedged bond.
• The Fischer projection is a less common alternative. By definition in these
drawings, the vertical bonds go into the paper, and the horizontal bonds come out
of the paper.
You do not always have to show the full stereochemistry (3-D shape) of a
Three-Dimensional Representation Around Atomic Centre

60. • PCl5 has trigonal bipyramidal shape.
• It should be noted that all the bond angles in trigonal bipyramidal geometry are
not equivalent.
• In PCl5 the five sp3d orbitals of phosphorus overlap with the singly occupied p
orbitals of chlorine atoms to form five P-Cl sigma bonds. Three P-Cl bond lie in
one plane and make an angle of 120 with each other; these bonds are termed as
equatorial bonds.
• The remaining two P-Cl bonds-one lying above and the other lying below the
equatorial plane, make an angle of 90 with the plane. These bonds are called
axial bonds.
• As the axial bond pairs suffer more repulsive interaction form the equatorial bond
pairs, therefore axial bonds have been found to be slightly longer and hence
slightly weaker than the equatorial bonds; which makes PCl5 molecule more
reactive.

61. SP3D2

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

63. • xample: SF6
• In SF6, one electron each from 3s and 3p orbitals is promoted to 3d
orbitals The six orbitals get hybridized to form six sp3d2 hybrid
orbitals. Each of these sp3d2 hybrid orbitals overlaps with 2p orbital
of fluorine to form S–F bond.
• Thus, SF6 molecule has octahedral structure. The dotted electrons
represent electrons from F–atoms.

65. Important of Hybridisation
• It use to determine geometric arrangement of molecules or
compounds.
• To determine shape or structure of molecules or compounds.
• To determine reactivity of molecule
• Describe the sharing of electrons before atom in a covalent bond.

66. • Assignment
• Write on shape and bond in water
• SP3D3 hybridization